IGCSE CHEMISTRY NOTES
IGCSE Chemistry-Dr. D. Bampilis
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IGCSE Chemistry-Dr. D. Bampilis
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1 Τhe particulate nature of matter Atom: The smallest particle of matter Molecule: A small particle made from more than one atom bonded together Element: A substance made of only one type of atom Compound: A substance made from two or more different elements bonded together States of matter: Solid: 1. Strong forces of attraction between particles 2. Have a fixed pattern (lattice) 3. Atoms vibrate but can’t change position therefore fixed volume and shape 4. Can’t be compressed Liquid: 1. Weaker attractive forces than solids 2. No fixed pattern, liquids take up the shape of their container but have a fixed volume 3. Particles slide past each other. 4. Can’t be compressed Gas: 1. Almost no intermolecular forces 2. Particles are far apart, and move quickly 3. They collide with each other and bounce in all directions. 4. Can be compressed The Kinetic Theory of Matter States: The kinetic theory is a theory put together by the finest chemists and physicians of all time. It consists of a number of true facts related to matter and their states. The theory explains the behavior of matter and their physical properties. The kinetic theory of matter states: All matter is made up of tiny, microscopic moving particles. And each matter has a different type of particles with different size and mass. Particles are in continuous movement. All particles are moving all the time in random directions (Brownian motion). The speed of movement depends on the mass of the particle, temperature and several other factors that you will know later on. Kinetic means movement, and so kinetic energy means movement energy.
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Comparing Properties of Solids, Liquids and Gases: Solid
Liquid
Gas
Molecular Structure
Particles Arrangement Intermolecular Spaces Intermolecular Forces Movement of Particles Shape
Volume
Very closely packed Closely packed Regularly arranged in Irregular arrangement lattice Almost none Minimal Negligible Tiny spaces Not weak Extremely strong Weaker than in solids Vibrating in a fixed Slowly slide over each position other randomly No fixed shape Fixed definite shape Depends on the container fixed
Compressibility Cannot be compressed Diffusion
Cannot diffuse
fixed Can be hardly compressed Diffuses slowly
Very far apart Very irregular arrangement Very large Very weak Fast movement in random direction No fixed shape No fixed volume Depends on the container Very compressible Diffuses quickly
• melting – freezing – boiling – condensing – subliming – desubliming
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Diffusion is the random movement of liquid or gas particles to fill the available space and spread evenly. For instance, if you pass by a trash can, you can smell the ugly scent of trash. This is because molecules from the garbage diffused out of the can to the air which you breathed in.
Diffusion rate depends on several factors, these are: Mass of the substance. The lighter the substance (lower M r or Ar) the faster it diffuses Temperature. The more kinetic energy the particles have, the faster they move and diffuse. Presence of other substance. Diffusion is faster when it occurs in an area where there are fewer particles of other substances present. This is why diffusion is extremely fast in vacuums. This is because the diffusing particles have less other particles to stand in their way. Intermolecular spaces. This is why gases diffuse faster than liquids and solids do not diffuse.
TIPS ON SPECIFIC TOPICS • Remember that most of the particles in liquids are touching one another. It is a common error to think that they are well separated. •Diffusion is due to the random movement of particles so they spread out everywhere. In an exam, try not to give an answer involving movement of particles from high to low concentration as this suggests that the particles know where they are going.
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2 Experimental techniques 2.1 Αpparatus for the measurement Time(s): Stopwatch Clock Temperature(0C): Thermometer Mass(kgr-gr): Balance Volume(m3-dm3-cm3)of liquids: beaker -burette -pipettes -measuring cylinder -volumetric flask of gases: gas syringe-upturned measuring cylinder 2.2 (a) Criteria of purity • Paper chromatography is a technique that can be used to separate mixtures of dyes or pigments and is used to test the purity of a mixture or to see what it contains •different solubility in the solvent • different degrees of attraction for the filter paper • The importance of purity in substances in everyday life, e.g. foodstuffs and drugs Purity can be measured in a number of ways: • melting point/boiling point (impurities increase b.p. and decrease m.p.) • chromatography • Rf values
. Rf = distance moved by the compound ÷ distance moved by the solvent
• Chromatography techniques can be applied to colourless substances by exposing chromatograms to substances called locating agents(ninhydrin for aminoacids)
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2.2 (b) Methods of purification • Evaporation used to separate a solid from a solution. • Filtration to remove solid particles from a liquid.
• Crystallization removes the solvent, to leave the solute. • Distillation used to separate a solvent from a solution. 1. Salty water is heated
2. The water vapour cools in the condenser and drips into a beaker
3. The water has condensed and is now in the beaker, the salt stays behind
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• Fractional distillation used to separate liquids from each other, produces a number of substances from the original mixture (e.g. petroleum). 1. Water and ethanol solution is heated
2. The ethanol evaporates first, cools, then condenses
3. The water left evaporates, cools, then condenses
• Sedimentation allows an insoluble solid to separate out and sink to the bottom of a container
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• Centrifugation a spinning motion increases the force of gravity that quickly separates a solid from a suspension.
• Decanting pouring off liquid (e.g. pouring off excess water from a pot of peas).
• Magnetic Separation a method for separating one solid (usually iron) from a mixture of solids, very useful for separating aluminium cans from steel cans. • Solvent extraction used to separate two solutes dissolved in a solvent.
• Chromatography used to separate different substances from a solution
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TIPS ON SPECIFIC TOPICS •When measuring out volumes, think about the accuracy needed. A burette or volumetric pipette is far more accurate than a measuring cylinder. •When drawing chromatography apparatus, you must draw the origin line on the chromatogram so that it is above the starting level of the solvent. •Remember that pure substances have definite sharp melting points and boiling points. Impure substances melt and boil over a range of temperatures. •When describing crystallization, the answer ‘heat the solution’ is not enough. You need to write ‘evaporate off some of the water and then leave to cool’. •When choosing a method to purify a mixture, think about the states and solubilities of the substances in the mixture. • If you are distilling an aqueous solution of a salt, the salt itself does not evaporate as it has too high a boiling point. Only the water evaporates
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3 Atoms, Elements and Compounds 3.1 Atomic structure and the Periodic Table • Protons, neutrons and electrons • Electrons shells / energy levels • Proton/atomic number and nucleon/mass number • Periodic Table
• Isotopes • Types of isotopes as being radioactive and non-radioactive • Medical and industrial uses of radioactive isotopes • Electron configuration
• Period(number of shells)
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• Group(number of e in outer shell) • Valence electrons – chemical properties- group • Noble gases electronic structures 3.2 Bonding: the structure of matter • Element a substance that cannot be split into anything simpler, in a chemical reaction. Each element has a unique proton number. • Mixture two or more elements mixed together BUT that are not chemically combined • Compound a substance in which two or more different elements are chemically combined (molecular – ionic). Differences are made up of two or more
compound pure substances elements combined chemically a fixed ratio fixed , different from its constituents only by chemical methods
composition properties can be separated
mixture impure substances substances mixed physically varying ratios no fixed ,same of its constituents. easily by physical methods
• Density Differences in Physical Properties conductors of heat and electricity malleable - ductile lustrous at room temperature melting point density sonorous
Metal
Non-metal
good
poor
yes yes - shiny solids (exception is mercury)
no - brittle no -dull solids or gases(exception bromine) low(exception C,Si) low(most) no
high(exception group I, Hg) high(exception group I, Ga) yes
is
• Alloy, such as brass, a mixture of a metal with other elements
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3.2 (a) Ions and ionic bonds • Formation of ions by electron loss(cations) or gain(anions) • Ionic bonds between metallic and non-metallic elements Group Group Group Group Group Group Group Group 1 2 3 4 5 6 7 0 Example element Na
Mg
Al
C
N
O
Cl
He
Charge
+
2+
3+
Note 1 3-
2-
-
Note 2
Symbol of ion
Na+
Mg2+
Al3+
Note 1 N3-
O2-
Cl-
Note 2
Note 1: Carbon and silicon in Group 4 usually form covalent bondsby sharing electrons. Note 2: The elements in Group 0 do not react with other elements to form ions. • Octet rule • Lattice structure of ionic bond The diagram shows part of the crystal lattice of sodium chloride:
Properties of Ionic Compounds: Hard solids at room temperature, High melting and boiling points because of strong attraction forces, When solid they are electrical insulators but conduct electricity when molten or aqueous, Water soluble.
3.2 (b) Molecules and covalent bonds • Single covalent bonds in H2, Cl2 , H2O, CH4 and HCl as the sharing of pairs of electrons leading to the noble gas configuration (non metal- non metal) • Shared pair- lone pair • Electron arrangement in more complex covalent molecules such as N 2, C2H4, CH3OH and CO2 , Double- triple bond
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How many bonds? Element
Number of bonds
Group 4
Carbon
8-4=4
Group 5
Nitrogen
8-5=3
Group 6
Oxygen
8-6=2
Group 7
Chlorine
8-7=1
Hydrogen forms one covalent bond. The noble gases in Group 0 do not form any. • Intermolecular forces: weak – m.p. , b.p.
• Valency of an atom: the number of electrons that would be gained, lost or share if it reacts with other atoms. Types of Covalent Structures: There are two types of covalent structures: Simple Molecular Structure Giant Molecular Structure
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Simple Molecular Structure: They are simple and contain only a few atoms in one molecule. Covalent bonds between the atoms within a molecule (intermolecular bonds) are strong but they have weak bonds between molecules (intermolecular bonds). These forces increase as the size of the molecule increases. Giant Molecular Structure: They are also known as macromolecular structures. One molecule contains hundreds of thousands of atoms. They have extremely strong bonds between the atoms (intermolecular bonds). Properties of Covalent Compounds: Simple molecular structures are usually gases or liquids and sometimes solids with low melting points; this is because of weak forces of attraction between the molecules which can be broken easily. Giant molecular structures have very high melting points because the whole structure is held together with very strong covalent bonds. Most of them do not conduct electricity Most of them are insoluble in water
Differences in Chemical Properties electrons in the outer shell valence electrons form
Metal 1-3(exception is hydrogen) lose easily cations
form oxides react with acid
basic form hydrogen
Non-metal 4-8 gain or share anions(exception is hydrogen) acidic no
Differences
ionic compounds
covalent compounds
volatility melting points and boiling points solubility in water
high high
low low
usually soluble
electrical conductivity
molten or dissolved in water ionic solids are good insulators crystal lattice are hard
the majority do not dissolve don't conduct electricity in an aqueous solution
form
molecules tend to be soft and relatively flexible
3.2 (c) Macromolecules • Giant covalent structures of graphite and diamond(allotropes of C) What are allotropes? When an element exists in several physical forms of the same state, it is said to exhibit allotropy. Each form of this element is an allotrope. Lots of elements exhibit allotropy. Carbon has two very popular allotropes, diamond and graphite. Diamond and graphite are both made of carbon only. However, they look very different and have different physical properties. They are both giant molecular structures.
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• Graphite: 3 covalent bonds hexagon – layers(weak bonding) delocalized e conductor(electods) high m.p., b.p. used as a lubricant and in pencil leads(can flake off easily) • Diamond : 4 covalent bonds insulator high m.p., b.p. very hard used for cutting and drilling • Macromolecular structure of silicon (IV) oxide (silicon dioxide-sand - quartz) Similarity in properties between diamond and silicon (IV) oxide, related to their structures
The graphic shows the molecular structure of graphite and diamond(two allotropes of carbon) and of silica (silicon dioxide). 3.2 (d) Metallic bonding • Metallic bonding is a lattice of positive ions in a ‘sea of electrons’ • The electrical conductivity, malleability and high m.p. and b.p. of metals The electrical conductivity of a metal decreases with increasing temperature(the vibration of cations inhibit the move of electrons)
TIPS ON SPECIFIC TOPICS •In an exam you will always be given a Periodic Table .You can use your Periodic Table to find out the number of protons in an atom. You can also use it to calculate the number of neutrons. •You did not need to know the details about radioactivity or about α-, β- or γ-radiation. Don’t try to remember lots of uses for radioisotopes –just remember one medical and one industrial use. •Make sure that you can draw the electronic structure of the first 20 elements in rings containing electrons. If you are simply asked ‘what is the electronic structure of sodium?’ •You should learn the definitions of elements, compounds and mixtures. You may be asked to write these definitions in an exam. •If you are asked how to tell the difference between a metal and a non-metal it is best to select conductivity, malleability or ductility as properties. These have fewer exceptions to the general rules.
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•When drawing the electronic structure for an ion, make sure that the charge of the ion is shown at the top- right-hand corner just outside the square brackets. Do NOT put the charge in the nucleus. • When drawing dot-and-cross diagrams remember to pair up the bonding electrons in the overlap area between the atoms. Don’t put them outside the area where the atoms join. • When drawing the electronic structure of compounds with double and triple bonds, make sure that you draw the atoms large enough so that all the bonding electrons can fit into the overlap area of the atoms. •Remember that compounds of metals with non-metals are likely to be ionic. Compounds of non-metals with other non-metals are covalent. •When explaining why graphite conducts electricity, make sure that you state that electrons in the layers can move along. Do not write ‘The electrons move.’- that suggest that the electrons in the covalent bonds can move through the structure as well. •It is a common error to suggest that conduction in metals is due to moving ions. Remember that it is only the delocalized electrons which move. The positive ions remain fixed in position within the giant lattice. •When writing symbols containing two letters, make sure that the second letter is a small one. Cl is correct for chlorine. CL is wrong. •Take care when writing the second atom in a formula. Co 2 is not acceptable for carbon dioxide and neither is H2o for water. The symbol for oxygen is always a capital O. •When asked to write the formula of an ionic compound from a diagram of its structure, make sure that you write the formula as the simplest ratio. For example, CaBr 2 not a Ca8Br16. • It is a common mistake to count the bonds and not the electrons when asked about the number of electrons shared between the atoms in a molecule. For example, the number of shared electrons in methane is eight not four • Take care when writing electronic structures including hydrogen. Always show the hydrogen atom either as a circle or (if ionic) by its symbol. It is best practice to write the symbol of the atom in the centre so it is clear to the examiner which atom is which • When writing dot-and-cross diagrams for ionic structures, put the charge outside of the brackets, at the top, not in the centre of the atom • When asked about the number of covalent bonds in a compound, focus on the outer energy level / shell electrons that are shared, not the total number of electrons. Remember that some molecules have non-bonding pairs of electrons e.g. nitrogen • When drawing dot-and-cross diagrams for molecules such as nitrogen which have only three bonding pairs of electrons, don’t forget to draw in the lone pairs of electrons. Remember that there must be eight electrons surrounding each atom • Practice drawing diagrams of giant molecule structures, including silicon dioxide, diamond and graphite, as these are nearly always drawn badly. You must show the continuation bonds
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4 Stoichiometry • Symbols of the elements • Formulae of simple compounds (group- valency-formula) • Naming compounds AXBY : A B ide • Molecular formula – Empirical formula - Structural formula • Deduce the formula of a simple compound from a model or a diagrammatic representation Molecular compounds from structural formula Ionic compounds from simplest ratio • Determine the formula of an ionic compound from the charges on the ions present The Periodic Table and Charges:
Group (Charge)
Ions present
1 (+1)
Li+ Na+ K+
2 (+2)
Transition metals
Be2+ Mg2+ Ca2+ Ba2+
Cu2+ / Cu+ Fe2+ / Fe3+ Zn2+ Ag+
3 (+3)
4 (±4)
Al3+
C Si Pb2+
5 (3)
6 (2)
3-
2-
N P3-
O S2-
Compound Ions: Oxidation State
Name
Symbol
+1
Ammonium Ion
NH4+
-1
Hydroxide Ion Nitrate Ion Nitrite Ion Manganate(VII) Oxide Ion Hydrogen Carbonate Ion
OHNO3NO2MnO4HCO3-
-2
Carbonate Ion Sulfate Ion Sulfite Ion Dichromate (Vi) Ion
CO32SO42SO32Cr2O72-
-3
Phosphate Ion Phostphite Ion
PO43PO33-
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7 (1)
FClBrI-
• Word equations • Simple balanced chemical equations • Diatomic elements :H2, O2, N2, F2, Cl2 , Br2, I2 • Construct equations with state symbols • Ionic equations Substances form ions: -metals/non metals -acids -ammonium compounds Spectators ions • Relative atomic mass, Ar
Calculating relative atomic mass from isotopic abundance
• Relative molecular mass, Mr , as the sum of the relative atomic masses • Relative formula mass or Mr for ionic compounds • Calculations involving reacting masses in simple proportions 4.1 The mole concept • mole • Avogadro constant • Molar mass: the relative formula mass in g
n= or m=nM n:number of mole m: mass in g M:molar mass • Molar gas volume, taken as 24 dm3 at room temperature and pressure (200C – 1 atm) • Solution concentrations expressed in g/dm3 and mol/dm3 C= n=CV C: concentration in mol/dm3 V:volume of solution in dm3 n:mole of solute
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• Stoichiometry • Limiting reactants • Calculate stoichiometric reacting masses and volumes of gases and solutions • Titration • Percentage by mass of an element in a compound : using Ar and relative formula mass • Calculate empirical formulae and molecular formulae • % yield=
100
• % purity=
100
TIPS ON SPECIFIC TOPICS •When balancing symbol equations you must not change any of the formulae. Always balance by putting large numbers in front of the formulae. For example, balancing CaΟ by making it into CaO2 is wrong. •When writing ionic equations, first identify the reactants or products that are not ionic. These will be solids, liquids or simple molecules like chlorine. It is only then that you can separate the other compounds into ions. •If a formula has brackets, first work out the atomic masses inside the brackets then multiply by the number outside. Finally, add the atomic masses which were not bracketed. •When doing calculations put the relative formula masses or moles below the appropriate reactants or products in the symbol equation so that you can see the reactants or products are relevant. Be sure to take the stoichiometry of the equation into account. •The limiting reactant is the reactant that is NOT in excess. It has the smaller number of moles. Be careful though – you must also take into account the ratio in which the reactants combine. •When out gas volumes first find the number of moles and then multiply this by 24. The answer is then in dm3. Remember that the molar gas volume is given at the bottom of your Periodic Table. •Always show your working in calculations if a question is worth more than one mark. If you make an error at the start – for example use an incorrect molar mass – you can still gain marks. •When calculating empirical formulae, make sure that between steps 1 and 2 you don’t round up the figures. This often leads to errors. •Mole calculations involving concentrations are easier if you change cm3 to dm3 and then use the formula concentration= number of moles/ volume of solution in dm3 • If asked for a word equation, do not write a symbol equation. A word equation tests knowledge of chemical names. Although a correct symbol equation is often accepted this is not guaranteed and if you make an error, you won’t get the mark • A common error is to think that a nitrate ion has a 2- charge. The formula for the nitrate ion is NO3-.This makes the formula for nitric acid HNO3
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• The charge on a silver ion is 1+. A common mistake is to think that silver has a 2+ charge • When working out formulae, don’t be confused by oxidation numbers. A common mistake is to think that the formula for lead(IV) oxide is PbO 4 or that lead(II) nitrate is Pb2(NO3). In a formula you have to balance the positive and negative charges. Lead(IV) = 4+, lead(II) = 2+, oxide = 2- and nitrate = 1-. So lead(IV) oxide is PbO2, and lead(II) nitrate is Pb(NO3)2 • If asked to name a salt formed in a particular reaction, don’t put down any other product or you will lose a mark • When calculating moles, if you are given an equation such as: Mg + 2CH3CO2H → (CH3CO2)2Mg + H2 ignore the 2 in the equation when calculating the molar mass of ethanoic acid. The molar mass of ethanoic acid is 60, not 120. However, remember when calculating reacting masses that the 2 needs to be taken into account
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5 Electricity and chemistry • Electrolysis is a process in which electricity is used to break compounds down into their elements. The mixture being electrolysed is called an electrolyte and must be liquid (either melted or dissolved) to allow the ions to move. Electrolysis cell Electrodes • General principle that metals or hydrogen are formed at the negative electrode (cathode), and that non-metals (other than hydrogen) are formed at the positive electrode (anode) • Describe the electrode products in the electrolysis of: – molten lead(II) bromide – concentrated hydrochloric acid – concentrated aqueous sodium chloride between inert electrodes (platinum or carbon) Electrolysis of Molten Ionic Compounds: An idealized cell for the electrolysis of sodium chloride is shown in the figure below. A source of direct current is connected to a pair of inert electrodes immersed in molten sodium chloride. Because the salt has been heated until it melts, the Na + ions flow toward the negative electrode and the Cl- ions flow toward the positive electrode. Negative electrode (cathode): Na+ + e- → Na Cl- ions that collide with the positive electrode are oxidized to Cl 2gas, which bubbles off at this electrode. Positive electrode (anode): 2Cl- → Cl2 + 2eThe net effect of passing an electric current through the molten salt in this cell is to decompose sodium chloride into its elements, sodium metal and chlorine gas. 2NaCl(l) → 2 Na(l) + Cl2(g)
This example explains why the process is called electrolysis. The suffix -lysis comes from the Greek stem meaning to loosen or split up. Electrolysis literally uses an electric current to split a compound into its elements.
