Libro de texto_Chemistry and Physics 4 ESO by Edicions Talaiots

4 ESO

Chemistry and Physics Projecte de Seccions Europees Illes Balears MarĂa Paz Terrasa Sagrera

Edita: EDICIONS TALAIOT, S.L. C. Castelló de la Plana, 30 07181 Palmanova (Calvià) T. 971 75 16 04 edicionstalaiots@gmail.com www.edicionstalaiots.com Autora: María Paz Terrasa Sagrera Il·lustració i maquetació: Sara Socías Impressió: INSTITUT | Imprenta Digital. Palma ISBN 978-84-939484-9-8 D.L. PM-650-2012

4 ESO

Chemistry and Physics Projecte de Seccions Europees Illes Balears MarĂa Paz Terrasa Sagrera

Chemistry and Physics 3 ESO is aimed at CLIL students CLIL (Content and Language Integrated Learning) "CLIL refers to situations where subjects, or part of subjects, are taught through a foreign language with dual-focussed aims, namely the learning of content, and the simultaneous learning of a foreign language" (Marsh, 1994) This approach involves learning Chemistry and Physics through an additional language. It can be very successful in enhancing the learning of languages and other subjets, and developing in the youngsters a positive "can do" attitude towards themselves as language learners. (Marsh, 2000) This text is being updated and prepared to help those teachers involved in the experience of implementing bilinguism in European schools and in Spanish schools mainly.

Content Unit 1. Chemical bonding................................................................................................................................................ 7 1. Atoms, molecules, elements and compounds....................................................................................8 1.1 Atoms...........................................................................................................................................8 1.2 Elements.......................................................................................................................................8 1.3 Atomic Number and Atomic Weight.........................................................................................8 1.4 Molecules.....................................................................................................................................8 1.5 Compounds.................................................................................................................................9 2. Atom models.........................................................................................................................................10 2.1 Bohr atom model......................................................................................................................10 2.2 Types of atomic orbitals...........................................................................................................10 3. The periodic table.................................................................................................................................11 4. Periodic properties...............................................................................................................................12 4.1 Atomic radius.............................................................................................................................12 4.2 Ionization energy......................................................................................................................12 5. Ionic bonding.........................................................................................................................................12 5.1 The importance of noble gas structures................................................................................12 5.2 Ionic bonding in sodium chloride...........................................................................................12 5.3 Some other examples of ionic bonding.................................................................................13 5.4 Properties of ionic compounds...............................................................................................13 6. Covalent bonding..................................................................................................................................14 6.1 Simple covalent molecules......................................................................................................14 6.2 Covalent molecules properties................................................................................................15 7. Metallic bonding...................................................................................................................................15 Unit 2. Chemical changes.............................................................................................................................................19 1. Chemical reactions...............................................................................................................................20 1.1 Balancing chemical equations.................................................................................................20 1.2 Relative atomic mass................................................................................................................20 1.3 Relative formula mass..............................................................................................................21 1.4 The mole....................................................................................................................................21 2. Calculations from equations...............................................................................................................22 2.1 Calculations involving only masses.........................................................................................22 2.2 Calculations involving gas volumes........................................................................................22 2.3 Calculations involving solutions..............................................................................................23 3. Energy in chemical reactions...............................................................................................................24 4. Rates of reaction...................................................................................................................................24 4.1 Concentration of Reactants.....................................................................................................25 4.2 Temperature..............................................................................................................................25 4.3 Medium......................................................................................................................................25 4.4 Presence of Catalysts and Competitors.................................................................................25 5. Acid base reactions...............................................................................................................................25 6. Redox reactions....................................................................................................................................26 Unit 3. Introducing Organic Chemistry................................................................................................................ 31 1. Carbon compounds..............................................................................................................................32 2. Classification of organic compounds.................................................................................................32 2.1 Hydrocarbons............................................................................................................................32 2.2 Conjugated isolated. Double bonds - double bonds............................................................32 2.3 Compounds with single-bonded functional groups.............................................................32 2.4 Carbonyl compounds...............................................................................................................32 3. Organic compounds in Biology...........................................................................................................33 3.1 Carbohydrates...........................................................................................................................33 3.2 Lipids...........................................................................................................................................33 3.3 Proteins......................................................................................................................................35 3.4 Nucleic acids..............................................................................................................................35 4. Polymers................................................................................................................................................35 5. Naming organic compouds.................................................................................................................36 Unit 4. Motion and Position.........................................................................................................................................37 1. Coordinate system...............................................................................................................................38 2. Motion....................................................................................................................................................39 3. Speed and velocity................................................................................................................................40 3.1 Vector Basics..............................................................................................................................40 3.2 Instantaneous velocity..............................................................................................................40 4. Motion in a straight line at a constant speed....................................................................................41 4.1 Distance-time graphs................................................................................................................41 4.2 Velocity-time graphs.................................................................................................................42

