Chemistry
1
ATOMIC STRUCTURE AND THE PERIODIC TABLE IDEAS YOU HAVE MET BEFORE:
ELEMENTS, MIXTURES AND COMPOUNDS
WHAT MODEL DO WE USE TO REPRESENT AN ATOM?
• Mixtures can be separated easily by filtering and other ways. • Elements cannot be broken down by chemical means. • Compounds are made from elements chemically combined.
• Electrons fill the shells around the nucleus in set pattern orders. • Protons and neutrons make up the nucleus. • Electrons can be lost from or gained into the outer shell.
ATOMS AND THEIR STRUCTURE
HOW DID THE MODEL OF THE ATOM DEVELOP?
• Electrons have a negative charge. • Atoms have a nucleus with a positive charge. • Electrons orbit the nucleus in shells.
C
Li
• Atoms used to be thought of as small unbreakable spheres. • Experiments led to ideas of atoms with a nucleus and electrons. • Electrons in shells and the discovery of the neutron came later.
SOME ELEMENTS AND THEIR COMPOUNDS
WHY CAN WE USE CARBON DATING?
• Helium is unreactive and used in balloons. • Sodium chloride is used to flavour and preserve food. • Chlorine is used to kill bacteria in swimming pools.
• Atoms of an element always have the same number of protons. • They do not always have the same numbers of neutrons. • Elements exist as different isotopes.
C
WHY IS HELIUM SO UNREACTIVE AND SODIUM SO REACTIVE?
• Gold, silver and platinum are precious metals. • Mercury is a liquid metal. • Zinc, copper and iron are used to make many useful objects.
• The outer shell of helium can take no more electrons. • The outer shell of sodium has 1 electron which it needs to lose. • Metals need to lose electrons, non-metals do not.
METALS AND NON-METALS
WHAT DO TRANSITION METAL COMPOUND SOLUTIONS LOOK LIKE?
AQA GCSE Chemistry: Student Book
C
Spot the difference in these isotopes
METALS IN THE PERIODIC TABLE
• Gold, iron, copper and lead are metals known for centuries. • Oxygen and nitrogen are gases of the air. • Sulfur is a yellow non-metal.
12
IN THIS CHAPTER YOU WILL FIND OUT ABOUT:
• Transition metals are harder and stronger than Group 1 metals. • Transition metals are often used as catalysts. • Transition metal compounds often form coloured solutions.
Atomic structure and the periodic table
13
Chemistry
Check your progress
Worked example
You should be able to:
■■describe three main types of bonding
1
➞
■■represent an ionic bond with a diagram
■■explain how electrons are
used in the three types of bonding
■■draw a dot and cross
diagram for ionic compounds
➞
properties are linked
➞
■■identify single bonds in
draw dot and cross diagrams deduce molecular formula from ➞ ■■for small molecules ➞ ■■models and diagrams
■■describe that metals form giant structures
➞
■■explain how metal ions are held together
➞
relate their melting points to explain when ionic compounds ➞ ■■forces between ions ➞ ■■can conduct electricity
■■identify small molecules from formulae
■■recognise giant covalent
structures from diagrams
➞
■■identify polymers from their unit formula
explain the properties of ➞ ■■giant covalent structures
■■identify metal elements and ➞ ■■describe the purpose of a metal alloys lead–tin alloy ■■explain how the properties relate to the bonding in diamond
■■describe the structure of graphene
➞ ➞
■■explain why diamond differs from graphite
■■explain the structure and
AQA GCSE Chemistry: Student Book
uses of fullerenes
➞
2
X
Y
Explain why both diamond and silicon dioxide are hard with high melting points.
3
Draw diagrams to show how substances change from solids to liquids.
The particles move faster when they are heated and break away from the solid structure to move more freely.
to the bulk properties of a substance
This answer shows both diagrams and has an added explanation.
explain why alloys have ➞ ■■different properties to elements
➞
This answer explains what the bonding is like but needs to be linked to the energy required to break the bonds for melting to happen.