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Electrolysis of Aqueous Ionic Compounds: Electrolysing an ionic compound in its solution is very much different to electrolysing it when it’s molten. This is because in a solution we have 4 ions, H +and OH- from water and a positive and a negative ion from the compound. But only one type of ions gets discharged at each electrode. For the positive ions, the one that gets discharged at the cathode is the least reactive one. This is because least reactive elements have more tendencies to be an atom. So if the ion from the ionic compound is above hydrogen in the reactivity series (more reactive), H+ gets discharged at the anode And if the ion from the compound is below hydrogen in the reactivity series (less reactive), this ion gets discharged at the cathode. So for example if we are electrolysing aqueous sodium chloride, H+ ions will get discharged at the cathode because sodium is more reactive than hydrogen. And if we are electrolysing aqueous copper iodide, Cu2+ ions will get discharged at the cathode because copper is less reactive than hydrogen.
For the negative ions however it is different. Oxygen from OH - from water is always discharged at the anode except in one case, this is if the other negative ion is a halide. If oxygen from OH is discharged, the equation will be: 4OH- - 4e → O2 + H2O If the other negative ion is a halide, there are two probabilities: 1.Oxygen from OH- gets discharged at the cathode, 2.The halide ion gets discharged at the cathode. It all depends on the concentration of the halide. If the electrolyte is a concentrated solution, then there are many of the halide ions, more than OH-. So the halide ion gets discharged at the cathode. If the electrolyte is a dilute solution, then there are more OH - ions than halide ions, so oxygen from OH- gets discharged.
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So for example if the electrolyte is a concentrated solution of sodium chloride, hydrogen gas is formed at the cathode because hydrogen is less reactive than sodium. And chlorine gas is formed at the anode because the solution is concentrated. If the electrolyte is a dilute solution of silver sulfate, silver is formed at the cathode because it is less reactive than hydrogen and oxygen gas is formed at the anode. • Predict the products of electrolysis of a specified halide in concentrated or dilute aqueous solution Water is a weak electrolyte: H2O(l) Discharge series:
H+(aq)+OH–(aq)
Cu2+, H+,Al3+, Mg2+,Na+ I–,Br–,Cl–,OH–,NO3–,SO42–
Anode(+) - anions(–) - lose e– - oxidation Cathode(–) - cations(+) – gain e– - reduction Half equation Aqueous solutions: 2H+(aq)+ 2e– H2(g) 4OH–(aq) O2(g)+ 2 H2O (l) + 4e– This table shows some common ionic compounds (in solution), and the elements released when their solutions are electrolysed using inert electrodes, eg carbon electrodes: Ionic substance
Element at -
Element at +
Copper chloride, CuCl2
Copper, Cu
Chlorine, Cl2
Copper sulfate, CuSO4
Copper, Cu
Oxygen, O2
Sodium chloride, NaCl
Hydrogen, H2
Chlorine, Cl2
Hydrochloric acid, HCl
Hydrogen, H2
Chlorine, Cl2
Sulfuric acid, H2SO4
Hydrogen, H2
Oxygen, O2
Very dilute solutions of halide compounds If a halide solution is very dilute (eg NaCl), then oxygen will be given off instead of the halogen. This is because the halide ions are outnumbered by the hydroxide ions from the water. • The manufacture of chlorine and sodium hydroxide from concentrated aqueous sodium chloride (brine) The ions in solution: Na+, H+, Cl–,OH– Anode(+): 2Cl–(aq) Cl2(g)+2e– Cathode(–): 2H+(aq) +2e– H2(g) remain in solution: Na+, OH– • Refining of copper(purification by electrolysis) electrolysis aqueous copper(II) sulfate using copper electrodes
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Anode(+):Cu(s) Cu2+(aq) +2e– Cathode(–):Cu2+(aq) +2e– Cu(s) the electrolyte remains the same deep blue colour The pure copper rod is connected to the negative terminal of a battery, and the impure rod is connected to the positive terminal
The pure copper rod has increased in size, while the impure rod has deteriorated, leaving a pool of anode sludge at the bottom of the beaker
The electrolysis aqueous copper(II) sulfate using carbon electrodes Anode(+):4OH–(aq) O2(g)+ 2 H2O (l) + 4e– Cathode(–):Cu2+(aq) +2e– Cu(s) the electrolyte gradually loses its blue colour • The electroplating of metals How it works The negative electrode should be the object to be electroplated. The positive electrode should be the metal that you want to coat the object with. The electrolyte should be a solution of the coating metal, such as its metal nitrate or sulfate. Anode(+):Me(s) Mex+(aq) +xe– Cathode(–):Mex+(aq) +xe– Me(s) Me: Ag,Au,Cu,Ni,Sn,Cr • Uses of electroplating: protection from corrosion - appearance • The manufacture of aluminium from pure aluminium oxide in molten cryolite The ore crushed and mixed with NaOH Al2O3(s) + 2 NaOH(aq) 2NaAlO2(aq) + H2O (l) The impurities are insoluble The sodium aluminate heated to make up Al 2O3 The Al2O3 dissolved in molten cryolite(Na3AlF6) and CaF2 to lower its m. p. and improves the conductivity Anode/graphite(+):2O2– O2(g)+ 4e–(O2 react with C to form CO2) Cathode/ graphite( (–):Al3+ +3e– Al
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The overall reaction: Al2O3 4Al + 3O2 • Conductors copper : good conductor – ductile – easily purified by electrolysis steel-cored aluminium in cables Al: good conductor – low density – resistant to corrosion steel : additional strength Thin – thick wires: larger electric current – heat /melt • Insulators plastics : flexible – not biodegradable – high electric current :thermosetting ceramics: high m.p.- not affected by water /air
TIPS ON SPECIFIC TOPICS •Remember that in an electrolyte, it is the ions that move, not the electrons. • Remember that when a solution of sodium chloride is electrolysed, hydrogen is formed at the cathode whereas with molten sodium chloride, sodium is formed. • Make sure that you know the difference in the products at each electrode when dilute and concentrated aqueous sodium chloride and molten sodium chloride are electrolysed. • Remember that in electrolysis the electrodes are usually inert (graphite or platinum). If the anode is not inert, it will react and decrease in size. •When asked questions about what you observe during electroplating, the answer expected is what you see happening at each electrode and any changes in the colour of the electrolyte. •You do not have to learn the diagram of the shell used to extract aluminium but you should be able to label the different parts. You should also be able to write half equations for the reactions at the electrodes. •It is a common mistake to think that the steel core in electricity cables just conducts electricity. It is also there to strengthen the cables. • A common mistake is to think that sulphate ions break up during the electrolysis of aqueous solutions into sulphur dioxide. In fact, oxygen is given off at the positive electrode (from the electrolysis of the water) • If the exam paper shows an electrical circuit to test conduction, observations can also include what can be seen to be happening in the circuit e.g. ‘the bulb lights up’ •It
is a common error to muddle cells with electrolysis. In electrolysis an electric current is used to decompose the electrolyte. In a cell the different reactivity of the electrodes makes an electric current flow. • You do not need to remember details about the construction of a fuel cell, but you may be asked questions based on diagrams and relevant half equations.
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6 Chemical energetics 6.1 Energetics of a reaction • Exothermic and endothermic reactions • Bond breaking is endothermic and bond forming is exothermic •ΔH : kJ/mol For an exothermic reaction, the enthalpy change is always negative.
in exothermic reactions the reactants are higher than the products For an endothermic reaction, the enthalpy change is always positive.
in endothermic reactions the reactants are lower than the products •Bond energy 6.2 Production of energy • Production of heat energy by burning fuels coal: very polluting – acid rain – global warming petroleum: less polluting – global warming natural gas: less polluting – global warming hydrogen: non polluting – lot of energy – explosive mixture •Calorimeter Using the ideas you learn in physics about specific heat capacity, you may have to calculate the amount of energy released by one mole of a substance. Heat evolved = m.c.ΔT
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Then calculate heat released per mole: Heat per mole = heat evolved / moles *ΔT is the temperature rise, m is the mass of the solution in grams which is assumed to equal its volume in cm3, c is the specific heat capacity of water which is 4.2 J K-1 g-1 Fair testing When comparing different fuels, it is important to carry out a fair test. Several variables should be kept constant. They include: the volume of water used the starting temperature of the water the temperature increase the distance of the flame from the calorimeter • ΔH=ΣΒbroken - ΣΒformed • Radioactive isotopes, such as 235U, as a source of energy • The production of electrical energy from simple cells. Electrochemical cell Zn/Cu in dilute H2SO4 ( –):Zn(s) Zn2+(aq) + 2e– (+): 2H+(aq) + 2e– H2(g) The more reactive Me is always the negative electrode Disadvantanges: -lose power(reactants used up) -bulky -have to be recharged -harmful -difficult to dispose of safety • Fuel cell: hydrogen (is bubbled through negative electrode) react with oxygen(is bubbled through positive electrode) to generate electricity. Acidic electrolyte ( H+ produced at the negative electrode and reacts at the positive) ( –):2H2(g) 4H+(aq) + 4e– (+):O2(g) +4H+(aq) + 4e– 2H2O(l) Alkaline electrolyte ( OH–reacts at the negative electrode and produced at the positive) ( –):2H2(g) +4OH–(aq) 4H2O(l) + 4e– (+):O2(g) +2H2O(l) + 4e– 4OH–(aq) Advantanges: no pollutants are formed – more energy /gr – lightweight – not recharging – high efficiency TIPS ON SPECIFIC TOPICS •Remember that burning is exothermic. • If asked whether a reaction is endothermic or exothermic, remember the following: endothermic – heat is put in (e.g. you have to heat with a Bunsen to get a reaction); exothermic – heat is given out (e.g. burning fuels and neutralisation reactions are always exothermic)
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7 Chemical reactions 7.1 Rate (speed) of reaction • Measuring rate of reaction mass of the reaction mixture volume of gas amount of light transmitted change the pH change in pressure time taken for a precipitate to make a letter disappear • Controlled variables – Independent variable • Calculating rate of reaction rate of reaction = Graf: near the start reaction is faster- then gets slower- finally reaction stops A reaction stops when the limiting reactant is completely used up
• The effect of particle size- surface area on the rate of reactions Increasing the surface area of a solid reactant increases the rate of reaction Smaller particles of solid have a larger surface area than larger ones with the same total volume Danger of explosive combustion with fine powders (e.g. flour mills) and gases (e.g. mines)
• The effect of catalysts - enzymes on the rate of reactions A catalyst speeds up the rate of a chemical reaction but is not used up itself We need tiny amounts of catalyst There are 2 types of catalyst: solid – in solution A solid catalyst works by allowing the reactants to get close together The reaction occurs more quickly at a lower temperature
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Activation energy Activation energy is the minimum energy needed for a reaction to occur when two particles collide. It can be represented on an energy level diagram.
• The effect of concentration on the rate of reactions: C ↑ rate ↑ Collision theory: Enough energy – number of successful collisions per second C ↑ frequency of collisions↑ rate ↑ Reactions involving gases : P↑ means C ↑
• The effect of temperature on the rate of reactions: T↑ rate ↑ Collision theory: T↑ E↑ frequency of collisions↑ rate ↑ Activation energy Ea T↑ more particles have E>Ea number of effective collisions ↑
rate ↑(more important)
the reactant particles move more quickly they have more energy the particles collide more often, and more of the collisions are successful the rate of reaction increases
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• Photochemical reactions ( the effect of light on the rate of reactions) • The use of silver salts in photography as a process of reduction of silver ions to silver 2Ag+Br- (crystal) + hv (radiation) 2Ag + Br2 2Ag+ + 2e- 2Ag : reduction 2BrBr2 + 2e- : oxidation • Photosynthesis the reaction between carbon dioxide and water in the presence of chlorophyll(catalyst) and sunlight to produce glucose and oxygen 6 CO2 + 6 H2O + Light
C6H12O6 + 6 O2
TIPS ON SPECIFIC TOPICS • Many students have difficulty explaining what is meant by rate of reaction. Remember two points: it is the change in volume or mass etc over a fixed period of time. Time is often omitted • Remember that the total volume of gas released by the same amount of metal is always the same. A common error is to think that powdered metal, when reacted with acid, gives off more gas than larger lumps of the same amount of metal • The total volume of gas released by a catalysed reaction is exactly the same as for an uncatalysed reaction. The same amount of reactants is the important factor • In rate questions, when asked to analyse graphs of volume of gas against time for the reaction of an acid with a metal or carbonate, a common error is to state the volume is increasing and not mention the rate. Remember that the rate is getting less and less with time because rate is the difference in volume divided by time •Remember that rate of reaction depends on two things: 1. the change in amount or concentration of reactants or products and 2.the time taken for this change to occur. •Make sure that you know how to interpret the different parts of a graph of volume of gas released or loss in mass of the reactants against time. For the Extension you should also be able to calculate the rate of reaction from these graph. •It is a common error to think that larger particles have a larger surface area than smaller ones. Think of a large cube cut up –by cutting, you are exposing more surfaces. •When defining a catalyst, the best answer is ‘a substance that speeds up a reaction but remains chemically unchanged at the end of the reaction’. Phrases such as ‘a substance which changes the rate of a reaction’ are rather vague. •When explaining the effect of concentration on reaction rate don’t just refer to more collisions between the particles. It is the more frequent collision of the particles which is important. • When writing answers to questions about rates of reaction, it is important to use words like faster or slower not just fast or slow. •Note that as temperature increases, each particle collides with a greater force. It is also more accurate to write that there are more frequent collisions than just more collisions. •It is important to realize that light only affects a few reactions. The only ones you have to know about are the photosynthesis, the conservation of silver bromine to silver and the reaction of alkanes with clorine.