5. Acceleration...........................................................................................................................................44 6. Accelerated motion..............................................................................................................................46 6.1 Equations of accelerated motion............................................................................................47 7. Circular motion ....................................................................................................................................47 7.1 Formulas for uniform circular motion....................................................................................48 Unit 5. Forces.......................................................................................................................................................................53 1. Forces and its effects............................................................................................................................54 2. Measuring forces..................................................................................................................................54 3. Forces as vectors..................................................................................................................................56 3.1 Vector Addition..........................................................................................................................56 3.2 The Pythagorean Theorem......................................................................................................56 4. Newton´s laws......................................................................................................................................57 4.1 Newton's First Law of Motion..................................................................................................57 4.2 Newton's Second Law of Motion.............................................................................................57 4.3 Newton's Third Law of Motion................................................................................................58 5. Friction force.........................................................................................................................................58 6. Forces and motion................................................................................................................................59 6.1 Forces in rectilinear movement..............................................................................................59 6.2 Forces in circular motion..........................................................................................................60 7. Gravitational forces..............................................................................................................................61 8. Free fall..................................................................................................................................................62 8.1 Mass and weight.......................................................................................................................62 8.2 Motion under free fall..............................................................................................................63 Unit 6. Fluid dynamics...................................................................................................................................................69 1. Density and pressure...........................................................................................................................70 1.1 Density........................................................................................................................................70 1.2 Pressure.....................................................................................................................................70 2. Pressure in liquids. Hydraulic pressure.............................................................................................71 3. Pascal´s principle.................................................................................................................................73 4. Archimedes´principle and flotation...................................................................................................74 5. Pressure from the atmosphere..........................................................................................................75 Unit 7. Work, energy and heat...................................................................................................................................79 1. Work.......................................................................................................................................................80 2. Power.....................................................................................................................................................80 3. Energy. Gravitational potential energy. Kinetic energy....................................................................81 3.1 Gravitational Potential Energy.................................................................................................81 3.2 Kinetic energy............................................................................................................................83 4. Mechanical energy. Energy transfers.Conservation of energy.......................................................83 5. Thermal energy.....................................................................................................................................85 5.1 Temperature..............................................................................................................................85 5.2 Heat............................................................................................................................................86 6. Specific heat capacity...........................................................................................................................86 7. Latent heat of fusion and vaporization..............................................................................................88 8. Thermal expansion...............................................................................................................................89 9. Conduction, convection, radiation......................................................................................................90 Unit 8. Waves, light and sound................................................................................................................................. 95 1. Vibrational motion................................................................................................................................96 2. Categories of waves.............................................................................................................................97 2.1 Longitudinal and transverse waves........................................................................................97 2.2 Electromagnetic and mechanical waves................................................................................98 3. Properties of waves..............................................................................................................................98 3.1 Amplitude...................................................................................................................................98 3.2 Wavelength................................................................................................................................99 3.3 Period.........................................................................................................................................99 3.4 Frequency..................................................................................................................................99 3.5 Speed of a wave..................................................................................................................... 100 4. Sound waves and properties........................................................................................................... 101 4.1 The speed of sound............................................................................................................... 101 4.2 Sound properties................................................................................................................... 102 5. Light waves and properties.............................................................................................................. 103 5.1 Reflection of light................................................................................................................... 103 5.2 Refraction of light................................................................................................................... 104

UNIT

Chemical bonding

1. Atoms, molecules, elements and compounds 2. Atom models 3. The periodic table 4. Periodic properties 5. Ionic bonding 6. Covalent bonding 7. Metallic bonding

1 UNIT Chemical bonding

1. Atoms, molecules, elements and compounds 1.1 Atoms All substances are made up of matter and the fundamental unit of matter is the atom. The atom constitutes the smallest particle of an element. The atom is made up of a central nucleus containing protons positively-charged and neutrons with no charge. The electrons, negatively-charged with negligible mass, revolve around the nucleus in different imaginary paths called orbits or shells.

1.2 Elements An element is a substance made up of atoms of one kind. There are about 82 naturallyoccurring elements and about 31 artificially-made elements as listed in the Periodic Table.

1.3 Atomic Number and Atomic Weight Atomic number of an element is the number of protons in the nucleus of an atom. Since atoms are electrically neutral, the number of protons equals the number of electrons in an atom. Atomic weight or relative atomic mass of an element is the number of times an atom of that element is heavier than an atom of hydrogen. The atomic weight of hydrogen is taken to be unity. Mass number of an element is the sum of the number of protons and neutrons in the nucleus of an atom.

1.4 Molecules A molecule is formed when atoms of the same or different elements combine. A molecule is the smallest particle of a substance that can normally exist independently. Examples : â&#x20AC;˘ Two atoms of oxygen combine to form a molecule of oxygen O2. â&#x20AC;˘ One atom of carbon combines with two atoms of oxygen to form a molecule of carbon dioxide CO2. â&#x20AC;˘ Two hydrogen atoms combine to form a molecule of hydrogen H2.

Chemical bonding UNIT

1.5 Compounds A compound is formed when atoms or molecules of different elements combine. In a compound, elements are chemically combined in a fixed proportion. Examples are : â&#x20AC;˘ Hydrogen and oxygen are combined in a fixed proportion of 2:1 to form the compound water H2O. â&#x20AC;˘ Carbon and oxygen are combined in a fixed proportion of 1:2 to form the compound carbon dioxide CO2.

H2O

NaCl

ACTIVITY 1. Complete the following table: Name Aluminium cation

Symbol

Number of protons

Number of electrons

Number of neutrons

Charge

27 Al 3+ 13

29 35

1 UNIT Chemical bonding

2. Atom models During the twentieth century scientific discoveries demonstrated that Rutherford´s atom model was not accurate. The main find was that the electrons spinned at a certain distance from the nucleus and they have stable orbits.

2.1 Bohr atom model The electrons, negative charged, are attracted by the nucleus which has a positive charge, but since they are always spinning around it, they do not fall on the nucleus. Bohr demonstrated that the electrons turn around the nucleus only in a determined orbit and they do not lose energy. So, the electrons are arranged in different shells with a determined energy, so that the shells are called energy levels. • In the first level next to the nucleus there are only 2 electrons • In the second one, 8 electrons • In the third one up to 18 electrons • In the fourth level there are 32 electrons N= 32

M= 18

L= 6 K= 2 Nucleus

The electrons occupy the lowest energy level which is empty.Then if an atom has only one electron, it will be placed in the first energy level.