■■relate the intermolecular forces
explain the strength of covalent ➞ ■■bonds
➞
X has lost an electron so the 1+ charge is missing. Y has gained an electron so the 1– charge is missing. Both charges should be at the top right outside the brackets. Electrons should be kept in pairs in a circle.
They have covalent bonds that act in all directions.
enabled by the delocalisation of electrons
■■describe the properties of ionic compounds
bD raw the dot and cross diagram for the resulting compound. Use X to represent the metal and Y to represent the nonmetal. Show the outer shell only.
■■explain how metallic bonding is
explain the changes of state use state symbols in chemical ➞ ■■ ➞ ■■equations
The answer ionic is correct.
a) metallic b) ionic c) covalent d) giant
work out the empirical formula ➞ ■■of an ionic compound
■■use data to predict the states of substances
ions of metal and non-metals from the group number of the element
explain the limitations of ➞ ■■diagrams and models
molecules and structure
a Identify the type of bonding.
■■work out the charge on the
■■identify ionic compounds from structures
92
➞
■■explain how bonding and
An element from Group 1, X, bonds with an element from Group 7, Y.
solid
liquid
■■explain the similarity of graphite to metals
4
■■compare ‘nano’ dimensions
to dimensions of atoms and molecules
Fill in the missing data in the table. Substance
Metal
Small molecule
Giant covalent
Ionic
Melting point
High
Low
High
High
Conducts electricity
Yes
No
Yes
Yes when melted, no when solid
The metal, small molecule and ionic columns are all correct. The student has correctly stated the difference in conductivity between an ionic solid and liquid. Giant covalent structures do not normally conduct electricity. Graphite is an exception and this should be explained.
Worked example
93
Chemistry
Concentration of solutions Learning objectives: • relate mass, volume and concentration • calculate the mass of solute in solution • relate concentration in mol/dm3 to mass and volume.
solute solvent solution concentration
Many chemical reactions take place in solutions and often the concentration needs to be known. If a volume of solvent is used, say 100 cm3, the number of particles of solute in the solvent is lower if the solution is more dilute and is higher if the solution is more concentrated.
or concentration =
Figure 3.15 The concentration increases as the number of solute particles in a fixed volume increases.
mass of solute volume
The standard unit of concentration used in the laboratory is: grams per dm3 so the units are then g/dm3. 1 dm3 is 1000 cm3 so the concentration of X is 25 g/dm3 1
Put the following solutions into order with the most dilute first : a) 20 g/100 cm3 b) 20 g/1000cm3 c) 8 g / 50cm3
Key information
2
Calculate the concentration of the following solutions in g/dm3. a) 3.2 g in 100 cm3 b) 3.2 g in 250 cm3 c) 6.4 g in 500 cm3
1 dm3 is commonly called 1 litre but we use 1 dm3 when using units in scientific contexts.
Calculating from concentrations A solution has a concentration of 6 g/dm . What is the mass of solute that is in 100 cm3 of this solution? mass Concentration = volume 6 Concentration = g/dm3 and Concentration = X g/100 cm3 1000 X 6 6 × 100 Therefore: = so X = = 0.6 g 100 1000 1000 3
112
AQA GCSE Chemistry: Student Book
4
A solution has a concentration of 5.4 g/100 cm . Calculate the mass of solute dissolved in 35 cm3 of solution.
3.8
3
Higher TIER Only
Using moles in concentrations
Checking the correct units
This is because the concentration of a solution can be measured as the mass of solute dissolved per specified volume of solution
A solution has a concentration of 4.2 g/dm3. Calculate the mass of solute dissolved in 250 cm3 of solution.
Key words
We have learned before that a solution forms when a solute is dissolved in a solvent and that solutions can be dilute or concentrated. Dilutions are often critical, such as making the correct formulations of medicines or baby milk formula. How do we make sure we have the correct concentration?