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8 Reversible reactions Are ones that can go forward and backwards depending on the conditions Dehydration and Hydration: Assume we have a hydrated salt, copper sulphate for example. If you heat the salt you get two products. They are water and anhydrous copper sulphate. This is a reversible reaction because if you cool the mixture of the products again, you get hydrated copper sulphate back. CuSO4 . 5H2O⇋ CuSO4 + 5H2O →Heating→ ←Cooling← Note: hydrated copper sulphate is blue crystals. Anhydrous copper sulphate is white powder but it forms a blue solution with water. Equilibrium: Some reversible reactions are very unique, at a certain point, the reaction will be going forward and backwards at the same time and at the same rate. This is called the state of equilibrium. In the state of equilibrium, the rate of forward reaction is equal to the rate of backward reaction and the amount of products and reactants remain constant. Dynamic equilibrium 1. Rate of forward reaction = rate of reverse reaction 2. Concentrations of all reactants and products remain constant. 3. The system is closed, and on the large scale (macroscopic) everything is constant. • the effect of changing the concentration, on reversible reactions Increasing the concentration of a reactant moves the reaction in the direction of the products If we remove the products from an equilibrium mixture, more reactants are converted into products. If a catalyst is used, the reaction reaches equilibrium much sooner, because the catalyst speeds up the forward and reverse reactions by the same amount. • the effect of changing the temperature on reversible reactions If the temperature is increased, the position of equilibrium moves in the direction of the endothermic reaction if the temperature is reduced, the position of equilibrium moves in the direction of the exothermic reaction • the effect of changing the pressure on other reversible reactions If the pressure is increased, the position of equilibrium moves in the direction of the fewest moles of gas. TIPS ON SPECIFIC TOPICS • Make sure that you understand the term hydrated anhydrous and the water of crystallization • Remember that if the equilibrium conditions are changed the reaction always tries to act in the opposite direction
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• A common mistake is to say that in an equilibrium reaction, a catalyst increases the rate of the forward reaction more than the back reaction. One of the characteristics of equilibrium is that the backward and forward reactions go at the same speed. This applies to catalyzed as well as unanalyzed reactions
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9 Redox Oxidation is gain of oxygen. Reduction is loss of oxygen.
Oxidation is the loss of electrons from a substance. It is also the gain of oxygen by a substance Reduction is the gain of electrons by a substance. It is also the loss of oxygen from a substance. Usually, oxidation and reduction take place at the same time in a reaction. We call this type of reaction a redox reaction. Note that: the oxidising agent is the chemical that causes oxidation the reducing agent causes the other chemical to be reduced • In a redox reaction involving ions, tow half equations can be writen • Redox reactions by changes in oxidation state Assigning oxidation numbers Rule
Examples
1. The oxidation number of each atom in a pure element is zero.
Zn, O in O2, and P in P4 all have an oxidation number of zero.
2. The oxidation number of an atom in a monatomic ion is equal to the charge on the ion.
Na+ has an oxidation number of +1.
3. In compounds containing oxygen, each oxygen atom has an oxidation number of -2
In H2O and CO2 each oxygen atom has an oxidation number of -2.
4. In compounds containing hydrogen, each hydrogen atom has an oxidation number of +1
In NH3 and H2O each hydrogen atom has an oxidation number of +1.
5. For a molecule, the sum of the oxidation numbers of the atoms equals zero.
The sum of the oxidation numbers of the atoms in CH4 is zero. As such hydrogen atom has an oxidation number of +1, the oxidation number of the carbon atom is -4: (x + (4x + 1) = 0, x = -4).
6. For a polyatomic ion, the sum of the oxidation number of the atoms equals the charge in the ion.
The sum of the oxidation numbers of the atoms in PO43- is -3. As each oxygen atom has an oxidation number of -2, the oxidation number of the phosphorus atom is +5: (x + (4x – 2) = -3, x = +5).
7. In a compound, the most electronegative atom is assigned the negative oxidation number.
In SF6, the oxidation number of each fluorine atom is -1. The oxidation number of the sulfur atom is +6: (x + (6x – 1) = 0, x = +6).
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S2- has an oxidation number of -2.
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Oxidation is an increase in oxidation state Reduction is a reduction in oxidation state • redox reactions by the colour changes involved when using acidified potassium manganate(VII) in acidic solution is a good oxidant, when it oxidizes a substance is color change from purple to colourless potassium iodide in acidic solution is a good reductant, when it reduces a substance is color change from colourless to brown
TIPS ON SPECIFIC TOPICS • When explaining redox reactions, make sure you understand exactly what is being asked, especially if the question says ‘use the equation…’. Don’t just give a definition of redox in terms of electron loss or gain. If a question says ‘use the equation to explain why the iron oxide is reduced’, you must refer to the species in the equation in your answer, e.g. ‘the iron oxide loses its oxygen’. ‘Iron oxide gains electrons’ is incorrect
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10 Acids, bases, salts 10.1 The characteristic properties of acids and bases Acids are substances made of a hydrogen ion and non-metal ions. They have the following properties: They dissolve in water producing a hydrogen ion H +, They have a sour taste, Strong ones are corrosive, Their pH is less than 7. Turns blue litmus paper/ solution red All acids must be in aqueous form to be called an acid. For example Hydrochloric acid is hydrogen chloride gas dissolved in water. The most common acids are: Hydrochloric acid HCl, Sulphuric Acid H2SO4, Nitric Acid HNO3, Cirtric Acid, Carbonic Acid H2CO3. Dilute acids react with relatively reactive metals such as magnesium, aluminium, zinc and iron. The products of the reaction are a salt plus hydrogen gas. metal + acid → salt + hydrogen In general, the more reactive the metal, the faster the reaction. However, aluminium has a protective oxide layer, so it reacts slowly with acids to begin with. Acids react with metal oxides and hydroxides, a salt and water are made: acid + metal oxide → salt + water Acids react with carbonates, such as calcium carbonate (found in chalk, limestone and marble), a salt, water and carbon dioxide are made. In general: acid + metal carbonate → salt + water + carbon dioxide Bases are substances made of hydroxide OH- ions and a metal. Bases can be made of: Metal hydroxide (metal ion & OH- ion) Metal oxides Metal carbonates (metal ion & CO32-) Metal hydrogen carbonate (Bicarbonate) Ammonium hydroxide (NH4OH) Ammonium Carbonate ((NH4)2CO3) Properties of bases: Bitter taste Soapy feel Have pH’s above 7 Strong ones are corrosive Turns red litmus paper/ solution blue
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Some bases are water soluble and some bases are water insoluble. Water soluble bases are also called alkalis. Reactions of bases Alkalis react with acids to produce a salt and water (neutralization) e.g. NaOH(aq) + HCl(aq)
NaCl(aq) + H2O(l)
Metal oxides react with acids to produce a salt and water (neutralization) e.g. MgO(s) + 2HCl(aq)
MgCl2(aq) + H2O(l)
Metal carbonates react with acids to produce a salt, water and carbon dioxide e.g. Na2CO3(s) + 2HCl(aq)
2NaCl(aq) + H2O(l) + CO2(g)
Metal hydrogen carbonates react with acids to produce a salt, water and carbon dioxide e.g. NaHCO3(s) + HCl(aq)
NaCl(aq) + H2O(l) + CO2(g)
Displacement of ammonia from ammonium salts NH4Cl(s) + NaOH(aq)
NaCl(aq) + H2O(l) + NH3(g)/(aq)
Ammonia reacts with acids to produce an ammonium salt e.g. NH3(aq) + HCl(aq)
NH4Cl(aq)
• Neutrality and relative acidity and alkalinity in terms of pH - Measured using Universal Indicator paper
Controlling Soil pH: If the pH of the soil goes below or above 7, it has to be neutralized using an acid or a base. If the pH of the soil goes below 7, calcium carbonate (lime stone) or calcium oxide (lime) is used to neutralize it. The pH of the soil can be measured by taking a sample from the soil, crushing it, dissolving in water then measuring the pH of the solution. Acids and bases in terms of proton transfer acid is a hydrogen ion (proton) donor. base is a hydrogen ion (proton) acceptor
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Strong acid: an acid that ionizes completely in aqueous solution. e.g. HCl, HNO3, H2SO4 Weak acid: an acid that ionizes to a small extent (partially)in aqueous solution. Strong base: a base that almost completely dissociated in aqueous solution , are group 1 hydroxides (ie NaOH etc), or lower group 2 hydroxides Ba(OH) 2. e.g. NaOH, KOH, Ba(OH)2 Weak base: a base that accepts a hydrogen ion from water with difficulty. Distinguish between equimolar solutions of strong and weak acid. Strong acid: has a higher conductivity, better electrical conductor react more rapidly with magnesium results to a greater increase in temperature during the neutralization has lower pH, higher concentration of H+
10.2 Types of oxides Acidic Oxides They are all non-metal oxides except non-metal monoxides They are gases They react with an alkali to form salt and water Note: metal monoxides are neutral oxides Examples: CO2, NO2, SO2 (acidic oxides) & CO, NO,H2O (neutral oxides) Basic Oxides They are metal oxides They react with acids forming a salt and water They are solids They are insoluble in water except group 1 metal oxides. They react with an acid forming salt and water Examples: Na2O, CaO and CuO Amphoteric Oxides These are oxides of Aluminum, Zinc & Lead They act as an acid when reacting with an alkali & vice versa Their element’s hydroxides are amphoteric too They produce salt and water when reacting with an acid or an alkali. Al2O3(s) + 6HCl(aq)
2AlCl3(aq) + 3H2O(l)
Al2O3(s) + 2NaOH
2NaAlO2(aq) + H2O(l)
ZnO(s) + 2HCl(aq)
ZnCl2(aq) + 2H2O(l)
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ZnO(s) + 2NaOH
Na2ZnO2(aq) + H2O(l)
Neutral Oxides These are N2O, NO,CO They do not act as an acid or base 10.3 Preparation of salts Soluble
Insoluble
All nitrates
None
All common sodium, potassium and ammonium salts
None
Most common sulfates
Calcium , Barium and Lead
Most common chlorides, bromides, iodides
Silver , Lead
Sodium, Potassium and Ammonium
Most common carbonates
Group I and Ammonium, (Calcium is slightly soluble)
Hydroxides
Group I and Group II react with water
Most metal oxides
Preparing Soluble Salts: Displacement Method (Excess Metal Method): Metal + Acid → Salt + Hydrogen Note: this type of method is suitable to for making salts of moderately reactive metals because highly reactive metals like K, Na and Ca will cause an explosion. This method is used with the MAZIT (Magnesium, Aluminum, Zinc, Iron and Tin) metals only. Example: set up an experiment to obtain magnesium chloride salt. Mg + 2HCl → MgCl2 + H2 Observations of this type of reactions: Bubbles of colorless gas evolve (hydrogen). To test approach a lighted splint if hydrogen is present it makes a pop sound The temperature rises (exothermic reaction) The metal disappears You know the reaction is over when: No more gas evolves No more magnesium can dissolve The temperature stops rising The solution becomes neutral Proton Donor and Acceptor Theory: When an acid and a base react, water is formed. The acid gives away an H + ion and the base accepts it to form water by bonding it with the OH- ion. A hydrogen ion is also called a proton this is why an acid can be called Proton Donor and a base can be called Proton Acceptor.
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Neutralization Method: Acis + Base → Salt + Water Note: This method is used to make salts of metals below hydrogen in the reactivity series. If the base is a metal oxide or metal hydroxide, the products will be salt and water only. If the base is a metal carbonate, the products will be salt, water and carbon dioxide. Type 1: Acid + Metal Oxide → Salt + Water To obtain copper sulfate salt given copper oxide and sulfuric acid: CuO + H2SO4 → CuSO4 + H2O Observations of this reaction: The amount of copper oxide decreases The solution changes color from colorless to blue The temperature rises You know the reaction is over when No more copper oxide dissolves The temperature stops rising The solution become neutral Type 2: Acid + Metal Hydroxide → Salt + Water to obtain sodium chloride crystals given sodium hydroxide and hydrochloric acid: HCl + NaOH → NaCl + H2O Observations: Sodium hydroxide starts disappearing Temperature rises You know the reaction is over when: The temperature stops rising No more sodium hydroxide can dissolve The pH of the solution becomes neutral Type 3: Acid + Metal Carbonate → Salt + Water + Carbon Dioxide To obtain copper sulfate salt given copper carbonate and sulfuric acid: CuCO3 + H2SO4 → CuSO4 + H2O + CO2 Observations: Bubbles of colorless gas (carbon dioxide) evolve, test by approaching lighted splint, if the CO2 is present the flame will be put off Green Copper carbonate starts to disappear The temperature rises The solution turns blue
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You know the reaction is finished when: No more bubbles are evolving The temperature stops rising No more copper carbonate can dissolve The pH of the solution becomes neutral Titration Method: This is a method to make a neutralization reaction between a base and an acid producing a salt without any excess. is used to make a soluble salt the experiment is preformed twice, the first time, using an indicator ,is to find the amounts of reactants to use, and the second experiment is the actual one.
Other indicators Indicator
Acidic
Neutral
Alkaline
Methyl orange
Red
Yellow
Yellow
Phenolphthalein
Colourless
Colourless
Pink
Preparing Insoluble Salts: Precipitation Method:A precipitation reaction is a reaction between two soluble salts. The products of a precipitation reaction are two other salts, one of them is soluble and one is insoluble (precipitate). Example: To obtain barium sulfate salt given barium chloride and sodium sulfate: BaCl2 + Na2SO4 → BaSO4 + 2NaCl Ionic Equation: Ba2+ + SO42- → BaSO4 Observations: Temperature increases An insoluble solid precipitate (Barium sulfate) forms
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You know the reaction is over when: The temperature stops rising No more precipitate is being formed • Suggest a method of making a given salt from suitable starting material, given appropriate information
10.4 Identification of ions and gases Colors of Salts: Salt Hydrated copper sulfate Anhydrous copper sulfate Copper nitrate Copper chloride Copper carbonate Copper oxide Iron(II) salts
Formula CuSO4.5H2O
Solid Blue crystals
In Solution Blue
CuSO4
White powder
Blue
Cu(NO3)2 CuCl2 CuCO3 CuO E.g.: FeSO4, Fe(NO3)2
Blue Green Insoluble Insoluble Pale green
Iron(III) salts
E.g.: Fe(NO3)3
Blue crystals Green Green Black Pale green crystals Reddish brown
Reddish brown
Tests for Gases: Gas Ammonia Carbon dioxide Oxygen Hydrogen Chlorine Nitrogen dioxide Sulfur dioxide
Formula NH3 CO2
Tests Turns damp red litmus paper blue Turns limewater milky
O2 H2 Cl2 NO2
Relights a glowing splint ‘Pops’ with a lighted splint Bleaches damp litmus paper Turns damp blue litmus paper red
SO2
Turns acidified aqueous potassium dichromate(VI) from orange to green
Tests for Anions: Anion Carbonate (CO32-) Chloride (Cl-)(in solution) Iodide (I-)(in solution) Nitrate (NO3-)(in solution) Sulfate (SO42-)
Test Add dilute acid Acidify with dilute nitric acid, then add aqueous silver nitrate Acidify with dilute nitric acid, then add aqueous silver nitrate Add aqueous sodium hydroxide, then aluminium foil; warm carefully Acidify, then add aqueous barium nitrate
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Result Effervescence, carbon dioxide produced White ppt. Yellow ppt. Ammonia produced
White ppt.