2.2 Types of atomic orbitals There are different types of orbitals represented by s, p, d, f. The shape and size depend on the energy level and the sublevel. • Orbitals s have a spherical shape • Orbitals p have a figure-eight • Orbitals d and f are more complex 10

Chemical bonding UNIT

Moreover, orbital s from second level called 2s, has a larger size than orbital s from level one, called 1s. The picture below shows a diagram to arrange the electrons in an atom. It is known as Moeller´s diagram. Observe that in each level there are a determined number of orbitals. In level 1, there are only one s orbital. In level 2, there are orbitals s and p. In level 3, there are s, p and d. In level 4 there are orbital s, p, d, and f. The same happens in levels 5, 6 and 7.

WORKED EXAMPLE 1. Write down the electronic configuration of the lithium atom; Z = 3. We must arrange three electrons. So in the first level, there are 2 and in the second there is only 1. The arrangement for lithium is Li:1s²2s¹.

3. The periodic table The periodic table is one of the most important tools that a chemist has. The elements are arranged in order of atomic number in such a way that the periodic properties, chemical periodicity of the elements, are made clear. The standard form of the table includes periods, usually horizontal in the periodic table, and groups usually vertical. Elements in groups have some similar properties to each other. There is no one single or best structure for the periodic table but by whatever consensus there is, the form used here is very useful. The periodic table is a masterpiece of organised chemical information. The evolution of chemistry's periodic table into the current form is an astonishing achievement with major contributions from many famous chemists and other eminent scientists. 1

Hydrogen 1.007 94 (7)

Nitrogen

Oxygen

Fluorine

6.941 (2)

9.012 182 (3)

10.811 (7)

12.0107 (8)

14.0067 (2)

15.9994 (3)

18.998 4032 (5)

20.1797 (6)

22.989 770 (2)

24.3050 (6)

otassium 39.0983 (1)

Cal ium 40.078 (4)

S andium 44.955 910 (8)

itanium 47.867 (1)

anadium 50.9415 (1)

C romium 51.9961 (6)

Manganese 54.938 049 (9)

ron 55.845 (2)

Cobalt 58.933 200 (9)

Ni el 58.6934 (2)

ubidium

85.4678 (3)

Strontium 87.62 (1)

88.905 85 (2)

57-71

Caesium 132.905 45 (2)

Fran ium (223)

ttrium

Lant anoids

Barium 137.327 (7)

ir onium 91.224 (2)

Ha nium 178.49 (2)

89-103 A tinoids

104

ut er ordium (261)

adium (226)

Lant anoids

A tinoids

Niobium 92.906 38 (2)

antalum 180.9479 (1)

105

ubnium (262)

Molybdenium 95.94 (2)

ungsten 183.84 (1)

106

Seaborgium (266)

netium (98)

enium 186.207 (1)

107

Bo rium (264)

Lant anum 138.9055 (2)

Cerium 140.116 (1)

raseodymium 140.907 65 (2)

Neodymium 144.24 (3)

A tinium (227)

orium 232.0381 (1)

rota tinium 231.035 88 (2)

ranium 236.028 91 (3)

utenium 101.07 (2)

Osmium 190.23 (3)

108

Hassium (277)

romet ium (145)

Ne tunium (237)

odium 102.905 50 (2)

ridium 192.217 (3)

109

Meitnerium (268)

alladium 106.42 (1)

latinum 195.078 (2)

110

armstadtium (271)

Cu Co

Aluminium

Sili on

26.981 538 (2)

28.0855 (3)

Helium 4.002 602 (2)

Carbon

Magnesium

Boron

Sodium

Beryllium

Litium

30.973 761 (2)

35.453 (2)

39.948 (1)

orous

C lorine

allium 69.723 (1)

ermanium 72.64 (1)

Arseni 74.921 60 (2)

Selenium

Bromine

63.546 (3)

in 65.409 (4)

78.96 (3)

79.904 (1)

ellurium 127.60 (3)

odine 126.904 47 (3)

Sil er 107.8682 (2)

old 196.966 55 (2)

Cadmium 112.411 (8)

Mer ury 200.59 (2)

ndium 114.818 (3)

allium 204.3833 (2)

in 118.710 (7)

Lead 207.2 (1)

Antimony 121.760 (1)

Bismut 208.980 38 (2)

Te 84

olonium (209)

Astatine (210)

adoliniu m 157.25 (3)

erbium 158.925 34 (2)

ys rosium 162.500 (1)

Holmium 164.930 32 (2)

rbium 167.259 (3)

ulium 168.934 21 (2)

tterbium 173.04 (3)

100

101

102

ry ton 83.798 (2)

enon 131.293 (6)

Rn adon (222)

111

uro ium 151.964 (1)

Ameri ium (243)

lutonium (244)

Argon

oentgenium (272)

Samarium 150.36 (3)

Sul ur 32.065 (5)

Neon

Cm Curium (247)

Ber elium (247)

Cali ornium (251)

insteinium (252)

Fm Fermium (257)

Mendele ium (258)

Nobelium (259)

Lutetium 174.967 (1)

103

La ren ium (262)

1 UNIT Chemical bonding

4. Periodic properties 4.1 Atomic radius Unlike a ball, an atom does not have a fixed radius. The radius of an atom can only be found by measuring the distance between the nuclei of two touching atoms, and then halving that distance. The trends in atomic radius in Periods 2 and 3 are represented in the figure below.