If a 100cm3 solution contains 2.5g of solute X, then the concentration is 2.5 g/100 cm3.
3
Many solutions in used in chemical reactions have solutes expressed as amounts of substances, measured in moles. The concentration of these solutions can then be measured in mol/dm3. A relationship exists between the amount in moles, concentration in mol/ dm3 and volume in dm3: concentration = amount in moles ÷ volume amount in moles = concentration × volume volume = amount in moles ÷ concentration. A way to remember this is by using a formula triangle, as in Figure 3.16. Don’t forget these formula triangles are only an aid to help you check your answer in a test. You should work from the first principles of the equation and carry out the same operations on both sides of the equation in order to calculate a new value in the relationship. For example: A solution has a concentration of 0.2 mol/dm3 To find the amount of moles in 500 cm3 of a 0.2 mol/dm3 solution we need first to make sure that the units match. The volume in cm3 converts into dm3 so: 500 cm3 = 0.5 dm3. Then we use the formula: amount in moles = concentration × volume = 0.2 mol/dm3 × 0.5 dm3 = 0.1 mol 5
A solution with a concentration of 2 mol/dm3 contains 1 mole of solute. Calculate the volume, in dm3, of the solution.
6
A solution of 0.18 moles has a volume of 0.6 dm3. Calculate the concentration in mol/dm3
7
Calculate the amount of solute, in moles in 500 cm3 of a solution with a concentration of 3 mol/dm3
8
4.90g of H2SO4 were dissolved in water to make 200 cm3 of solution. Calculate the resulting concentration of the acid solution in mol/dm3
9
0.025 moles of HNO3 were dissolved in water to make 125cm of solution. Calculate:
amount in moles concentration x volume
Figure 3.16 Cover the quantity needed to find the formula to use.
You’ll need the relative atomic masses of the following elements for some of these questions: H = 1 N = 14 O = 16
b the concentration of HNO3 in g/dm3 10
Remember Avogadro’s number? Concentrations are expressed in mol/dm3 because chemical equations are used to express the ratios of reactants and products in moles. The number of particles in 1 mole is always the same.
Remember!
3
a the concentration of HNO3 in mol/dm3
Did you know?
The concentration of HCl is 18.25 g/dm3
S = 32 CI = 35.5
a Calculate the concentration of HCl in mol/dm
3
b Work out the number of molecules of HCl per dm3. The Avogadro number is 6.02 × 1023 mol−1 Google search: 'calculating concentrations from moles'
113
Chemistry
REQUIRED PRACTICAL
REQUIRED PRACTICAL
The readings on the burette are taken by looking directly in line with the meniscus of the liquid. They are noted down in a recording table, measured in cm3. The order of recordings are:
Finding the reacting volumes of solutions of acid and alkali by titration
titration concentration acid alkali
•
describe how safety, the correct manipulation of apparatus and the accuracy of measurements are managed when titrations are carried out make and record observations and accurate measurements using burettes calculate the concentration of a solution from the concentration and volume of another.
Using instruments for measuring volumes accurately is an important skill for many scientists. The ability to calculate the concentration of solutions in both mol/dm3 and g/dm3 from the reacting volumes is a skill necessary to achieve the higher grades. Can you recognise the neutralising volume from a titration curve when measuring pH?
These pages are designed to help you think about aspects of the investigation rather than to guide you through it step by step.
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Volume of nitric acid 1st reading cm
alkali, pH high
13 12
acid has reacted with some of the alkali pH falls
11 10 9 7
Rough volume
More accurate trials 1st trial
2nd trial
burette
20 30 40
acid
50
solution of alkali
reaction mixture and indicator
Figure 4.38
3rd trial
1.5
1.8
2.4
15.2
2nd reading cm3
28.5
28.4
29.1
41.8
Volume used cm3
27.0
26.6
26.7
26.6
3
5
Calculate the mean titre.
6
Jo calculated the mean titre as 26.725 cm3. Explain why this is incorrect.