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Tests for aqueous cations: Cation Aluminium (Al3+) Ammonium (NH4+) Calcium (Ca2+) Copper (Cu2+) Iron(II) (Fe2+) Iron(III) (Fe3+) Zinc (Zn2+)
Effect of aqueous sodium hydroxide White ppt., soluble in excess giving a colourless solution Ammonia produced on warming White ppt., insoluble in excess Light blue ppt., insoluble in excess Green ppt., insoluble in excess Red-brown ppt., insoluble in excess White ppt., soluble in excess, giving a colourless solution
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Effect of aqueous ammonia White ppt., insoluble in excess – No ppt. or very slight white ppt. Light blue ppt., soluble in excess, giving a dark blue solution Green ppt., insoluble in excess Red-brown ppt., insoluble in excess White ppt., soluble in excess, giving a colourless solution
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TIPS ON SPECIFIC TOPICS • Don’t confuse the pH scale with the degree of acidity. The more acidic the substance, the lower the pH – learn this by remembering that ‘a’ (for acid) is the lowest numbered letter of the alphabet • A common error is to think that less sodium hydroxide is needed to neutralise a weak acid than to neutralise a strong acid of the same concentration. The same amount is needed because the hydroxide is reacting with all the acidic hydrogens in the molecule, not just those that have ionised
• The phrase ‘explain why this acid is acting as a base’ demands a chemical reason (usually based on particle theory). The examiner is looking for an answer involving proton transfer. Vague answers (such as ‘it is neutralising the base’) are not accepted as they do not give an explanation • Simple inorganic salts such as sodium chloride are generally neutral when dissolved in water – they are not acidic • Nitric acid is a strong, not a weak, acid • A common error is to think that calcium hydroxide is insoluble in water. Remember that limewater is a solution of calcium hydroxide, so it must at least be slightly soluble • If you are asked to explain what the symbol (aq) means, write down more than ‘aqueous’. An answer such as ‘dissolved in water’ is needed • Look out for phrases such as ‘chemical test’ or ‘physical test’ – don’t just focus on the word ‘test’. For example, a chemical test for water could be ‘turns anhydrous copper sulphate blue’ (the word ‘anhydrous’ is essential). A physical test for water could be ‘a boiling point of 100oC’, using the correct units
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• When testing hydrogen chloride gas with litmus paper, many students think that the litmus paper is bleached first and then goes red. Remember that chlorine does this, not hydrogen chloride • The tests for ammonium and nitrate ions are commonly confused. Both require heating with sodium hydroxide, but to test for nitrate you need to add aluminium, as you need to remove the oxygen (reduce the nitrate) to make the ammonia. You don’t need to do this for the ammonium ion as it has no oxygen • Tests for aluminium ions and zinc ions are also often confused. Remember PANDA (precipitate of aluminium (hydroxide) does not dissolve in ammonia). Both zinc and aluminium ions form a white precipitate with sodium hydroxide, which re-dissolves in excess, but in ammonia only the zinc precipitate re-dissolves • Questions involving the height of precipitates when sodium hydroxide is added to a solution of metal ions often cause problems. Remember, as you add more hydroxide to a solution of suitable metal ions (e.g. iron(II) ions) there will be more precipitate until all the metal ions are used up. However, with excess sodium hydroxide, some hydroxides re-dissolve e.g. aluminium hydroxide. In these cases the height of the precipitate will then decrease as you add more hydroxide •Remember that a lower acidity gives a higher pH and a higher acidity gives a low pH. •Don’t forget that when acids react with carbonates, water is produced – as well as a salt and carbon dioxide. For extension you must be able to write the symbol equations. •It is incorrect to use the word ‘strong’ and ‘weak’ when referring to the concentration of acid or alkalis. Use ‘concentrated’ or ‘dilute’. Strong and weak refer to the degree of ionization of the acid or base, not the concentration. •When you make a salt using excess metal or metal oxide, you first have to filter off the excess solid reactant. You may be asked how to make a salt in any of the exam papers. •Make sure that you know what types of compound are soluble or insoluble. Without this knowledge you will not be able to select precipitation as the correct method to make a particular salt. •A common error is to confuse the tests for hydrogen and oxygen. It may help you to remember that ‘lighted’ (splint) has an ‘h’ in it for hydrogen and ‘glowing’ (splint) has an ‘o’ in it for oxygen. •When testing for metal ions using sodium hydroxide, make sure that you mention three things: (i) if there is a precipitate (ii) the colour of the precipitate (iii) what happens when you add excess sodium hydroxide •Remember that you add nitric acid and silver nitrate in the test for halide ions. If you add hydrochloric acid you will be adding chloride ions!
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11 Periodic table • The Periodic Table as a method of classifying elements and its use to predict properties of elements Periodic trends • the change from metallic to non-metallic character across a period • the relationship between Group number, number of valency electrons and metallic/nonmetallic character Special Elements: Alkali Metals: These elements lie in group 1 of the periodic table. They are Lithium, Sodium, Potassium, Rubidium, Caesium and Francium (radioactive). We will study the properties of the first three; Lithium, Sodium and Potassium. Like any metals they are all good conductors of heat and electricity. They are however, soft. Lithium is the hardest of them and potassium is the softest. They are extremely reactive; they have to be stored away from any air or water. They have low densities and melting points. They react with oxygen or air forming a metal oxide: 4Li +O2 → 2Li2O Their oxides can dissolve in water forming an alkaline solution of the metal hydroxide: + H2O
→ 2LiOH
(Lithium Oxide)
(Water)
Li2O
(Lithium Hydroxide)
They react with water vigorously forming metal hydroxide and hydrogen gas: 2K
+ 2H2O
→ 2KOH
+ H2
They React with Halogens forming a metal halide: 2Na + Cl2
→ 2NaCl
The reactivity of Group 1 elements increases as you go down the group because: the atoms get larger as you go down the group the outer electron gets further from the nucleus as you go down the group the attraction between the nucleus and outer electron gets weaker as you go down the group - so the electron is more easily lost
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Flame colors and the alkali metal ion they represent Flame colour
Ion present
Red
Lithium, Li+
Orange
Sodium, Na+
Lilac
Potassium , K+
Brick red
Calcium, Ca2+
The Halogens: These are elements of group 7; Fluorine, Chlorine, Bromine, Iodine and Astatine. We will study only properties of chlorine, bromine & iodine. They are colored and the color gets darker as we go down the group. They exist as diatomic molecules (Cl 2, Br2, I2). As you go down, they gradually change from gas to solid (chlorine is gas, bromine is liquid and iodine is solid). They react with hydrogen forming hydrogen halide, which is an acid if dissolved in water: H2
+
→
Cl2
(Hydrogen)
2HCl
(Chlorine)
(Hydrochloric Acid)
They react with metals forming metal halide: 2Fe +
→
3Cl2
2FeCl3
The reactivity also decreases as we do down, chlorine is most reactive, followed by bromine then iodine. If you bubble chlorine gas through a solution of potassium bromide, chlorine will take bromine’s place because it more reactive. This is a displacement reaction. 2KBr
+
Cl2
→
2KCl
+
Br2
Transition Elements: These are metals. They form a big part of the periodic table. Some of them are very common like copper, zinc and iron. They have the following properties: They are harder and stronger than metals of groups 1 & 2. They have much higher densities than metals other metals. They have high melting points except for mercury. They are less reactive than metals of group 1 & 2. Excellent conductors of heat and electricity. They show catalytic activity (act as catalysts) They react slowly with oxygen and water They form simple ions with several oxidation states and complicated ions with high oxidation states. They form coloured compounds
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Noble Gases: These are elements in group 8 of the periodic table. They are colorless gases. They are extremely unreactive; this is because they have their outer energy shell full with electrons. So they are stable, this is why they exist as single atoms. Noble gas
Uses
Helium
Party balloons, airships, cooling superconducting electromagnets (eg in MRI scanners), gas for scuba diving
Neon
Red neon signs, lasers
Argon
Shielding gas for welding, surrounding the filament in an old-fashioned lightbulb
Xenon
Lights, lasers
Krypton
Lights, photographic flashguns
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12 Metals 12.1 Properties of metals
Metallic bonds
Properties of Metals Good conductors of electricity : metals have delocalized electrons / sea of electrons which are mobile. Good conductors of heat : electrons jumping through cations and moving energy from M to M . Shiny : light absorbed by electrons and re-emitted at different Energy levels. Malleable : pushing layers, atoms/ions/layers (of positive ions) can slide over each other without change in the bonding forces / Ductile : moving the layers. (impurities – alloys: harder than the pure metals) Melting and boiling point : high • Alloys An alloy is a mixture of two or more elements, where at least one element is a metal. Most alloys are mixtures of two or more metals Alloys contain atoms of different sizes. These different sizes distort the regular arrangements of atoms. This makes it more difficult for the layers to slide over each other, so alloys are harder than the pure metal.
It is more difficult for layers of atoms to slide over each other in alloys
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12.2 Reactivity series Reactions with Dilute Hydrochloric Acid: Metal + HCl → Metal Chloride + Hydrogen Metals Potassium, Sodium & Calcium Magnesium & Aluminum Zinc, Iron & Lead Copper, Silver, Gold & Platinum
Reactivity with Dilute HCl React extremely violently with rapid effervescence and splashing React violently with rapid effervescence React slowly with bubbles Do not react
Reactions with Oxygen in Air: Most metals react with oxygen from air forming a metal oxide. Metal potassium, sodium, calcium and magnesium aluminum and zinc
React with oxygen with a very bright flame
Iron and copper
very slowly
copper lump
silver, gold and platinum
Product white ashes and their oxides are soluble. white powdered ashes but their oxides are insoluble. A layer of aluminum oxide adheres and covers the aluminum. At this point no further reaction can take place. rust which is reddish brown iron oxide - insoluble a white layer of black copper oxide forms on it. When the lump gets covered by this layer; the reaction stops- insoluble
do not react
Reactions of Metals with Water and Steam: Potassium, sodium and calcium react vigorously with cold water and may catch on fire. The products of these reactions are metal hydroxide and hydrogen gas. If hydrogen gas being produced accumulates it may ignite and cause an explosion. Metal + Water → Metal hydroxide + Hydrogen E.g.: 2Na + 2H2O → 2NaOH + H2 Magnesium, aluminum, zinc and iron are less reactive. They react with steam forming metal oxide and hydrogen. Magnesium and aluminum will react vigorously with steam while zinc and iron react slowly. Metal + Steam → Metal Oxide + Hydrogen E.g.: Magnesium + Steam → Magnesium oxide + Hydrogen Unreactive metals such as silver and gold do not react with water.
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Single Displacement Reactions in Solid State: Previously you’ve studied displacement reactions which are pre-formed in aqueous states. A very similar reaction takes place in the solid state, it is called thermite reaction. This reaction is used to repair damaged railway lines. In this reaction, aluminum and iron (III) oxide are the reactants. In the reaction, aluminum removes the oxygen ion from iron and bonds with it. This happens because aluminum is more reactive than iron. The products are aluminum oxide and iron in molten form. In the fixing procedure, the reactants are put in the cut in the railway line and the reaction is triggered by heating using a magnesium fuse. The reaction leaves aluminum oxide and molten iron with then condenses in the cut welding it. Like displacement reactions, this reaction is exothermic. 2Al + Fe2O3 → Al2O3 +2Fe Single Displacement Reactions in Aqueous State: These are ordinary displacement reactions in which the two positive ions compete for the negative ion. The ion of the more reactive metal wins. Zinc is higher than copper in the reactivity series. If zinc is added to a solution of copper nitrate, a displacement reaction will take place in which the zinc will displace the copper ion from the solution in its salt. The products of this reaction are zinc nitrate and copper. Copper salt solutions have a blue color which fades away as the reaction proceeds because the concentration of the copper salt decreases. This type of reaction also helped in confirming reactivity of metals since the more reactive metal displaces the less reactive one. Zn + Cu(NO3)2 → Zn(NO3)2 + Cu Explaining reactivity The easy with which a metal loses its valency electrons depends on the distance of the valency electrons from the nucleus the nuclear charge (number of protons) the number of electrons shells Reducing metal oxides with carbon Metal oxides below C in the reactivity series are reduced by carbon when heated Action of Heat on Metal Compounds: Applying heat to a metal compound such as potassium nitrate will cause it to decompose into potassium nitrite and oxygen. This is a thermal decomposition reaction. Metal: Potassium Sodium Calcium Magnesium Aluminum Zinc Iron Lead Copper Silver Gold
Nitrate (NO3) Metal Nitrate → Metal nitrite + Oxygen Metal Nitrate → Metal oxide + Nitrogen dioxide + Oxygen
Metal Nitrate → Metal + Nitrogen dioxide + Oxygen
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Anion: Carbonate (CO3) Hydroxide (OH) NO DECOMPOSITION Metal Carbonate → Metal oxide + Carbon dioxide
Metal Carbonate → Metal + Carbon dioxide + Oxygen
Metal hydroxide → Metal oxide + Hydrogen
Silver and gold hydroxides do not exist.
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Ions of more reactive metals tend to hold on tightly to their anions and do not decompose easily this is why lots of heat is needed. 12.3 Extraction and Uses of metals Metal
Method
Potassium
Electrolysis
Sodium
Electrolysis
Calcium
Electrolysis
Magnesium
Electrolysis
Aluminium
Electrolysis
(Carbon)
(Non-metal)
Zinc
Reduction by carbon or carbon monoxide
Iron
Reduction by carbon or carbon monoxide
Tin
Reduction by carbon or carbon monoxide
Lead
Reduction by carbon or carbon monoxide
(Hydrogen)
(Non-metal)
Copper
Various chemical reactions
Silver
Various chemical reactions
Gold
Various chemical reactions
Platinum
Various chemical reactions
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Extraction of Iron: The ore of iron is called hematite. It consists of 60% iron in form of Iron oxide (Fe 2O3) with other impurities such as silicon oxide (SiO2). This process takes place in a tower called a Blast furnace.
Substances Iron ore (Hematite) Coke (heated coal) Lime stone (Calcium carbonate) Hot Air
Products and Waste Materials Pure Iron Carbon dioxide Air Slag (Calcium silicate)
Substances are put in the blast furnace The process starts by blowing in hot air at the bottom of the furnace Coke burns in oxygen from the hot air producing carbon dioxide; C + O2 → CO2 Heat makes lime stone decompose into calcium oxide and carbon dioxide; CaCO3 → CaO + CO2 Carbon dioxide produced goes up the furnace and reacts with more coke up there producing carbon monoxide; CO2 + C → 2CO Carbon monoxide is a reducing agent. It rises further up the furnace where it meets iron oxide and starts reducing it producing iron and carbon dioxide; Fe2O3 + 3CO → 2Fe + 3CO2 Calcium oxide which was produced from the thermal decomposition of lime stone is a base. It reacts with impurities of hematite such as silicon oxide which is acidic forming calcium silicate which is called slag; CaO + SiO2 → CaSiO3 Molten Iron and slag produced trickles down and settles at the bottom of the furnace. Iron is denser than slag so it settles beneath it. Iron and slag are tapped off separately at regular intervals and pure iron is collected alone Waste gases such as carbon dioxide formed in the process and nitrogen and other gases from air blown in escape at the top of the furnace. Conversion of Iron into Steel: Iron produced in the blast furnace is called pig iron. It contains 4% carbon as well as other impurities such as sulfur, silicon and phosphorus which make it hard and brittle. It got that name from the fact that it has to be poured into mould called pigs before it is converted into steel. Most of produced iron is converted into steel because steel has better properties. If all the impurities are removed, the iron becomes very soft In this condition, it easily shaped but is too soft for many uses. Pure iron also rust very easily. Making steel out of pig iron is a process done in a basic oxygen furnace: Molten pig iron is poured into the oxygen furnace A water cooled lance is introduced which blows oxygen onto the surface of the molten iron Impurities start to react Carbon is oxidized into carbon monoxide and carbon dioxide and escape Sulfur is oxidized into sulfur dioxide and escapes Silicon and phosphorus are oxidized into silicon oxide and phosphorus pentoxide which are solids. Calcium oxide (lime) is added to remove the solid impurities as slag which is skimmed off the surface Throughout the process, sample of the iron are being taken and analyzed for the percentage of carbon present in it. When the percentage of carbon desired is reached, the furnace is switched off and the steel is collected.
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There are many different forms of steel. Steel Mild Steel
Composition 99.5% Iron 0.5% Carbon
Properties Easily worked lost brittleness
Hard Steel
99% Iron 1% Carbon 87% Iron 13% Manganese
Tough and brittle
Manganese Steel
74% Iron 18% Chromium 8% Nickel
Tough and resistant to corrosion
Tungsten Steel
95% Iron 5% Tungsten
Tough and hard even at high temperatures
Stainless Steel
Tough and springy
Uses Car bodies large structures Machinery Cutting tools and chisels Drill bits and springs and chemical plants Cutlery and surgical tools, kitchen sinks Edges of high speed cutting tools
Extraction of Zinc: The ore of zinc is called zinc blende and it is made of zinc sulfide. Zinc is obtained from zinc sulfide by converting it into zinc oxide then reducing it using coke, but first zinc sulfide must be concentrated. Zinc sulfide from zinc blende is concentrated by a process called froth floatation. In this process, the ore is crushed and put into tanks of water containing a frothing agent which makes the mixture froth up. Hot air is blown in and froth starts to form. Rock impurities in the ore get soaked and sink to the bottom of the tank. Zinc sulfide particles cannot be soaked by water; they are lifted by the bubbles of air up with the froth and are then skimmed off. This is now concentrated zinc sulfide. Then, zinc sulfide gets heated very strongly with hot air in a furnace. Zinc sulfide reacts with oxygen from the air to produce zinc oxide and sulfur dioxide gas which escapes as waste gas. 2ZnS + 3O2 → 2ZnO + 2SO2 Sulfur dioxide is used in the manufacture of sulfuric acid. Zinc oxide produced is put into a furnace with powdered coke. The mixture is heated till 1400oC. Carbon from the coke reduces the zinc oxide into zinc producing carbon monoxide which escapes as waste gas. ZnO + C → Zn + CO
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Carbon monoxide produced is hot and is used to heat the furnace to reduce heating costs. The pure zinc produced is collected and left to cool down. Zinc is used in many ways like the production of the alloy brass, galvanization and making car batteries. Uses of Zinc: for galvanising and for making brass
Extraction of Aluminum: Aluminum exists naturally as aluminum oxide (alumina) in its ore, which is called bauxite. Because aluminum is a very reactive metal, it holds on very tightly to the anion it bonds with, which is oxide in this case. This is why the best way to extract and purify aluminum is by electrolysis in a cell like the one below.