4.2 Ionization energy The ionization energy is a measure of the energy required to remove one electron from one mole of gaseous atoms or ions. Noble gases have the more stable electronic configuration. It is difficult to remove one electron from them. So, their ionization energy is large. The opposite occurs in alkali metals because it is easier to remove one electron from them. ACTIVITY 2. Observe the pairs of elements and say:

Na and K

B and Al Ga and Br P and Ar

a. The element that has larger atomic radius b. The element that has a larger ionization energy c. The most electronegative atom

5. Ionic bonding 5.1 The importance of noble gas structures At a simple level of knowledge, a lot of importance is attached to the electronic structures of noble gases like neon or argon which have eight electrons in their outer energy levels or two in the case of helium. These noble gas structures are thought of as being in some way a "desirable" thing for an atom to have. You may well have been left with the strong impression that when other atoms react, they try to organise things such that their outer levels are either completely full or completely empty.

5.2 Ionic bonding in sodium chloride Sodium, Na : 1s2 2s2 2p6 3s1, has 1 electron more than a stable noble gas structure (s2p6). If it gave away that electron it would become more stable. The outer shell disappears, so that the new outer one has 8 electrons. 12

Chemical bonding UNIT

Chlorine, Cl : 1s22s22p63s23p5, has 1 electron short of a stable noble gas structure. If it could gain an electron from somewhere it too would become more stable. The answer is obvious. If a sodium atom gives an electron to a chlorine atom, both become more stable. The sodium has lost an electron, so it no longer has equal numbers of electrons and protons. Because it has one more proton than electron, it has a charge of 1+. If electrons are lost from an atom, positive ions are formed. They are called cations. The chlorine has gained an electron, so it now has one more electron than proton. It therefore has a charge of 1-. If electrons are gained by an atom, negative ions are formed. They are called anions. The nature of the bond The sodium ions and chloride ions are held together by the strong electrostatic attractions between the positive and negative charges. The formula of sodium chloride You need one sodium atom to provide the extra electron for one chlorine atom, so they combine together 1:1. The formula is therefore NaCl.

5.3 Some other examples of ionic bonding Magnesium oxide Again, noble gas structures are formed, and the magnesium oxide is held together by very strong attractions between the ions. The ionic bonding is stronger than in sodium chloride because this time you have 2+ ions attracting 2Ë&#x2030; ions. The greater the charge, the greater the attraction. The formula of magnesium oxide is MgO. Calcium chloride This time you need two chlorines to use up the two outer electrons in the calcium. The formula of calcium chloride is therefore CaCl2.

5.4 Properties of ionic compounds Ionic compounds are usually hard, brittle, and they conduct electricity when molten or dissolved. Also, they have high melting and boiling points. Most are soluble in polar solvents such as water. Also they have a crystal structure or crystal lattice. The figure represents the arrangement of ions in a crystal lattice where the blue spheres show the anion and the grey ones the cation.

ACTIVITIES 3. Explain what happens when potassium bonds with chlorine. 4. Why is the bonding between a metal and a non metal called ionic bonding ? 5. a) Why are ions attracted and packed in a regular pattern ?

b) Why do ionic solids have high melting points ?

1 UNIT Chemical bonding

6. Covalent bonding As well as achieving noble gas structures by transferring electrons from one atom to another as in ionic bonding, it is also possible for atoms to reach these stable structures by sharing electrons to give covalent bonds.

6.1 Simple covalent molecules Hydrogen Hydrogen atoms only need two electrons in their outer level to reach the noble gas structure of helium. Once again, the covalent bond holds the two atoms together because the pair of electrons is attracted to both nuclei. There is a strong force of attraction between them holding them together. This force is called covalent bond. The bonded atoms form a molecule. Hydrogen gas is made up of hydrogen molecules, and for this reason its formula is H2.

Chlorine Two chlorine atoms could both achieve stable structures by sharing their single unpaired electron as in the diagram.

The fact that one chlorine has been drawn with electrons marked as crosses and the other as dots is simply to show where all the electrons come from. In reality there is no difference between them. Hydrogen chloride The hydrogen has a helium structure, and the chlorine an argon structure.

Chemical bonding UNIT

6.2 Covalent molecules properties • Covalent compound exists as a separate molecules because they are formed by neutral atoms, they are electrically neutral and the force of attraction between these molecules is small. • Due to weak intermolecular forces, generally covalent molecules or covalent compounds are liquids and gases. However, some covalent substances are solids like iodine, liquid like H2O, Br2 or gas e.g. CO2, H2, Cl2, NH3. • They are volatile. Melting point and boiling point are low. • Covalent compounds are generally insoluble in water. • Covalent compounds are non-electrolyte because they do not conduct electricity. • Non-polar covalent compounds do not conduct electricity. Polar covalent compounds conduct small amount of electricity. ACTIVITIES 7. Represent the (F3) molecule by Lewis structure. Does it exist ? 8. Represent by Lewis structures the following molecules : H2, NH3, Br2, CH4, CCl4, CO2, O2, N2.

7. Metallic bonding When sodium atoms bond together to form the solid metal, the outer electron on each sodium atom becomes free to move throughout the whole structure. The electrons are said to be delocalised. These electrons are no longer attached to a particular atoms. Instead, you can think of them as flowing around throughout the whole metal. The sodium becomes an ion. The attraction of each positive ion to the delocalised electrons holds the structure together. Metallic bonding is sometimes described as an array of positive ions in a « sea of electrons »

1 UNIT Chemical bonding

In magnesium, two outer electrons are delocalised, leaving behind ions that carry a charge of 2+. It packs more efficiently. The properties of metals are : • They allow electricity and heat to pass. They are good conductors. • They are solids at room temperature, except mercury, and most of them have high melting points. • Most of them can be hammered into different shapes, malleable, and drawn into wires known as ductile. ACTIVITIES 9. Explain what these words mean: malleable, ductile, brittle. 10. Why do metals have in general, large densities ? 11. Write down the similarities of the three types of bonding.