In this cell, the electrodes are made of graphite (Carbon). The cathode is a layer at the bottom of the cell and the anodes are bars dipped in the electrolyte. The electrolyte in this process is a molten mixture of aluminum oxide and cryolite. Aluminum oxide by its self has a very high melting point of 2050oC which is higher than the melting point of the steel container in which this process is done. That means the steel container will melt before the aluminum oxide. This is why aluminum oxide is mixed with cryolite which decreases the melting point of it to under 1000oC, thus saving a lot of money because heating is expensive and preventing the steel container from melting. Heat must be continuously supplied to the mixture to keep it molten. Aluminum oxide does not conduct electricity when solid because it does not have free mobile ions to carry the charge. Aluminum oxide is purified from impurities of oxide by adding sodium hydroxide Aluminum oxide is mixed with cryolite and put in the electrolysis cell Heat is given in until the mixture becomes molten Electrolysis start Oxide ions get attracted to the anode and discharged (oxidation); 2O2-, 4e → O2 Aluminum ions get attracted to the cathode and discharged and settle at the bottom of the container (reduction); Al3+ + 3e → Al Oxygen gas evolves and is collected with waste gases Aluminum is sucked out of the container at regular intervals Oxygen gas which evolves reacts with carbon from the cathode forming CO2. The cathode gets worn away. To solve this, the cathode is replaced at regular intervals. Heat supply is very expensive; this is why cryolite is used to decrease the melting point of aluminum oxide and this process is done in plants which use hydroelectric energy because it is cheap.
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Uses of aluminum: Construction of air-craft bodies because aluminum is very strong and very light and it is resistant to corrosion Food containers because it is resistant to corrosion Overhead power cables because it conducts electricity, is very light, malleable and ductile. Although it is strengthened with steel core Extraction of Copper: Copper is one of the most popular metals. Native copper occurs in some regions in the world. Otherwise, copper exists in its ore, copper pyrites (2CuFeS2). You have studied before that copper can be purified by electrolysis. It can also be extracted from it ore by converting pyrites into copper sulfide by reacting it with oxygen: 2CuFeS2 + 4O2 → Cu2S + 3SO2 + 2FeO Sulfur oxide produced escapes as waste gas and iron oxide impurities are removed by heating the mixture with silicon converting it in to iron silicate which is run off. The remaining copper sulfide is then heated strongly with air. Copper sulfide reacts with oxygen from air producing sulfur oxide which escapes as waste gas and pure copper. Cu2S + O2 → 2Cu + SO2 Thus copper is extracted. Uses of Copper: In electrical wires because it is a perfect electrical conductor and very ductile, malleable and cheap Making alloys such as bronze and brass Cooking utensils because it conducts heat and it is has high melting and boiling points and also resists corrosion Electrodes because it is a good conductor of electricity Water pipes because it is resistant to corrosion TIPS ON SPECIFIC TOPICS • Don’t confuse the properties of elements with those of their compounds (especially when they appear in the same question). For example, if asked about the properties of the element oxygen, don’t give the properties of an oxide • The properties of transition elements often cause problems. Remember that transition elements themselves are NOT coloured, it is their compounds that are coloured • When trying to distinguish between a transition metal and a non-transition metal, information on boiling points is more important than information on density. Some non-transition elements (such as lead) are very dense • If asked about the specific properties of transition metals, don’t list general properties of metals, such as ‘shiny’, ‘malleable’, etc. • In questions about sacrificial protection, remember that the more reactive metal of the pair will corrode. To answer this sort of question, know the order of common metals in the reactivity series
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• ‘Corrosive’ and ‘corrosion’ are often confused. ‘Corrosive’ means that a chemical ‘eats away’ another substance – acids and alkalis are corrosive. ‘Corrosion’ is the process of ‘eating away’. A statement such as ‘iron is corrosive’ is therefore incorrect • The source of an element is where it is found (i.e. a particular place or in a particular substance) – a source of sulphur is the southern USA, or petrol. It does not mean the process of extraction. Don’t write vague statements such as ‘underground’ • Sulphur dioxide is not used ‘to make wood pulp’, it is used to bleach wood pulp •You need to know where the metals and the non-metals appear in the Periodic Table. You do not have to remember exactly the dividing line between metals and non-metals in Groups III to VI. •When you describe observations concentrate on what you see, hear smell or feel by touch. •When you compare Group I metals, remember that they have ‘similar properties’ NOT ‘the same properties’. The properties change slightly down the group. •Make sure that you can distinguish between the halogens (elements) and halides (compounds). It is a common error to write chlorine ions instead of chloride ions. •It is better to write that the noble gases are unreactive ‘because they have a full outer shell of electrons’, which is inaccurate. •It is a common error to suggest that transition elements are highly coloured. It is the compounds of transition elements which have a range of colours. •Remember that oxidation state does not always refer to the charge on the ions. For example, in potassium manganate (VII), KMnO4 , the oxidation state of manganese is +7 but the manganese ion with the highest charge is Mn+2. • It is a common error to think that all metals are hard and have very high melting points. Remember that Group I metals are soft and have low melting points. •Remember that metals that react with cold water form metal hydroxides. When a metal is heated in steam, an oxide is formed. •In your exam you will usually be given the reactivity series to help you answer questions about the ease of formation of ions. •Remember that aluminium is a reactive metal. It must be reactive if it forms an unreactive oxide layer on its surface so quickly. •You need to remember the products from the thermal decomposition of nitrates. If you don’t know these, you won’t be able to write equations for thermal decomposition. •You will not be asked to draw the furnace used for the extraction of zinc but you should be prepared to label a diagram and write relevant equations. • You will not be asked to draw the blast furnace. You should be prepared to answer questions related to a diagram of the blast furnace and the reactions involved. •Do not confuse steelmaking with the blast furnace. In steelmaking the impurities are removed from the impure iron we get from the blast furnace. In the blast furnace the impure iron is extracted from the iron ore.
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13 Air and water Chemical tests for water Pure copper(II) sulfate is white. It is also known as anhydrous copper(II) sulfate because it has no water in it. When water is present in a sample of copper(II) sulfate it turns blue. It is still a dry solid, because the individual water molecules are trapped within the ionic lattice surrounding the copper(II) ions.
Solutions of copper(II) sulfate are also blue. Water can also be detected using blue anhydrous cobalt(II) chloride. This turns pink in the presence of water.
Uses of Water: The uses of water are many, from drinking and cleaning to irrigating crops and landscapes. Water is used for cooling, for recreation, and dust control. Water is needed for restaurants, most industrial processes, and even some religious ceremonies. On another level, the splash and flow of water in streams and fountains soothes and inspires. In one way or another, water is a part of almost everything humans make and do. Washing a load of laundry uses 40 gallons, filling a backyard pool takes about 25,000 gallons, growing a pound of cotton consumes 1,000 gallons, while producing a pound of copper uses 20 gallons. Uses where water is consumed, usually through evaporation or plant growth, are consumptive uses. Examples include water used for irrigation or in evaporative coolers. Non-consumptive uses, such as bathing, hydropower generation and recreation, do no t use up water. Used nonconsumptively, the same water can be used again and again, although some uses lower the quality of the water. Once used, wastewater can be treated and used again as reclaimed water or effluent.
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The main categories of water use are agricultural, municipal and industrial. Municipal and industrial uses currently are much less, but are growing rapidly. Mining activities and cooling towers used for power generation account for most of the remaining water use. Water Purification: 1.Water that exists naturally in earth is never pure. There are always impurities in it, sometimes in large amounts. In fact water could very well be contaminated with diseases and bacteria. This is why water has to be purified before it is put to use. Water purification involves two processes (Filtration & Chlorination) done in several steps: 2.Water is taken from reservoirs or any other source to the water treatment plant 3.Water is passed through filters to remove large, floating objects such as pieces of rocks or mud 4.Smaller particles are removed by adding aluminum sulfate which makes them stick together in large pieces and settle down 5.Water is passed through sand and gravel filters which filter off small particles and may kill some bacteria (filtration is done) 6.Chlorine gas is bubbled through the water to kill all bacteria living in the water making the water sterile 7.The water may end to be slightly acidic, small amounts of sodium hydroxide are added to treat this. Fluoride might be added to because it helps in preventing tooth decay 8.Water is then delivered to homes Composition of clean air
Fractional distillation of air About 78 per cent of the air is nitrogen and 21 per cent is oxygen. These two gases can be separated by fractional distillation of liquid air. Liquefying the air Air is filtered to remove dust, and then cooled in stages until it reaches –200°C. At this temperature it is a liquid. The air has been liquefied. Here's what happens as the air liquefies: water vapour condenses, and is removed using absorbent filters carbon dioxide freezes at -79°C, and is removed oxygen liquefies at -183°C nitrogen liquefies at -196°C The liquid nitrogen and oxygen are then separated by fractional distillation. The liquefied air is passed into the bottom of a fractionating column. Just as in the columns used to separate oil fractions, the column is warmer at the bottom than it is at the top.
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Fractional distillation
Air Pollution: Pollution is the presence of harmful substances. Air pollution is the presence of pollutant gases in the air. A pollutant is a substance that causes pollution. These are: Carbon monoxide Oxides of nitrogen Sulphur dioxide Lead compounds Carbon Monoxide: Carbon monoxide (CO) is one of the poisonous pollutants of air. It is considered a pollutant because it can kill living organisms. The main source of carbon monoxide is factories which burn carbon-containing fossil fuels since CO is one of the products of the incomplete combustion of fossil fuels. Carbon monoxide could be treated by installing catalytic converters in chimneys of the factories. Sulphur Dioxide: Sulphur dioxide (SO2) is considered a pollutant since it contributes to acidic rain. Sulphur dioxide is a product of two process, these are combustion of sulphur –containing fossil fuels and extraction of metals from their sulphide ores (such as zinc sulphide). The problem associated with sulphur dioxide is that when it rises in the air from chimneys of factories, it mixes with water vapour of clouds and air. This results in the formation of sulphuric acid (H2SO4). When it rains, rain water which falls becomes acidic. Acid rain causes death to water creatures since it makes water acidic, acidifies soil causing death to plants and deforestation, reacting with limestone from buildings and sculptures corroding it, and may also cause lung cancer. Sulphur dioxide could be treated before it leaves chimneys of factories by reacting it with limestone which is a neutralisation reaction. This process is called desulphurisation. SO2 + CaCO3 → CaSO3 + CO2 Oxides of Nitrogen (NO & NO2): Nitrogen oxides are formed at high temperatures as a result of nitrogen and oxygen reacting. In cars, engines have a very high temperature; this creates a chance for nitrogen and oxygen present in air in the engine to react forming nitrogen monoxide. N2 + O2 → 2NO
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The produced carbon monoxide is released through the exhaust with other waste fumes. Nitrogen monoxide reacts with more oxygen from air producing nitrogen dioxide. 2NO + O2 → 2NO2 The problem associated with nitrogen dioxide is similar to that of sulphur dioxide. It rises up in the air and mixes with rain water forming nitric acid. This causes acid rain. Nitrogen oxides can also cause health respiratory problems to humans and animals. To treat this issue, cars are now fitted with devices called catalytic converters which eliminate nitrogen oxides. Lead Compounds: Compounds of lead are waste products of fuel burning in cars. They are considered pollutants because they are poisonous and they are said to cause mental disabilities to young children. To treat this problem, gas stations now provide unleaded fuel. Catalytic Converters: Car fuels contain carbon; so carbon monoxide gas is released by cars as waste fumes, as well as nitrogen oxides. These are pollutant gases. To prevent these gases from polluting air, a device called catalytic converter is fitted at the end of the exhaust. This device contains a catalyst which catalyses the reaction between these two gases producing two harmless gases, nitrogen and carbon dioxide: 2NO + 2CO → 2CO2 + N2 2NO2 + 4CO → 4CO2 + N2 The catalyst of the device works best at temperature around 200°C. • State the adverse effect of common pollutants on buildings and on health The carbon cycle
The carbon cycle is a natural global cycle of the element carbon. It is what maintains a constant level of carbon dioxide in air (0.03%). The cycle goes as follows: Plants absorb carbon dioxide from air and undergo photosynthesis reaction which turns it into glucose and produces oxygen: 6CO2 + 6H2O → C6H12O6 + 6O2 The carbon is now stored in plants as glucose. One of two things happen, either the plants get eaten by animals or humans, or the plant dies and decays. If the plant is eaten by animals or humans, glucose in the plant is used by them in a process called respiration to release energy for their body.
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C6H12O6 + 6O2 → 6CO2 + 6H2O Respiration is the opposite of photosynthesis. Carbon dioxide is one of the products of it, which is released by the humans through breathing into the air. Thus carbon dioxide returns to the atmosphere. If the plant dies. It is buried underground and by time it decays forming coal and other fossil fuels. These substances contain the carbon which was made and stored by the plants and they are then taken by power stations which put them to use. Power stations burn carbon-containing fuels that were obtained as coal or fossil fuels formed by dead plants. This is a combustion reaction. C + O2 → CO2 Carbon dioxide is result of these reactions. Carbon dioxide produced is released to the air through chimneys of power stations. Thus the cycle is completed and all carbon dioxide returns to the atmosphere. Green House Gases:
The sun sends energy to the earth in two forms, light and heat. Some of the heat energy reflects back to the space, some however are trapped inside the Earth. This is caused by some gases and it is called the greenhouse effect. The main greenhouse gases are carbon dioxide and methane. • formation of carbon dioxide: – as a product of complete combustion of carbon containing substances – as a product of respiration – as a product of the reaction between an acid and a carbonate – from the thermal decomposition of a carbonate Methane, the other greenhouse gas is formed by animals. When animals eat and digest their food, methane gas is one of the waste products of this process. It is released to the atmosphere by animals. When plants die and decompose over many years, methane gas is also produced. The greenhouse effect poses a threat to the world now days. This is because greenhouse gases, especially carbon dioxide, have increased in amounts in the atmosphere due to activity of humans. Lots of fuel combustion is taking place around the world, increasing the levels of CO 2, while trees are being chopped off to made use of instead of leaving to replace CO 2 with oxygen. These activities cause an increase of the levels of CO2 in the atmosphere, which leads to more heat trapping in earth. This rises the global temperature of the earth causing what’s called global warming. Global warming is the increase of the temperature of the earth due to the increase of levels of greenhouse gases. Global warming has effects on the earth. To start with, it north and south poles, which are made of ice, will start to melt raising sea levels. The sea temperature will also
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rise causing death to marine lives. This is also accompanied by other natural disasters such as hurricanes and heavy rains. Humans could prevent this by reducing combustion of fossil fuels and leaving forests to live. Rusting: Rusting is the corrosion of iron as a result of reaction with oxygen from air and water. If iron objects are left uncovered and exposed to air & water, iron will react with oxygen forming hydrated iron oxide (also known as rust). Rust is a reddish brown flaky solid which will fall of the object making it thinner and loses it its shape. Iron must come in contact with air and water in order for rusting to happen. The formula of rust is Fe 2O3. xH2O. Steel can also rust since it is made up of mostly iron. Rusting can become very dangerous in some cases. For example, bridges that cross rivers stand on columns that are made of iron. The conditions of rusting are present in this case (Water from the river and oxygen from the air). There is a risk that the columns will rust and collapse with the whole bridge. In another case, ships are made of iron. Again, the conditions of rusting are present (water from the sea and oxygen from the air). In fact, this situation is more critical because sea water contains minerals that act as a catalyst to speed up the reaction of rusting. There some available methods to prevent rusting. These methods are based on covering the iron object with another substance to create a barrier between iron and oxygen and water so that rusting does not take place: Painting: The iron or steel object is painted all over. The paint creates the desired barrier to prevent iron or steel coming in contact with air and water. This method is used in car bodies and bridges. Electroplating: The iron or steel object gets electroplated with another metal that doesn’t corrode. The object is usually electroplated with tin or chromium since they are very unreactive. This method is used in food cans and car bumpers. Sacrificial Protection: This method is based on the idea that metals that are higher than iron in the reactivity series will react in preference to it and thus that metal is corroded and the iron is protected. Metals usually used as protectors in this method are zinc and magnesium since they are higher than iron in the reactivity series. In ships for example, zinc or magnesium bars are attached to the iron base of the ship which is in contact with water and oxygen from air. But rusting doesn’t take place since zinc or magnesium is the one that gets corroded. These bars must be replaced from time to time because once they all get corroded, iron becomes unprotected and rusts. This method is usually used in ships or bridge columns. The zinc or magnesium bars do not have to completely cover the iron or steel because as long as they are attached to each other the zinc or magnesium bars get corroded and not the iron. Galvanisation: Galvanisation is a very reliable method for preventing rusting. It is basically covering the whole object by a protective layer of zinc. This can be done either by electroplating the object with zinc or dipping it into molten zinc. The zinc layer provides a barrier that prevents iron or steel from coming in contact with air and water. The zinc gets corroded instead iron thus protecting it. If the a part of the zinc coat falls off and the iron or steel gets exposed to air and water, the bare part still doesn’t get corroded since it is protected by sacrificial protection now. Fertilisers Chemicals applied to plants to improve their growth and increase the amounts of products such as fruits, nuts, leaves, roots and flowers that they produce for us. They work by supplying plants with the vital elements they need including
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Nitrogen - in the form or nitrate (NO3- containing) salts; phosphorous – in the form of phosphate (PO43- containing) salts and potassium (K+ containing) salts. Displacement of ammonia from its salts Ammonia (NH3) is a smelly gas. One way to produce it is to react ammonium (NH4+) salts with an alkali (OH-) eg: NH4Cl + NaOH NH3 + H2O . + NaCl
The Haber process The raw materials for the process of making ammonia are hydrogen and nitrogen. Hydrogen is obtained by reacting natural gas (mostly methane) with steam, or from cracking oil fractions. Nitrogen is obtained from the air. Air is 78 per cent nitrogen and nearly all the rest is oxygen. When hydrogen is burned in air, the oxygen combines with the hydrogen - leaving nitrogen behind. In the Haber process, nitrogen and hydrogen react together under these conditions: a high temperature - about 450°C a high pressure - about 200 atmospheres (200 times normal pressure) an iron catalyst In addition, any unreacted nitrogen and hydrogen are recycled. The reaction is reversible. In a chemical equation, the symbol is used instead of an ordinary arrow if the reaction is reversible: This equation summarises the Haber process:
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Stage 1
Having obtained the hydrogen and nitrogen gases (from natural gas and the air respectively), they are pumped into the compressor through pipes.