Chemical bonding UNIT

Questions 1. Which are the similarities and the differences between a proton and an electron ? 2. Write down the electronic configuration of the following atoms : B (Z=5) ; C (Z=6) ; Ne (Z= 10) ; S (Z =16) 3. Group VIII elements are called noble gases. a. What do they have in common ? b. How many electrons do kripton, xenon and radon have ? c. How many electrons do they have in their outer shell ? 4. Answer the following questions : a. What are ions ? b. And cations and anions ? c. Is the same a crystal as a molecule ? d. What does Ca²+ mean ? 5. Make a drawing of the electronic configuration including details from the nucleus of Neon and the Phosphorus. 6. Write down the electronic configuration of the O (Z=8) and the O²ˉ, of the Na (Z= 11) and the Na+. 7. How is an anion formed? And a cation ? 8. Which electronic configuration acquires the calcium when loses two electrons ? 9. Which type of compound will form O and Mg when combining ?. And the Na with O? Look up the information you need into the periodic table. 10. Why do atoms bond ?. Explain the main difference between ionic and covalent bonding. 11. Complete the table :

Name

N P Ne S Ca

protons

electrons

neutrons

Elec. conf.

Group

Period

Ion Symbol

7 15

16 1s22s22p6

20 16

16 20

12. Propose simple laboratory experiences to determine if a substance is ionic.

1 UNIT Chemical bonding Questions

13. Is it possible to produce an ionic bonding between two non metals ? Give reasons. 14. Ionic compounds do not form molecules. What does NaCl mean ? 15. Write down the ions that form each of the following compounds : NaCl, CaCl2, FeCl3, FeCl2. 16. Explain the type of bonding in the following substances : Cl2, CaO, PbO, HBr, Sn, SO2, Ni, CaCl2, N2. 17. Classify the following elements without consulting the Periodic Table, in metals, non metals, noble gases and write down the symbol for their ions. a. Z=7 b. Z=10 c. Z=12 d. Z= 11 e. Z=17 18. Write down the electronic configuration of the elements which atomic numbers are 12 and 17. a. Which group do they belong to ? b. Which type of ion do they form ? c. Write down the formula for the compound when they bond together d. Write down the number of the period where they are found 19. Explain why sodium chloride conducts electricity when dissolved in water whereas it does not in solid state.

2 UNIT

Chemical changes

1. Chemical reactions 2. Calculations from equations 3. Energy in chemical reactions 4. Rates of reaction 5. Acid base reactions 6. Redox reactions

2 UNIT Chemical bonding

1. Chemical reactions When two substances combine to produce a new one, the mass of the reactants is the same as the mass of the products. This is known as the Lavoisier´s law.

1.1 Balancing chemical equations A chemical equation is balanced when the number of atoms of each type on each side of the equation is the same. This means that if you have 4 hydrogens on the left hand side of the equation, you must have 4 hydrogens on the right hand side, if there are 2 oxygens on the left, there must be 2 oxygens on the right, and so on. This is because of the law of conservation of mass; you can not make or destroy atoms during a chemical reaction.

But you can not just add atoms at random to each side, you have to work with the molecules of the reactants. Also, you will find it very tricky to try to balance a word equation, it is very much easier to use a chemical equation with chemical symbols, as then you will be able to see how many atoms of each type are in each chemical. A good example to show how to balance equations is the reaction between propane and oxygen. The products are carbon dioxide and water. The unbalanced equation is below.

C3H8 + O2 → H2O + CO2 There are three carbons on the left, but only one on the right. There are eight hydrogens on the left but only two on the right. There are two oxygens on the left but three on the right. The formula of the compounds can not be changed. The numbers 5, 4 and 3 that appear in the balanced equation below, represent the number of molecules involving during the reaction. Observe that if you multiply these numbers by the numbers representing the atoms in a compound, you can see that the numbers of carbon, hydrogen and oxygen are the same on both sides of the equation. The balanced equation is :

C3H8 + 5 O2 → 4 H2O + 3 CO2 1.2 Relative atomic mass It is the ratio of the average mass per atom of the naturally occurring form of an element to one-twelfth the mass of an atom of carbon-12. Symbol Ar. Abbreviation is R.A.M. Former name is atomic weight.

Chemical bonding UNIT

1.3 Relative formula mass Most atoms exist in molecules. You can use the Relative Atom Masses of elements to work out the mass of the molecule they make up. To work out the Relative Formula Mass you simply add up the Relative Atomic Masses of each atom in the molecule. Sodium chloride (NaCl) has 1 sodium atom and 1 chlorine atom. RAM for Na is 23 and RAM for Cl is 35,5. Then, the formula mass is 58,5 atomic mass unit. Water (H2O) has 2 hydrogen atoms and 1 oxygen atom. RAM for H is 1 and for O is 16, then the formula mass for water is 18 atomic mass unit.

1.4 The mole In Chemistry, the mole is the amount of substance. The abbreviation is mol. To find the mass of 1 mole of any substance, work out the relative formula mass and attach the units in grams. In one mole there are 6,022·10²³ units of any substance. This number is called Avogadro´s number. WORKED EXAMPLE 1 Work out the mass of one mole of oxygen gas, O2. If we look at the Periodic Table, we take down the RAM of O = 16. We work out that RFM O2 = 32 atomic mass unit. Then, 1 mole of O2 weighs 32 grams and it contains 6,022·10²³ molecules of oxygen. To interconvert between a mass in grams and a number of moles for a given substance, you can use the formula :

number of moles=

mass ( g ) mass of 1 mole (g )

Then, in 1kg of oxygen gas, there are 1.000 g / 32 g = 31,25 mol of oxygen molecules, and it represents the huge amount of 1,88 ·10²⁵molecules of oxygen. Since each one has two atoms, the amount is doubled, then 3,76 ·10²⁵ oxygen atoms. WORKED EXAMPLE 2 Find the mass of 0,2 moles of calcium carbonate (CaCO3). RAM´s are : Ca = 40, C =12, O = 16. First, find the RFM for CaCO3 : 40 + 12 + 3· 16 = 100 atomic mass unit. Mass of 1 mole of CaCO3 = 100 grams. Mass = number of moles x mass of 1 mole = 0,2 · 100 = 20 grams

2 UNIT Chemical bonding WORKED EXAMPLE 3 Work out the amount of substance in 800 g of CO2. How many molecules are there? RAM´s are : C = 12, O = 16. Mass of 1 mole of CO2 = 44 grams. Number of moles = 800 grams / 44 grams per mole = 18,18 mol. In 18,18 moles there are 18,18 · 6,022 ·10²³ = 1,09 ·10²³ molecules.

2. Calculations from equations In laboratories and in industry it is neccesary to know the masses and the volumes of the reactants and products involved in chemical reactions.

2.1 Calculations involving only masses When limestone CaCO3 is heated, calcium oxide is formed. Suppose you wanted to calculate the mass of calcium oxide produced by heating 25 g of limestone. RAM´s are:

Ca = 40; C = 12; O = 16. First write down the equation: CaCO3 → CaO + CO2 Interpret the equation in terms of moles. One mole of CaCO3 produces one mole of CaO and one mole of CO2. Calculate the formula mass of CaCO3 and CaO. The mass of CO2 is not useful in this case, so you do not need to calculate it. Mass CaCO3 = 40 + 12 + 3·16 = 100 atomic mass units. Mass of 1 mole = 100 grams. Mass CaO = 40 + 16 = 56 atomic mass units. Mass of 1 mole = 56 grams. Convert moles to masses. From 100 g of CaCO3, 56 g of CaO are produced. But we want to know the mass of CaO from 25 g of CaCO3. Instead of doing a simple proportion, we use conversion factors in order to understand more complicated equations:

25 g CaCO3

1 molCaCO 3 1 mol CaO

56 g CaO

100 g CaCO 3 1 mol CaCO 3 1 mol CaO

=14 g CaO

The answer is that from 25 g of CaCO3, 14 grams of CaO are obtained.

2.2 Calculations involving gas volumes As a rule, gas volumes are measured in cubic centimetres (cm³), cubic decimetres (dm³) or litres (l).

1 litre = 1 dm³ = 1.000 cm³

Chemical bonding UNIT

Thanks to Avogadro we know this fact about the gases: STP stands for standard temperature and pressure, and it means a temperature of 0ºC and 1 atmosphere of pressure. One mole of any gas at stp has a volume of 22,4 litres. This is called the molar volume. Let us carry out the following activity to explain it. Calculate the volume of carbon dioxide produced at 0ºC and 1 atm when excess of dilute HCl is added to 1 g of CaCO³. RAM´s are: C = 12; Ca = 40; H = 1. First write the equation and make sure it is balanced.

CaCO3 + 2 HCl → CaCl2 + CO2 + H2O Interpret the equation in terms of moles: 1 mole of CaCO3 produces 1 mole of CO2. Then, calculate the formula mass of CaCO3 and CO2. Mass CaCO3 = 40 + 12 + 3 ·16 = 100 atomic mass units. Mass CO2 = 12 + 16 = 44 atomic mass units. Convert moles to masses: 100 g of CaCO³ produce 44 g of CO2. Therefore from 1 g:

1 g CaCO 3

1 mol CaCO 3 1 mol CO2 22,4lCO 2 100 g CaCO3 1 mol CaCO3 1 mol CO2

=0,224 lCO2

2.3 Calculations involving solutions Molarity is probably the most commonly used unit of concentration. It is the number of moles of solute per litre of solution, not necessarily the same as the volume of solvent. The molarity of a solution made when water is added to 10 g CaCl2 to make 100 mL of solution is calculated as follows: 10 g CaCl2 / (111 g CaCl2 per mol CaCl2) = 0.10 mol CaCl2 100 mL x 1 L / 1.000 mL = 0.10 L molarity = 0.10 mol / 0.10 L molarity = 1.0 M WORKED EXAMPLE 4

We dissolve 45 g of glucose (C6H12O6) in water to complete 500 cm³ of solution. Which is its molarity? n = m/M = 45/180 = 0,25 mol M = n/V = 0,25/0,5 = 0,5 M Where, n = number of moles; m = mass in grams; M = formula mass; V = volume.

2 UNIT Chemical bonding

3. Energy in chemical reactions All chemical reactions are accompanied by a change in energy. Some reactions release energy to their surroundings usually in the form of heat, and are called exothermic. For example, sodium and chlorine react so violently that flames can be seen as the exothermic reaction gives off heat. On the other hand, some reactions need to absorb heat from their surroundings to proceed. These reactions are called endothermic. A good example of an endothermic reaction is that which takes place inside of an instant "cold pack." Commercial cold packs usually consist of two compounds, urea and ammonium chloride in separate containers within a plastic bag. When the bag is bent and the inside containers are broken, the two compounds mix together and begin to react. Because the reaction is endothermic, it absorbs heat from the surrounding environment and the bag gets cold.