Stage 2
The gases are pressurised to about 200 atmospheres of pressure inside the compressor.
Stage 3
The pressurised gases are pumped into a tank containing beds of iron catalyst at about 450°C. In these conditions, some of the hydrogen and nitrogen will react to form ammonia.
Stage 4
The unreacted nitrogen and hydrogen, together with the ammonia, pass into a cooling tank. The cooling tank liquefies the ammonia, which can be removed into pressurised storage vessels.
Stage 5
The unreacted hydrogen and nitrogen gases are recycled by being fed back through pipes to pass through the hot iron catalyst beds again.
The reaction mixture contains some ammonia, plus a lot of unreacted nitrogen and hydrogen. The mixture is cooled and compressed, causing the ammonia gas to condense into a liquid. The liquefied ammonia is separated and removed. The unreacted nitrogen and hydrogen are then recycled back into the reactor.
TIPS ON SPECIFIC TOPICS • To remember that carbon monoxide is poisonous (it binds to haemoglobin), think of the ‘nox’ in carbon monoxide as being short for noxious (poisonous). The effects of pollutant gases on nature are often confused, as not all pollutant gases are acidic. Know the different effects of carbon monoxide, sulphur dioxide and carbon dioxide • A common error is to think that fume cupboards keep air away from a reaction. Fume cupboards have a continuous airflow to allow poisonous vapours to escape through the fan •You do not need to know all the details about water treatment. Filtration and chlorination are the usual points examined. •The separation of gases from the air is complex. You will only be asked questions about boiling points and distillations – not about details of the distillation plant. •It is a common error to suggest that sulfur rather than sulfur dioxide is responsible for acid rain. Comments such as ‘sulfur dissolves in water to form acid rain’ are incorrect. •The reactions in the catalytic converter are not well understood. The best equations to remember are the reactions of nitrogen oxides with carbon monoxide to form nitrogen and carbon dioxide. •It is important that you do not muddle the effects of different pollutants: carbon dioxide and methane are linked to global warming and sulfur dioxide is linked to acid rain. •The two important regulating features of the carbon cycle are the uptake of carbon dioxide by photosynthesis and the production of carbon dioxide during respiration.
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•If you are asked to choose two methods that prevent rusting, try not to choose two that are similar. Don’t give answers such as ‘removing water and air’. These are explanations not methods •When you write equations for the formation of ammonium salts remember that no water is formed as a product. For example: ammonia + sulfuric acid ammonium sulfate. •You need to know the conditions used in the Haber process and why these particular conditions are used by referring to the equilibrium reaction.
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14 Sulfur Sulphur is a non metal element in group 6 of the periodic table. Sulphur has many useful properties which make it widely used in the industry. Sources of Sulphur: Sulphur is found in many places in the world in different forms. It usually exists in volcanic regions in USA, Mexico and Sicily. Sulphur could also be obtained from some metal ores like Copper pyrites (CuFeS2) and Blende (ZnS). Properties of Sulphur: In room temperature, sulphur is a yellow, brittle solid which doesn’t conduct electricity as it is a non-metal. Sulphur is insoluble in water. It is able to react with both metals and non-metals. Sulphur Dioxide: Sulphur dioxide is the product of combustion of sulphur or sulphur-containing fuels. As you have studied in the previous chapter, it is an air pollutant as it causes acid rain. However, SO2 has important uses too: Bleaching wood pulp for the manufacturing of paper It is used as a food preservative as it kills bacteria Manufacturing of Sulphuric acid Contact Process (Manufacturing of Sulphuric Acid): Sulphuric acid is one of the most important chemicals in the industry since it has a role in the manufacturing of almost every product. Sulphuric acid is manufactured by a process called Contact Process and it involves several steps: 1.Making the sulphur dioxide 2.Converting the sulphur dioxide into sulphur trioxide 3.Converting the sulphur trioxide into sulphuric acid 1. Making the sulphur dioxide Sulphur is first burned in air producing sulphur dioxide: S(s)+ O2(g)→ SO2(g) 2. Converting the sulphur dioxide into sulphur trioxide: This is a reversible reaction, and the formation of the sulphur trioxide is exothermic. 2SO2(g) + O2(g) ⇌ 2SO3(g) 3. Converting the sulphur trioxide into sulphuric acid This can't be done by simply adding water to the sulphur trioxide - the reaction is so uncontrollable that it creates a fog of sulphuric acid. Instead, the sulphur trioxide is first dissolved in concentrated sulphuric acid:
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H2SO4(l) + SO3(g) → H2S2O7(l) The product is known as fuming sulphuric acid or oleum. This can then be reacted safely with water to produce concentrated sulphuric acid - twice as much as you originally used to make the fuming sulphuric acid. H2S2O7(l) + H2O(l) → 2 H2SO4(l) The average percentage yield of this reaction is around 30%. Properties & Uses of Sulphuric Acid: Sulphuric acid is a very strong acid. It is a dibasic acid which means it every molecule of it produces two hydrogen ions when it is dissolved in water. Sulphuric acid has some other unique properties. For example, it is a dehydrating agent. This means it eliminates water from compounds.
E.g.: CuSO4.5H2O
CuSO4 + 5H2O
E.g.: C6H12O6
6C + 6H2O
It is also a drying agent. This means it removes water from mixtures. Don’t confuse that dehydrating agent.
TIPS ON SPECIFIC TOPICS •Sulfuric acid has two hydrogen ions that can be replaced. Make sure that you remember this when writing symbol equations for the reaction of sulfuric acid with metals, metal oxides and metal carbonates. •You will need to know the main reactions in the contact process and be able to write relevant equations. You also need to know why the particular conditions of temperature and pressure are used.
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15 Carbonates Carbonates are salts of carbonic acids (H2CO3). Carbonates are very useful salts, specially calcium carbonate (CaCO3). Sources of Calcium Carbonate: Calcium carbonate can be found in large amounts in the Peak District. It is found as a type of rocks called limestone near rivers. Forms of Calcium Carbonate: Limestone is not the only form of calcium carbonate. Marble and chalk are also other forms of this valuable salt. Chalk is made of shells of marine algae. Marble on the other hand, is a metaphoric rock made of limestone at high pressure. Uses of Calcium Carbonate: in the manufacture of iron and of cement Manufacture of Lime: One of the industrial uses of calcium carbonate is the manufacturing of lime from it. Lime is calcium oxide salt. This process takes place in a device called lime kiln and it is based on the thermal decomposition of calcium carbonate. Limestone is inserted in the kiln and heating starts. At the bottom of the kiln air is being blown in. this is also where lime is collected. The other product of this reaction, carbon dioxide gas, evolves and escapes at the top of the kiln. CaCO3 ⇌
CaO
(Limestone)
(Lime)
+ CO2 (Carbon Dioxide)
Uses of Lime: Lime can be used to neutralise soil acidity in farms. This is because it is a basic oxide. Slaked lime (Calcium hydroxide; Ca(OH)2) is also a basic oxide can be used as an alternative to lime for neutralising soil acidity. Another use of lime is neutralising sulphur dioxide waste in power stations. This is because sulphur dioxide is an acidic oxide while lime is a basic one. This process is called desulphurisation which you have studied earlier.
TIPS ON SPECIFIC TOPICS •It is a common error to suggest that oxygen is a reactant or product in the production of lime. The reaction is a thermal decomposition. Oxygen does not react with calcium carbonate and oxygen gas is NOT given off in the thermal decomposition.
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16 Organic chemistry 16.1 Names of compounds Alkanes share the same general formula: Alkanes are saturated hydrocarbons. This means that their carbon atoms are joined to each other by single bonds. Alkenes are a homologous series of hydrocarbons that contain a carbon-carbon double bond. The number of hydrogen atoms in an alkene is double the number of carbon atoms, so they have the general formula . Alcohols contain hydrogen and carbon but also possess one hydroxyl group (-OH). Their general formula is CnH(2n+1)OH. The names of alcohols end with ‘ol’, eg ethanol. Carboxylic acids contain the carboxyl functional group (-COOH). Carboxylic acids end in '-oic acid'. Their general formula is CnH2O2. 16.2 Fuels • Combustion Fuels are substances that react with oxygen to release useful energy. Most of the energy is released as heat, but light energy is also released. About 21 per cent of the air is oxygen. When a fuel burns in plenty of air, it receives enough oxygen for complete combustion.
Complete combustion Complete combustion needs a plentiful supply of air so that the elements in the fuel react fully with oxygen. Fuels such as natural gas and petrol contain hydrocarbons. When hydrocarbons burn completely: the carbon oxidises to carbon dioxide the hydrogen oxidises to water (remember that water, H2O, is an oxide of hydrogen) Here are the equations for the complete combustion of propane, used in bottled gas: propane + oxygen → carbon dioxide + water C3H8 + 5O2 → 3CO2 + 4H2O Incomplete combustion Incomplete combustion occurs when the supply of air or oxygen is poor. Water is still produced, but carbon monoxide and carbon are produced instead of carbon dioxide. In general, for incomplete combustion: hydrocarbon + oxygen → carbon monoxide + carbon + water The carbon is released as soot. Carbon monoxide is a poisonous gas, which is one reason why complete combustion is preferred to incomplete combustion. Gas fires and boilers must be serviced regularly to ensure they do not produce carbon monoxide. Here are the equations for the incomplete combustion of propane, where carbon is produced rather than carbon monoxide: propane + oxygen → carbon + water C3H8 + 2O2 → 3C + 4H2O
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Fractional distillation of crude oil Fractional distillation separates a mixture into a number of different parts, called fractions. A tall fractionating column is fitted above the mixture, with several condensers coming off at different heights. The column is hot at the bottom and cool at the top. Substances with high boiling points condense at the bottom and substances with lower boiling points condense on the way to the top. Crude oil is a mixture of hydrocarbons. The crude oil is evaporated and its vapours condense at different temperatures in the fractionating column. Each fraction contains hydrocarbon molecules with a similar number of carbon atoms and a similar range of boiling points. Oil fractions The diagram below summarises the main fractions from crude oil and their uses, and the trends in properties. Note that the gases leave at the top of the column, the liquids condense in the middle and thesolids stay at the bottom.
The fractionating column As you go up the fractionating column, the hydrocarbons have: 1.lower boiling points 2.lower viscosity (they flow more easily) 3.higher flammability (they ignite more easily). • Name the uses of the fractions as: – refinery gas for bottled gas for heating and cooking – gasoline fraction for fuel (petrol) in cars – naphtha fraction for making chemicals – kerosene/paraffin fraction for jet fuel – diesel oil/gas oil for fuel in diesel engines – fuel oil fraction for fuel for ships and home heating systems – lubricating fraction for lubricants, waxes and polishes – bitumen for making roads Other fossil fuels Crude oil is not the only fossil fuel.
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Natural gas mainly consists of methane. It is used in domestic boilers, cookers and Bunsen burners, as well as in some power stations. Coal was formed from the remains of ancient forests. It can be burned in power stations. Coal is mainly carbon but it may also contain sulfur compounds, which produce sulfur dioxide when the coal is burned. This gas is a cause of acid rain. Also, as all fossil fuels contain carbon, the burning of any fossil fuel will contribute to global warming due to the production of carbon dioxide. 16.3 Homologous series A homologous series is a family of compounds which have the same general formula and have a similar molecular structure and similar chemical properties because they have the same functional group of atoms (e.g. C=C alkene, C-OH alcohol or -COOH carboxylic acid). Members of a homologous series have similar physical properties such as appearance, melting/boiling points, solubility etc. BUT show trends in them e.g. steady increase in melting/boiling point with increase in carbon number or molecular mass. The functional group is a group atoms common to all members of a homologous series that confer a particular set of characteristic chemical reactions on each member molecule of the series. Characteristics of a homologous series: -all the compounds fit the same general formula -the chain length increases by 1 each time -as the chain gets longer, the compounds show a gradual change in properties. Structural isomers: have the same chemical formula, but different structures, they can be straight or branched. 16.4 Alkanes Alkanes are saturated hydrocarbons. This means that their carbon atoms are joined to each other by single bonds. This makes them relatively unreactive, apart from their reaction with oxygen in the air - which we call burning or combustion. Alkanes undergo a substitution reaction with halogens in the presence of light.
16.5 Alkenes Bromine water is an orange solution of bromine. It becomes colourless when it is shaken with an alkene. Alkenes can decolourise bromine water, but alkanes cannot. The reaction between bromine and alkenes is an example of a type of reaction called an addition reaction. The bromine is decolourised because a colourless dibromo compound forms. For example: ethene + bromine → dibromoethane C2H4 + Br2 → C2H4Br2 Other addition reactions of alkenes: Hydrogen can be added to a C=C double bond. This has the effect of ‘saturating’ the molecule, and will turn an alkene into an alkane. For example: C2H4 + H2 → C2H6 If steam (H2O) is added to an alkene, an alcohol is made. For example: C2H4 + H2O → C2H5OH
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Cracking Fuels made from oil mixtures containing large hydrocarbon molecules are not efficient as they do not flow easily and are difficult to ignite. Crude oil often contains too many large hydrocarbon molecules and not enough small hydrocarbon molecules to meet demand. This is where cracking comes in. Cracking allows large hydrocarbon molecules to be broken down into smaller, more useful hydrocarbon molecules. Fractions containing large hydrocarbon molecules are heated to vaporise them. They are then either: heated to 600-700°C passed over a catalyst of silica or alumina These processes break covalent bonds in the molecules, causing thermal decomposition reactions. Cracking produces smaller alkanes and alkenes (hydrocarbons that contain carbon-carbon double bonds). For example: hexane → butane + ethene C6H14 → C4H10 + C2H4 Some of the smaller hydrocarbons formed by cracking are used as fuels, and the alkenes are used to make polymers in plastics manufacture. Sometimes, hydrogen is also produced during cracking. Alkenes can be used to make polymers. Polymers are very large molecules made when many smaller molecules join together, end to end. The smaller molecules are called monomers. In general: lots of monomer molecules → a polymer molecule The polymers formed are called addition polymers. Alkenes can act as monomers because they are unsaturated: ethene can polymerise to form poly(ethene), also called polythene propene can polymerise to form poly(propene), also called polypropylene chloroethene can polymerise to form poly(chloroethene), also called PVC Polymer molecules are very large compared with most other molecules, so the idea of a repeat unit is used when drawing a displayed formula. When drawing one, you need to: 1.change the double bond in the monomer to a single bond in the repeat unit 2.add a bond to each end of the repeat unit
Addition polymerisation
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It can be tricky to draw the repeat unit of poly(propene). Propene is usually drawn like this:
It is easier to construct the repeat unit for poly(propene) if you redraw the monomer like this:
You can then see how to convert this into the repeat unit 16.6 Alcohols The alcohols are a homologous series of organic compounds. They all contain the functional group –OH, which is responsible for the properties of alcohols. The first three alcohols in the homologous series are methanol, ethanol and propanol. They are highly flammable, making them useful as fuels. They are also used as solvents in marker pens, medicines, and cosmetics (such as deodorants and perfumes). Ethanol is the alcohol found in alcoholic drinks such as wine and beer. Ethanol is mixed with petrol for use as a fuel. Ethanol from ethene and steam Ethanol can be manufactured by the hydration of ethene. In this reaction, ethene (which comes from cracking crude oil fractions) is heated with steam in the presence of a catalyst of phosphoric acid (to speed up the reaction):
This reaction typically uses a temperature of around 300°C and a pressure of around 60– 70 atmospheres. Notice that ethanol is the only product. The process is continuous – as long as ethene and steam are fed into one end of the reaction vessel, ethanol will be produced. These features make it an efficient process. However, ethene is made from crude oil, which is a non-renewable resource. Ethanol can also be made by a process called fermentation. Fermentation During fermentation, sugar (glucose) from plant material is converted into ethanol and carbon dioxide. This typically takes place at temperatures of around 30°C. The enzymes found in single-celled fungi (yeast) are the natural catalysts that can make this process happen:
Unlike ethene, sugar from plant material is a renewable resource.