4. Rates of reaction Why do some reactions, which by their free energy values should be spontaneous and far from equilibrium, sit inert and unreactive for years, whereas other reactions go with explosive rapidity? The decomposition of NO to nitrogen and oxygen is thermodynamically spontaneous, so why do we have photochemical smog from oxides of nitrogen? If all combustions with oxygen liberate free energy, and the atmosphere is full of oxygen, then why does not everything that is potentially flammable burn at once, including ourselves? The answer is that these decompositions and combustions, although thermodynamically spontaneous, occur at miniscule rates at room temperature. The rates of chemical reactions and the factors that affect them are the subjects of this chapter.

Chemical bonding UNIT

4.1 Concentration of Reactants A higher concentration of reactants leads to more effective collisions per unit time, which leads to an increasing reaction rate, except for zero order reactions. Similarly, a higher concentation of products tends to be associated with a lower reaction rate. Use the partial pressure of reactants in a gaseous state as a measure of their concentration.

4.2 Temperature Usually, an increase in temperature is accompanied by an increase in the reaction rate. Temperature is a measure of the kinetic energy of a system, so higher temperature implies higher average kinetic energy of molecules and more collisions per unit of time.

4.3 Medium The rate of a chemical reaction depends on the medium in which the reaction occurs. It may make a difference whether a medium is aqueous or organic; polar or nonpolar; or liquid, solid, or gaseous.

4.4 Presence of Catalysts and Competitors Catalysts e.g., enzymes lower the activation energy of a chemical reaction and increase the rate of a chemical reaction without being consumed in the process. Catalysts work by increasing the frequency of collisions between reactants, altering the orientation of reactants so that more collisions are effective, reducing intramolecular bonding within reactant molecules, or donating electron density to the reactants.

5. Acid base reactions Scientists use something called the pH scale to measure how acidic or basic a liquid is. Although there may be many types of ions in a solution, pH focuses on concentrations of hydrogen ions (H+) and hydroxide ions (OHË&#x2030;). The scale goes from values very close to 0 through 14. Distilled water is 7, right in the middle. Acids are found between a number very close to 0 and 7. Bases are from 7 to 14. Most of the liquids you find every day have a pH near 7. Neutralisation reaction When an acid and a base are placed together, they react to neutralize the acid and base properties, producing a salt. The H+ cation of the acid combines with the OHË&#x2030;anion of the base to form water. The compound formed by the cation of the base and the anion of the acid is called a salt. The combination of hydrochloric acid and sodium hydroxide produces common table salt, NaCl:

The word salt is a general term which applies to the products of all such acid-base reactions. 25

2 UNIT Chemical bonding

6. Redox reactions Redox reactions, or oxidation-reduction reactions, have a number of similarities to acidbase reactions. Fundamentally, redox reactions are a family of reactions that are concerned with the transfer of electrons between species. Like acid-base reactions, redox reactions are a matched set ; this means that you do not have an oxidation reaction without a reduction reaction happening at the same time. Oxidation refers to the loss of electrons, while reduction refers to the gain of electrons. Each reaction by itself is called a "half-reaction", simply because we need two halfreactions to form a whole reaction. In notating redox reactions, chemists typically write out the electrons explicitly:

Cu (s) → Cu²+ (aq) + 2 eThis half-reaction says that we have solid copper with no charge being oxidized, losing electrons, to form a copper ion with a plus 2 charge. Notice that, like the stoichiometry notation, we have a "balance" between both sides of the reaction. We have one copper atom on both sides, and the charges balance as well. The symbol "e-" represents a free electron with a negative charge that can now go out and reduce some other species, such as in the half-reaction: 2 Ag+ (aq) + 2 eˉ → 2 Ag (s) Here, two silver ions, silver with a positive charge, are being reduced through the addition of two electrons to form solid silver. The abbreviations "aq" and "s" mean aqueous and solid, respectively. We can now combine the two half-reactions to form a redox equation:

Cu (s) → Cu²+ (aq) + 2 eˉ 2 Ag+ (aq) + 2 eˉ → 2 Ag (s) Cu (s) + 2 Ag+ (aq) + 2 eˉ → Cu²+ (aq) + 2 Ag (s) + 2 eˉ Cu (s) + 2 Ag+ (aq) → Cu²+ (aq) + 2 Ag (s)

Chemical bonding UNIT

ACTIVITY 1. Say which is the oxidant and which one is the reducing agent in the reaction :

Fe + S → FeS Sulphur is a non metal, so it gains electrons whereas iron is a metal, then it loses electrons. Sulphur is the oxidant and iron is the reducing agent. After the reaction the iron will be oxidised by the sulphur and this will be reduced by the iron. As a result, iron is converted in iron cation and sulphur in sulphide anion. The charges are Fe ²+ and S²ˉ . Both will combine to form iron sulphide, FeS, with a net charge of zero, so that it is neutral.