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Hydration of ethene v fermentation Fermentation Type of raw materials Type of process Labour Rate of reaction Conditions needed
Renewable (glucose from plants)
Non-renewable (ethene from crude oil)
Batch (stop-start)
Continuous (runs all the time)
A lot of workers needed
Few workers needed
Slow
Fast
Warm (30°C), normal pressure (1 atm)
High temperature (300°C) and high pressure (60-70 atm)
Purity of product Impure (needs treatment) Energy needed
Hydration of ethene
A little
Pure (no by-products made) A lot
16.7 Acids Properties of carboxylic acids Short carboxylic acids are liquids and are soluble in water. Longer carboxylic acids are solids and are less soluble in water. The boiling point of a carboxylic acid is higher than that of the alkane with the same number of carbon atoms because the intermolecular forces are much stronger. Carboxylic acids are weak acids, so they can donate a hydrogen ion(H+) in acidbase reactions:
This means that they will react with carbonates to produce a salt, water and carbon dioxide: They will also react with reactive metals to produce a salt and hydrogen. Making a carboxylic acid Ethanoic acid can be made by oxidising ethanol (which is an alcohol). In this case, oxidation involves adding an oxygen atom and removing two hydrogen atoms. This can happen: during fermentation if air is present when ethanol is oxidised by an oxidising agent, such as acidified potassium manganate(VII) Making an ester Esters occur naturally - often as fats and oils - but they can be made in the laboratory by reacting an alcohol with an organic acid. A little sulfuric acid is needed as a catalyst. The general word equation for the reaction is: alcohol + organic acid → ester + water For example: methanol + butanoic acid → methyl butanoate + water The diagram shows how this happens, and where the water comes from:
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So, to make ethyl ethanoate, you would need to react ethanol with ethanoic acid.
Different esters have different smells. These smells are often fruity. Alcohol
Organic acid
Ester made
Smell of ester
Pentanol
Ethanoic acid
Pentyl ethanoate
Pears
Octanol
Ethanoic acid
Octyl ethanoate
Bananas
Pentanol
Butanoic acid
Pentyl butanoate
Strawberries
Methanol
Butanoic acid
Methyl butanoate
Pineapples
16.8 Macromolecules Macromolecules are large molecules built up from small units (monomers). Different macromolecules have different units and/or different linkages For example glucose (the small unit) can join together to make starch or cellulose (natural macromolecules). Examples of the small units: -glucose -amino acids -fatty acids and glycerol Examples of linkages: -amide -ester Examples of macromolecules: -starch -protein -lipids
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16.8 (a) Synthetic polymers
Uses of polymers Different polymers have different properties, so they have different uses. The table gives some examples: Polymer
Typical use
Poly(ethene)
Plastic bags ,bottles, gloves, cling film (low density), mugs, bowls, chairs, dustbins (high density)
Poly(propene)
Crates and ropes
Poly(chloroethene)
Water pipes, wellingtons, hoses and insulation on electricity cables
Polystyrene
Used as expanded polystyrene in fast-food cartons, packaging, and insulation for roofs and walls
Teflon
Coated on frying pans to make them non-stick, fabric protector, windscreen wipers, flooring nylon ropes, fishing nets and lines, tents, curtains
Terylene
Clothing (especially mixed with cotton), thread
Polymers have properties that depend on the chemicals they are made from and the conditions in which they are made. For example, there are two main types of poly(ethene) - LDPE, low-density poly(ethene), and HDPE, high-density poly(ethene). LDPE is weaker than HDPE and becomes softer at lower temperatures. Modern polymers are very useful. For instance, they can be used as: new packaging materials waterproof coatings for fabrics (eg for outdoor clothing) fillings for teeth dressings for cuts hydrogels (eg for soft contact lenses and disposable nappy liners) smart materials (eg shape memory polymers for shrink-wrap packaging) Pollution problems from plastics: -choke birds, fish and other animals that try to eat them. Or they fill up the animals’ stomachs so that they can’t eat proper food, and starve to death. -they clog up drains and sewers and cause flooding. -they collect in rivers, and get in the way of fish. Some river beds now contain a thick layer of plastic -they blow into trees and onto beaches. So the place looks a mess. Tourists become put off. • structure of the polymer product from a given alkene and vice versa
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Condensation polymers Some polymers are made via condensation polymerisation. In condensation polymerisation, a small molecule is formed as a by-product each time a bond is formed between two monomers. This small molecule is often water. An example of a condensation polymer is nylon.
Condensation polymerisation involves linking lots of small monomer molecules together by eliminating a small molecule. This is often water from two different monomers, a H from one monomer, and an OH from the other, the 'spare bonds' then link up to form the polymer chain. Nylon (a polyamide) is formed by condensation polymerisation, the structure of nylon represented below where the rectangles represent the rest of the carbon chains in each unit. For advanced molecular representations see Organic Nitrogen Compounds (A level Notes)
(3 units etc.) This is the same linkage (-CO-NH-) that is found in linked amino acids in naturally occurring macromolecules called proteins, where it is called the 'peptide' linkage.
Nylon-6,6 Terylene (a polyester) is formed by condensation polymerisation and the structure of Terylene represented as
(3 units etc.) This is the same kind of 'ester linkage' (-COOC-) found in fats which are combination of long chain fatty carboxylic acids and glycerol (alcohol with 3 -OH groups, a 'triol'). Terylene (polyester) and nylon are good for making 'artificial' or 'man-made' fibres used in the clothing and rope industries. In the manufacturing process the polymer chains are made to line up. This greatly increases the intermolecular forces between the 'aligned' polymer molecules and strong fibrestrands of the plastic can be made. Although these are actually thermoplastic polymers, nylon and terylene can be drawn out into thin strong fibres for use in clothing. Some important structure, strength and 1D to 3D dimension concepts are in the Chemical Bonding notes. Nylon and polyester are typical synthetic fibres which have, in many cases, replaced cotton, silk and wool fabrics in the clothing industry.
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They are cheap to make on an industrial scale compared to cotton from fields, silk from silkworms and wool from sheep. As well as being cheaper, the physical properties of synthetic fibres have several advantages compared to their natural predecessors like cotton, silk and wool. Compared to natural fibres, synthetic fibres tend to be .... lighter - outdoor or indoor clothing, more durable - harder tougher wearing fibres, water-resistant - better water-proofed fabrics, However, there are some disadvantages e.g. they are not very breathable and sweat builds up making you feel uncomfortable. A don't forget that silk fibres (for fabrics), rubber (for tyres and elastic objects) are very useful natural polymers. Wood, an extremely useful construction material, and is mainly a polymer mixture of cellulose (a polymer of glucose) and lignin (with a rigid cross-linked structure). The valuable crop of cotton (for fabrics) also has a molecular structure based on cellulose, in fact its the purest form of cellulose that occurs naturally.
16.8 (b) Natural macromolecules Food’s main constituents are proteins, fats and carbohydrates. Proteins contain the same linkages (amide links) as nylon, but with different units. Similarly, lipids and terylene both have ester links but different units. The structure of a protein is:
In digestion proteins are broken down into amino acids (hydrolysis). Fats are esters possessing the same linkage as Terylene (ester links) but with different units. Soap is a product of the hydrolysis of fat. It is done using sodium hydroxide (as opposed to acid, in digestion). The hydrolysis gives glycerol and the sodium salts of fatty acids. The salts are used as soaps. Complex carbohydrates: are a large number of joined sugar units (monosaccharide like glucose).
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Carbohydrates are a whole series naturally occurring molecules based on the elements carbon, hydrogen and oxygen. They are an important source of chemical energy in our diet. e.g. the respiration reaction glucose + oxygen ==> carbon dioxide + water C6H12O6 + 6O2 ==> 6CO2 + 6H2O Carbohydrates like glucose and fructose are used as sweeteners in food as well as sweets themselves. Historically the name 'carbohydrate' comes from the fact that all their formulae seemed to be based on Cx(H2O)y BUT this is not the way to think of their formula. They range from relatively small molecules called monosaccharide (means one basic unit), or disaccharide (two basic units combined) to very large natural polymers or macromolecules called polysaccharides (many units combined). The formation of complex carbohydrates: These are made of smaller carbon, hydrogen and oxygen based molecules combining together e.g. the polysaccharides starch and cellulose are formed from glucose, molecular formula C6H12O6. Their formation can be described in terms of a large number of sugar units joined together by condensation polymerisation Note: Condensation polymerisation means the joining together of many small 'monomer' molecules by eliminating an even smaller molecule between them to form the linkage. e.g. HO-XXXXX-OH + HO-XXXXX-OH
HO-XXXXX-O-XXXXX-OH + H2O etc.
n C6H12O6 ==> (C5H10O5)n + nH2O (where n is a very large number to form the natural polymer) The XXXXX or the [rectangles] below, represent the rest of the carbon chains in each unit (more detail in the 3rd diagram below).
plus many H2O etc.
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This diagram of starch or cellulose is in 'skeletal formula' style and both are polymers of glucose - can you see the connection between each 'unit' and the structure of glucose itself? The resulting natural polymer is called a polysaccharide. Acid hydrolysis of complex carbohydrates (e.g.. starch) gives simple sugars. This can be brought about by e.g. warming starch with hydrochloric acid solution to form glucose. (C5H10O5)n + nH2O ==> n C6H12O6 (where n is a very large number) The hydrolysis products from polysaccharides can be analysed with paper chromatography as in the case of amino acids. We can digest long molecules like starch, though they have to be broken down by enzyme action before the smaller molecules like glucose can be used in respiration. However, we cannot digest cellulose because we don't have the enzymes to effect this process. In digestion, the hydrolysis (Decomposition of a chemical compound by reaction with water, such as the dissociation of a dissolved salt or the catalytic conversion of starch to glucose, which can be accelerated by an acid or base) of starch happens in the mouth by the enzyme amylase to make glucose. In the lab, unless you have enzymes, you have to boil the complex carbohydrate (or proteins or fats) in acid the products will be the following: -starch → glucose -proteins → amino acids -fats → fatty acids and glycerol But if hydrolysis is not complete, the macromolecules are not completely broken down. So you get a mixture of molecules of different sizes for example for starch you get, glucose, maltose (2 glucose units) and maltotriose (3 glucose units). Chromatography can be used to identify the products and the substances. However, amino acids and sugars are colourless when dissolved in water, so a locating agent is used. The substances can be identified using the Rf values or by matching them with spots which are horizontal. The fermentation of simple sugars to produce ethanol Yeast is a microorganism containing an enzyme which will convert a sugar (glucose) solution into carbon dioxide and alcohol (ethanol). This process is called fermentation. The word equation for fermentation is glucose + yeast
carbon dioxide + ethanol.
Carbon dioxide gas bubbles out of the solution into the air leaving a mixture of ethanol and water Ethanol can be separated from the mixture by fractional distillation. Fermentation must be carried out in the absence of air to make alcohol. If air is present, ethanoic acid is made instead of alcohol. Fermentation works best at a neutral or slightly acidic pH.
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Fats and oils Fats and oils are naturally-occurring esters. Fats are solid at room temperature, whereas oils are liquids. Vegetable oils Vegetable oils are natural oils found in seeds, nuts and some fruit. The oil can be extracted. The plant material is crushed and pressed and the oil, eg olive oil, is squeezed out. Sometimes the oil is more difficult to extract and has to be dissolved in a solvent. Once the oil is dissolved, the solvent is removed by distillation and impurities (such as water) are also removed. This leaves pure vegetable oil, eg sunflower oil. Structure of vegetable oils Molecules of vegetable oils consist of glycerol and fatty acids. The diagram shows how three long chains of carbon atoms are attached to a glycerol molecule to make one molecule of vegetable oil.
The structure of a vegetable oil molecule
Fats and oils are esters possessing the same linkages as Terylene but with different monomer units formed from long chain fatty acids and the 'triol'
alcohol glycerol
, which has three C–O–H groups.
Glycerol is the alcohol plants and animals use to make oils and fats which are esters we use in food and soaps. Animals and plants combine glycerol and long chain fatty acids to make triglyceride esters – fats from animals and oils from plants. Most of them are esters of the tri–alcohol ('triol') glycerol (systematic name propane–1,2,3– triol, but that can wait until AS–A2 level). The carboxylic acids which combine with the glycerol are described as 'long–chain fatty acids'.
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The resulting ester is called a 'triester' or 'triglyceride' and they are the major components in animal fat, vegetable oils, and processed fats like margarine etc.. 'Traditional' soap is a product of the hydrolysis of fats from animals and vegetable oils from plants
'Soapy' soaps (not modern detergents) are the sodium salts of long chain fatty acids formed by heating fatty oils with concentrated alkalis like sodium hydroxide or potassium hydroxide to hydrolyse them. This is known as a saponification reaction and a typical equation is illustrated above and the general word equation quoted below. vegetable oil/animal fat + sodium hydroxide ==> soap molecule + glycerol This reaction breaks the fat molecule down into one glycerol molecule (a triol alcohol) and three sodium salts of the long chain carboxylic fatty acids that formed part of the original oil/fat ester.
TIPS ON SPECIFIC TOPICS • Examiners are often very particular. One way to please them is to use the word ‘only’ in the definition of a hydrocarbon i.e. the answer ‘a compound containing only carbon and hydrogen’. • Only one compound is formed in the reaction of ethene with steam. Remember, this is a simple addition reaction (one compound formed from two or more substances) – a common error is to say that hydrogen is also formed • When trying to identify ‘cracking’ reactions from a set a reactions given, look out for one molecule of reactant forming two or more molecules of product. Remember that cracking does not involve oxygen • ‘Clear’ does not mean ‘colourless’; when bromine is added to an alkene the colour change is red-brown to colourless, not red-brown to clear •When drawing the full structural formula of an organic compound you should show all atoms and all bonds. Don’t forget that there is a bond in the alcohol functional group - O – H. •When drawing alkenes make sure that there are not too many hydrogen atoms that form the double bond. Check to see that each carbon has four bonds. •Don’t get confused between petroleum and petrol. Petroleum is crude oil. Petrol, also known as gasoline, is a fraction obtained when we distil petroleum.
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•You do not have to remember the boiling range or typical number of carbon atoms in each fraction and where they condense in the fractionating column. •You will be expected to be able to balance symbol equations for the combustion of alkanes. Remember to balance the oxygen. •When describing cracking you must state that (i) large hydrocarbon molecules are broken down to smaller ones (ii) and alkenes (iii) using a high temperature (iv) a catalyst. •The test for an alkene is that it turns acqueous bromine colourless. Do not use the word ‘clear’ to. Aqueous bromine stays the same yellow or orange colour when an alkene is added. A comm mean colourlesson error is to write ‘no observations’. •When writing a symbol equation for the combustion of an alcohol, when you balance the oxygen, remember that the alcohol contains oxygen too. •Remember that in naming the carboxylic acids the carbon atom of the –COOH group is included. So CH3COOH is ethanoic acid because compounds with two carbon atoms have names beginning with ‘-eth’. •When plastics are burned, poisonous or toxic gases are given off– not just harmful gases. A common error is to suggest that sulfur dioxide is given off when plastics burn. Few plastics contain sulfur. •When writing the formula for an addition polymer don’t forget: (i) the double bond changes to a single bond and (ii) to include the continuation bonds. •A common error in writing formulae for polyamides and polyesters is to write all the bonding atoms in the same direction when the monomers each have only one type of functional group, is wrong •You should be able to recognize the repeating units in proteins as NHCOCH(R) and that this repeats along the chain. •The full structure for a fat and the equation for soap making do not need to be remembered. Sufficient information will be given to help you answer questions. •You do not need to know the structure of carbonhydrates but it is important to know how hydrolysis breaks down complex carbonhydrates using simplified formulae. •It is a common error to suggest that oxygen is required for the fermentation of glucose to ethanol.
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GLOSSARY acid = any substance that produces hydrogen ions, H+, when dissolved in water acidic solution = a solution with a pH less than 7 acid rain = rain with a pH less than 5.6; acid rain has been made more acidic than normal rain because sulfur oxides and nitrogen oxides have dissolved in it; acid rain causes damage to buildings made from limestone, damages metal structures, kills fish, damages leaves in trees so they photosynthesise less acidic soil = soil with a pH less than 7 activation energy = minimum amount of energy needed to start the reaction/for a successful collision. actual yield = the amount of product obtained when carrying out a reaction addition reaction = a reaction in which atoms are added to an unsaturated carbon compound; the atoms are added using the double bond as one of the double bonds breaks and is used to make two new bonds, e.g. alkenes and halogens addition polymer = polymer formed by addition polymerization; adding many unsaturated monomers using double bonds addition polymerization = the joining together of many unsaturated monomer molecules (double bonds) to form a long molecule; new monomers are added to the chain at the double bonds alcohol = a homologous series of organic compounds which has -OH as its functional group; ethanol is a member of this homologous series alkanes = a homologous series of hydrocarbons which are saturated as they have only single bonds between the carbon atoms alkenes = a homologous series of hydrocarbons which are unsaturated as they have at least 1 double bond somewhere in the chain allotrope = different forms of the same element e.g. diamond, graphite and the fullerenes are allotropes of carbon alloys = mixture of a metal and small amounts of other metals and non-metals, made to have certain improved properties eg harder, stronger, increased resistance to corrosion, increased heat or electrical resistance alkali = any base which is soluble in water alkali metal = any metal in group 1 of the Periodic Table, most reactive metals alkali solution = a solution with a pH larger than 7 anions = negative ions; attracted to anode anode = positive electrode in electrolysis arrangement = how particles are positioned compared to each other e.g. close together, far apart, in fixed positions atom = the smallest particle that can exits of an element atomic number = number of protons in the nucleus of an atom, determines the order and place of each element in the Periodic Table avogadro’s constant = 6.02 x 1023 balanced equation = numbers of atoms are the same on either side of the equation (any equation should be balanced as in any chemical reaction particles are only re-arranged and are not destroyed or created); also shows the ratio in which reactants react and products are produced during a chemical reaction base = a substance which can neutralise an acid to make a salt and water examples: metal oxides, metal hydroxides, bauxite = ore containing aluminium oxide from which aluminium is extracted blast furnace = a furnace used for getting iron from iron oxide with the help of carbon boiling = a process during which a liquid changes into a gas as its particles gain more energy and move a lot faster and also much farther apart from each other. further from gas to liquid; only happens at the boiling temperature as opposed to evaporation brine = concentrated sodium chloride solution catalyst = a substance which speeds up a reaction but which remains unchanged at the end of the reaction catalytic converter = a piece of equipment which is part of the exhaust of a car and which changes nitrogen oxides into nitrogen before they are released into the atmosphere
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cathode = negative electrode in electrolysis cation = positive ion ; attracted to cathode chain length = number of carbon atoms one after the other in an organic compound chemical bond = electrostatic attraction between atoms or ions chemical property = how it reacts chromatogram = the result of a chromatography chromatography = a separating technique which uses the difference in solubility in a given solvent between the different parts of a mixture to separate them; combustion = burning, the reacting of a substance with oxygen, exothermic complete combustion = combustion in sufficient oxygen which in the case of hydrocarbons produces carbon dioxide and water compound = a pure substance made from two or more different atoms joined together chemically concentration = the number of moles of per liter of solution; tells us how much solute is dissolved inn a solvent condensation = a process during which a gas changes into a liquid because its particles are having less energy, slow down and come much closer together condensation polymer = a long molecule formed by condensation polymerization e.g. nylon condensation polymerization = the joining together of many of two different monomer molecules to form one single long molecule during which a small molecule is removed for each link between the monomers. covalent bond = force of attraction between a pair of shared electrons and the nucleii of both atoms cracking = the breaking down of long-chain alkanes into smaller alkanes and alkenes using a catalyst and heat (500 C) crude oil (or petroleum) = a mixture of organic compounds formed, as a result of high temperatures and pressures, from the remains of living plants and animals which died millions of years ago; a fossil fuel crystallisation = the forming of crystals from a saturated solution decomposition = breaking down a compound into simpler substances delocalised electrons = electrons that can move between atoms; they are not part of 1 atom diamine = a type of amine with exactly two amino groups diatomic = 2 atoms only dicarboxylic acid = organic compounds that contain two carboxylic acid functional groups. diffusion = the movement of particles by which different substances mix as a result of the random motion of each of their particles displacement reaction = a reaction in which a more reactive metal or halogen takes the place of a less reactive metal or halogen in its compound distillate = the liquid obtained from distillation; the liquid which has evaporated and condensed distillation = a separating technique in which a mixture is heated, the substance with the lowest boiling point evaporates and is condensed back to liquid form ductile = can easily be drawn into wires, what metals are endothermic = absorbs energy electrical conductivity = conducts electricity for which it needs mobile charged particles electrodes = rods of ususally carbon which are used to make elctrical contact with the electrolyte electrolysis = a reaction which uses electricity to decompose a compound electrolyte = an ionic compound or acid which conducts electricity (molten or in solution) and which is decomposed as it conducts electrolytic cell = a beaker with an electrolyte, 2 electrodes, a power supply and leads which changes electrical energy into chemical energy electron = a sub-atomic particle which has a negative charge and no relative mass electronic configuration = the number of electrons on each energy level in an atom element = a pure substance that consists of 1 type of atom only empirical formula = the formula which gives the most simple ratio of atoms/ions in a molecule/formula unit
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equilibrium = is reached when the forward reaction and reverse reaction are going on at the same time; at this point the amount of reactant or product does not change. evaporation = a process during which a liquid changes into a gas as some of its particles at the surface gain more energy, move a lot faster and farther apart from each other and eventually escape from the liquid; happens at any temperature between melting and boiling point. exothermic = releases/gives out energy fermentation = the changing of sugars dissolved in water into alcohol and carbon dioxide by the enzymes in yeast at a temperature of between 30 C to 40 C. filtrate = the liquid/solution that goes through the filter paper fixed positions = particles in a solid cannot move from their positions because of the strong forces of attraction forces of attraction = forces which hold/pull particles together forward reaction = the reaction which produces the products fraction = a group of substances which has a specific boiling point/range/condenses at similar temperature (because they have a similar number of carbon atoms in them); fractional distillation = crude oil is heated to evaporate most components which then condense back at different levels in the fractionating column because they have differing boiling points; freezing = process during which a liquid changes into a solid as its particles lose energy, slow down and come closer together again fuel = a substance that can release a lot of energy e.g. by burning gas = a state of matter in which particles are far apart, have a lot of energy and move fast and randomly galvanising = the coating of steel or iron by zinc to protect it from rusting giant structure = a structure in which a very large number of atoms or ions are joined together strongly and continuously in all 3 directions; a large network of particles group = vertical column in Periodic Table half equation = equation showing what goes on at each electrode in electrolysis halogen = any element from group 7 in the Periodic table homologous series = a group of organic compounds which all have the same general formula, similar chemical properties because they have the same functional group, have a gradual trend in physical properties, and differ by one CH2 unit. hydrocarbon = a compound which has carbon and hydrogen only incomplete combustion = burning in not enough oxygen indicator = any chemical which can change colour when added to different chemicals, usually acids and bases inert = very unreactive inert gases = gases in group 0 intermolecular forces = weak forces of attraction between molecules ion = a charged atom or group of atoms (which has become charged because it has either lost or gained an electron(s)) ionic bond = strong electrostatic attraction between two oppositely charged ions, formed between metals and non-metals isomers = compounds with the same molecular formula but different structures or displayed formula and therefore different properties isotopes = atoms with the same number of protons and electrons but different number of neutrons; same mass number but different mass number lattice = regular 3-dimensional arrangement of the particles (atoms, ions or molecules) limestone = calcium carbonate liquid = a state of matter in which particles are close together but in a disorderly arrangement, they can move past one another and have energy to move from their positions lubricant = an oily but soft substance used to reduce friction between two moving surfaces malleable = easily shaped without breaking, what metals are mass number = the total number of protons and neutrons in the nucleus of an atom melting = a process during which a solid changes into a liquid as its particles have gained more energy and move from their positions and past one another into an irregular arrangment metallic bond = attraction between positive metal ions and delocalised (mobile ‘sea’ of electrons electrons
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metallic character = behaves like a metal, gives away electron (s) when it reacts to form a positive ion, conducts, shiny, malleable mixture = 2 or more substances mixed together which have not reacted and which are therefore easily separated by physical processes like evaporation/distillation/filtration molar mass = the actual mass of 1 mole or 6.02 x 1023 particles (atoms, ions, molecules or formula units) of that substance molar volume = the volume of 1 mole of a gas = 24L at rtp mole = the name given to a certain number and that number is 6.02 x 1023. molecular formula = shows the type of atoms/ions and their number/molar ratio in a molecule/formula unit molecule = a particle made up of 2 or more atoms held together by covalent bonds monomer = a small molecule which can be joined together to make a long molecule called a polymer; a monomer must have a double bond or a functional group at either end movement = how particles move e.g. fast, vibrate neutron = a sub-atomic particle with no charge and which is in the nucleus and has a relative mass of 1 neutralisation = a reaction between an acid and a base to produce water and a salt and sometimes also carbon dioxide noble gas = any element form the last group in the Periodic Table noble gas electronic configuration = the way in which electrons are arranged in the noble gas atoms which is that they have their outer shell full! This 2 electrons in the helium atom and 8 in the other noble gases ore = a mixture of rock which contains a useful chemical organic compounds = compounds that have the element CARBON in it oxidation = a reaction during which a substance gains oxygen; oxygen is added to the element or compound increasing its mass; also a reaction during which a substance loses an electron oxide = a compound which ends with oxygen oxidizing agent = a chemical which oxidises another chemical; it loses oxygen/gains electrons and becomes reduced oxide layer = layer of an oxygen compound % yield = the actual yield expressed as percentage of the theoretical yield period = horizontal row in the Periodic Table periodic trends = gradual changes in properties of the elements in the same period petroleum = a mixture of organic compounds formed, as a result of high temperatures and pressures, from the remains of living plants and animals which died millions of years ago; contains fossil fuels. pH = a number between 1 and 14 which tells us how strong or how weak an acid or alkali is pH scale = a scale running from 1- 14 used to show how acid or alkaline a substance is physical property = properties like melting and boiling point, volatility, conductivity, appearance, colour poly(ethene) = polymer made from polymerising ethene molecules - addition polymer polymer = a large molecule made from many small molecules that have been joined together; each polymer is made up of many repeated units polymerisation = a chemical reaction in which many small molecules or monomers are joined together to form a long molecule called a polymer position of equilibrium = gives an idea of how much reactant or product there is at equilibrium precipitation = a reaction between 2 salt solutions which produces an insoluble salt which sinks to the bottom of the test tube precipitate = insoluble solid formed during a reaction product = substance on right hand of equation pure substance = a single chemical element or compound which melts and boils at fixed temperatures rate of a reaction = amount of change in a reactant or product over a period of time; tells us how fast a reaction is going reactant = substance on left side of equation reactivity = refers to the ease with which a substance reacts with other substances
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reactivity series = a list of metals with the most reactive metal first based on results from experiments redox = a reaction during which both a chemical is oxidised and another is reduced reducing agent = a chemcial which reduces another chemical; it gains oxygen/loses electrons and becomes oxidised reduction = a reaction during which a susbstance loses oxygen and has its mass decreases; also a reaction during which a substance gains electrons relative atomic mass = the mass of an atom as compared to 1/12th of the mass of a carbon12 atom; it is also the average mass of all isotopes relative molecular mass = the sum of the relative atomic masses (multiplied by the number of times they are in the molecule) of the atoms in the molecule relative formula mass = the sum of the relative atomic masses (multiplied by the number of times they are in the formula) of the atoms or ions in the giant structure residue = the insoluble part that remains behind in the filter paper during a filtration or what is left in the flask reverse reaction = reaction which changes products back into reactants reversible reaction = a reaction during which products are made but are also changed back again into reactants rust = a loose orange brownflaky layer of hydrated iron oxide sacrificial protection = method of rust protection in which blocks of more reactive metal are attached to irn; the more reactive metal react with the air and water instead of the iron saturated solution = a solution which contains as much solute as possible saturated organic compound = each carbon atom in the organic compound has made 4 single covalent bonds simple molecular substance = substance made up of individual molecules held together by covalent bonds and has weak intermolecular forces between these molecules solid = a state of matter in which particles are close together and in a regular arrangement, can only vibrate in fixed positions and have little energy solvent = a liquid that does the dissolving solvent front = the height the solvent goes up to on the chromatography paper solute = a solid which dissolves solution = a mixture made by dissolving a solute in a solvent steel = an alloy of iron with other elements sub-atomic particle = very small particles from which atoms are made: electrons, protons and neutrons sublimation = a process during which a solid changes directly into a gas because its particles have alot more energy, move around very fast and are very far apart. system = the reactants and products of a reaction theoretical yield = the amount of product you should obtain according to the balanced equation and calculations thermal decomposition = breaking down of a compound by hetaing it transition element = metal in the transition block of the Periodic Table universal indicator = a mixture of indicators used to measure the pH because it goes different colours unsaturated organic compound = has at least one double bond; decolourizes brown bromine water valency = the combining power of an atom or group of atoms; in an ionic compound the valency of an ion is its charge; in a molecule the valency of an atom is the number of bonds it makes valency electrons = the electrons on the most outer shell; vapourise = change from liquid into gas vibrate = move forwards and backwards but in the same fixed position volatile = vapourises easily, low boiling point word equation = an equation in which the names of the chemicals are used
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EXAMINER TIPS PAPER 1 • Some questions may ask you to choose a combination of things in order to select the correct answer – understand exactly what is required before you start to answer • Within a single question, use a pencil to cross out the choices that are clearly incorrect, then choose between the others GENERAL TIPS FOR PAPERS 2, 3, 5 AND 6 • Read the question correctly. For example, if the question says ‘give two observations apart from temperature change’, don’t include temperature change in your answer • Check for contradictions within your answer. For example, a common error is to write ‘a white insoluble precipitate dissolves’ (6.5(a)(i)) Show any workings • In any calculation, the final answer should be to the correct number of significant figures – generally the same as the data. You may be penalized if you write an excess number of significant figures e.g. 1.257487 instead of 1.26 • Know your syllabus statements and definitions exactly – use the Revision Checklist on the website. Don’t add your own ideas to the statements. For example, the syllabus statement on batteries says ‘they are portable’, meaning they can be easily carried around: an answer such as ‘they are small’ may not be accepted, as something can be small yet heavy • If asked to ‘describe what you would observe’, write down what you see, hear or feel (e.g. ‘the test tube gets hot’). A common mistake is to write something like ‘a gas is given off’ or ‘copper is deposited’; these are not observations, these are conclusions • If asked to ‘describe what you would see’, don’t note observations about sounds or temperature • Learn your definitions! Questions such as “what is a compound?” or “Define the mole.” are often poorly answered. Define does not mean ‘give an example of.’ • When drawing diagrams: (i) make sure they fill the space given on the paper and are LABELLED (ii) when drawing apparatus for gas measurement, make sure that the gas cannot escape. For example, don’t draw a gas syringe with the plunger much smaller than the syringe barrel – this is a common error • When asked to give examples, give the number requested by the examiner. For example, if asked to give two examples, do not give three – if one is incorrect, you may lose a mark. If a question asks for a single use for a substance don’t write a list – the examiner will think you are ‘playing safe’ and you won’t get the mark • If you have to tick boxes to answer a question, make sure that you tick the correct number – don’t assume that a single answer is always required • In chemistry, when plotting a graph of reaction rate, you must draw a curve of best fit through your prints. Lines drawn with a ruler from point to point will not get a mark • Look out for ‘hidden words’ in questions such as ‘which of the following is a gas containing diatomic molecules?’ Many students focus on one or two words, and might forget ‘gas’. Underline key words and read the question slowly • Avoid vague statements. For example, if the question asks about the use of graphite, the answer ‘graphite is used for electrodes in electrolysis’ is appropriately specific. ‘Graphite is used in electrolysis,’ is too vague
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TIPS ON PRACTICAL PAPERS • Makes sure that you know the accuracy to which you can read burettes, measuring cylinders, etc. • If asked to describe the appearance of a substance, remember that there are generally two points to be made, the state and the colour • When making practical observations, use the words ‘precipitate’ rather than ‘cloudy’ if you cannot see through the test tube on adding one aqueous solution to another. Don’t forget the colour • ‘Test the pH’ means ‘give the pH number’, not just say whether something is acidic or alkaline •In (a) make sure that full and correct names are used. ‘Cylinder’, ‘stand’ and ‘spoon’ are not precise. ‘Measuring cylinder’, ‘tripod’ and ‘spatula’ are the correct names of apparatus. In (b) use a ruler and a sharp pencil and label the diagram clearly. A filter funnel and filter paper both need to be included in the answer. •Common mistakes are to label the box with the wrong acid. The solid is often incorrectly labeled as sodium sulfate, rather than sodium sulfite. In the answer to (b) an arrow needs to be positioned with its point touching the flask underneath the solid. Vague answers in (c) will not score credit, e.g. ‘There is no lid on the collecting vessel.’ Identifying clearly that the gas should be collected through water are the correct conclusions drawn from the supplied information. •Note the readings after checking the scale used in the diagrams. •A common error would be to misread the temperatures recorded when 10cm 3 and 40cm3 of Y have been added. Incorrect readings would be ‘36C’ and ‘46C’. •This question involves applying experience of common practical procedures to an unfamiliar situation. It also tests knowledge and understanding of chromatography. A diagram in (d) showing the solvent at the correct level and three separate colourings would indicate the ability to present information in a clear, logical form. When drawing chromatography apparatus, make sure that the place where you put the original spot of colour is above the level of the solvent. •The answer in (b) show that the knowledge of cation tests is good. The student correctly describes the effect of aqueous ammonia and sodium hydroxide on a solution containing Zn2+ (aq). •In the answers to (b) (i) and (ii) the examiners would prefer you to use the word ‘colourless’ instead of ‘clear’ when referring to a solution that is not coloured. •The anion tests in (e) and (f) are also recognized. However, the use of cobalt chloride to test for the presence of water was confused with the test for chlorine gas. •The standard of answers varied widely from excellent to very poor. Poor answers involved :
mixing the fuel with the water and measuring the temperature rise heating the fuel directly and measuring the temperature rise failure to ensure a fair test reference to the diagram with no detailed method.
Failure to show how a comparison of the results would indicate which fuel produces more energy was also common.
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•The response shows the ability to suggest suitable techniques and apparatus for the investigation. All measurements and observations to be made are clearly recorded. The idea of a fair test is clearly realized and the comparison of the results to draw appropriate conclusions made. This student also notes safety precautions and suggests repeating the experiment to check reliability.
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