Questions 1. Which is the difference between physical and chemical change ? 2. Balance the following equations : a. N2 + H2 → NH3 b. H2 + I2 → HI c. C + O2 → CO 3. Calculate the formula mass for NH3. How many molecules are there in 1 mole ? 4. Aluminium reacts with hydrochloric acid as follows : 2 Al + 6 HCl → 2 AlCl3 + 3 H2 If we measure 20 g of aluminium : a. How many grams of AlCl3 will appear ? b. How many litres of hydrogen will obtain at s.t.p. ? 5. What is an endothermic reaction ? and an exothermic ? 6. An iron nail is introduced in a silver nitrate solution. a. Write down the chemical reaction b. Which substance is oxidised ? 7. Given the chemical equation C + O2 → CO2 , if 240 g are burnt, how many grams of CO2 are obtained ? How many litres of CO2 at s. t.p. are produced? 8. One hundred g of potassium chlorate are decomposed by heating. The equation is:

KClO3

→

KCl + O2

2 UNIT Chemical bonding Questions a. Balance the equation and work out the grams of potassium chloride obtained. b. Which volume of oxygen is obtained at s.t.p.? 9. Write down the ions in the following acids and bases: a. NaOH b. HCl

c. KOH

d. Al (OH)3

e. H2S

10. Calculate the molarity of a solution made up of 21 grams of sodium chloride in one litre of distilled water. RAMÂ´s are: Na = 23, Cl = 35,5. 11. Calculate the number of moles of water in a 5 litre container. RAMÂ´s are: H = 1; O= 16. 12.Which is the molarity of a solution made up of 5 g of common salt in water up to a volume of 100 ml? 13. We would like to prepare 250 ml of an aqueous solution of calcium chloride in water 1,5 M. Calculate the number of grams of solute we need for it. 14. We have 15 ml of a solution of calcium hydroxide 0,5 M. Calculate the number of moles and the grams of calcium hydroxide we have got in the solution. 15. Work out the volume of a solution of hydrochloric acid 1,25 M that we need to take in order to obtain 0,5 mol of acid. 16. Explain the proceedings you would carry out to prepare an aqueous solution of sodium hydroxide 2 M in the laboratory. RAM's are Na = 23; O = 16; H = 1. 17. You are asked to prepare a 20% solution of sodium chloride and you think that this means that there are 20 g of salt into 100 g of water. If you proceed this way, which will the concentration of the solution be? 18. Calculate the molarity of a solution of sodium hydroxide if it contains 80 grams of solute in 2.000 ml of solution. RAM's are Na = 23; O = 16; H = 1. 19. We need a solution of nitric acid (HNO3) to prepare B3 vitamin, and it must be 0,1 M. How many grams of pure HNO3 are needed to prepare the solution? RAM's are: H = 1; N = 14; O = 16. 20. Potassium hydroxide is used to make soap. A solution of potassium hydroxide contains 112 g of solute per litre. a. Calculate the molarity b. If the density of the solution is 1,1 g/ml, expres the concentration in mass percentage.

Chemical bonding UNIT

Questions 21. Hydrochloric acid is inside humans digestive apparatus. Calculate the amount of hydrochloric acid needed to prepare 1 litre of solution: 1 M; 0,5 M; 2 M. 22. Hydrochloric acid reacts with calcium hydroxide to yield calcium chloride and water. We have 50 ml of a 0,5 M solution of calcium hydroxide. Work out: a. The mass of calcium chloride that it is obtained b. The volume of the solution of hydrochloric acid 0,25 M needed to react with the calcium hydroxide. 23. Say if the following changes of state are exothermic or endothermic: fusion, vaporisation, sublimation, condensation, solidification. Explain what happens. Is a different way to change the state of matter? 24. Titanium is manufacturated by heating titanium (IV) chloride with sodium.

TiCl4 (g) + Na (s) Ti (s) + NaCl (s) What mass of sodium is required to produce 85 g of titanium? RAMs: Na = 23; Ti = 48. 25. How many water molecules are there in one drop of water? Assume one drop of water is 0,05 cm³ and the density of water is 1 g/cm³. RAM´s: H = 1; O = 16. 26. Calculate the relative formula mass of the following compounds. RAM´s: H = 1; C= 12; N = 14; O = 16; Na = 23; S = 32; Ca = 40; Cr = 52; Fe = 56 a. CO2 b. (NH4) SO4 c. Na2CO3. 10 H2O d. Cr2(SO4)3 e. (NH4)2 SO4. FeSO4. 6 H2O 27. Find the percentage by mass of the named substance in each of the following examples. RAM´s: H = 1; C = 12; O = 16; Mg = 24; S = 32 a. Carbon in propane, C3H8 b. Water in magnesium sulfate crystals, MgSO4. 7 H2O 28. Work out the percentage of nitrogen in each of the following substances (all used as nitrogen fertilisers). RAM´s: H = 1; C = 12; N = 14; O = 16; S = 32; K = 39 a. urea, CO (NH2)2 b. potassium nitrate, KNO3 c. ammonium nitrate, NH4NO3 d. ammonium sulfate, (NH4)2SO4

2 UNIT Chemical bonding Questions 29. Work out the mass of the following substances. RAM´s: H = 1;C = 12; N = 14; O = 16; Na = 23; Pb = 207 a. 1 mole of lead (II) nitrate, Pb (NO3)2 b. 4,3 moles of methane, CH4 c. 0,24 moles of sodium carbonate crystals, Na2CO3·10 H2O 30. How many moles are represented by each of the following . RAM´s H = 1; O = 16; S = 32; Fe = 56; Cu = 64 a. 50 g of copper (II) sulfate crystals, CuSO4.5 H2O b. 1 tonne of iron, Fe c. 0,032 g of sulfur dioxide, SO2 31. Some more questions about converting between moles and grams of a substance RAM´s: H = 1; O = 16; Na = 23; Cl = 35,5; Ca = 40; Cu = 64 a. What is the mass of 4 mol of sodium chloride, NaCl? b. How many moles is 37 g of calcium, Ca? c. How many moles is 1 kg of calcium hydroxide, Ca(OH)2? d. What is the mass of 0,125 mol of copper (II) oxide, CuO? e. 0,1 mol of a substance weighed 4 g. What is the mass of 1 mole? f. 0,004 mol of a substance weighed 1 g. What is the relative formula mass of the compound?