7
The mean titre is 26.6 cm3 and not 26.63 cm3. Explain why.
A number of different skills are needed to carry out a titration with safety and great accuracy. The volumes of the solutions used are either measured as a fixed volume or a variable volume. The point of neutralisation is determined by use of an indicator.
3
HIGHER TIER ONLY
all the alkali is neutralised pH = 7
6 5 4
Determining the concentration of one of the solutions in mol/dm3 and g/dm3.
excess acid pH is below 7
• First write the equation of acid + alkali. • Check that it is an equation of the type: HA + BOH → BA + H2O. This means that 1 mole of acid reacts with 1 mole of alkali.
2 1 0
4
Think about these questions: 1
Name the type of reaction that takes place between the acid and the alkali.
2
Name the piece of apparatus that is needed to measure a fixed volume of the alkali.
3
Identify the piece of apparatus that is needed to measure the variable volume of the acid.
4
Suggest two safety measures that should be used during a titration.
8
12
16
20
24
28
32
36
40
44
volume of acid in cm3 end point
48
Jo had used acid and alkali that reacted 1 mole:1 mole ratio. 8
Figure 4.37
DID YOU KNOW? Acid is added from the burette to neutralise the alkali, seen by the indicator colour change.
This number of moles exactly neutralised the same number of moles of Jo’s acid. 9
DID YOU KNOW? If the reaction is of the type: H2A + 2BOH → B2A + H2O then 1 mole of acid reacts with 2 moles of alkali and you will need to take that into account in your calculation.
The concentration of Jo’s alkali in mol/dm3 is 0.12. Calculate the number of moles in 25 cm3.
This same number of moles was contained in Jo’s mean titre of acid. Calculate how many moles this would be in 1000 cm3.
This is the concentration of the acid in mol/dm3. 10
Jo used HNO3. Work out the molar mass of this acid. This will be the mass that is dissolved in 1 mol/dm3.
Obtaining accurate measurements in a titration
11
Calculate the concentration of Jo’s nitric acid in g/dm3.
When acid is added to the burette it does not need to be filled right up to the 0.0 cm3 mark as the volume used can be calculated, but the section below the tap needs to be filled.
12
In a titration, 10.0 cm3 of 0.250 mol/dm3 NaOH was titrated with HCl. The mean titre of HCl was 11.2 cm3. Calculate the concentration of HCl in mol/dm3.
AQA GCSE Chemistry: Student Book
pipette
8
Managing titrations safely
150
0 10
Jo had recorded these results and calculated the mean titre:
pH
•
• The final reading at neutralisation is taken and the ‘rough’ volume of acid used is calculated. • The procedure is repeated again but towards the final expected volume, the acid is added drop by drop until the colour just changes. • The ‘titre’ is noted and three more titres are obtained with volume reading within 0.1 cm3.
KEY WORDS
Learning objectives: •
4.10
KEY SKILL The answer to question 9 was the number of moles in 1 dm3 of Jo’s acid. The molar mass of Jo’s acid was calculated in question 10. You can work out the mass of acid by rearranging the equation: Moles = mass/molar mass.
Google search: ‘using titrations to measure concentration’
151
Chemistry
End of chapter questions Getting started 1
Heavy fuel oil exits near the bottom of the fractionating column during fractional distillation of crude oil. Which temperature range is this? a 80–120 °C b 110–190 °C c 130–230 °C d 240–340 °C
1 Mark
2
Define the term hydrocarbon.
1 Mark
3
Identify the products for complete and incomplete combustion of an alkane.
2 Marks
4
Identify the alkene. a C 3H 8 b C 3H 6 c C 3H 8O d C 3H 6O2
5 6 7
16
Describe the structure of one strand of DNA.
2 Marks
17
Write a balanced equation for the incomplete combustion of heptane, C7H16.
2 Marks
18
The following table contains the results of some chemical tests on organic compounds. Determine the most likely functional group for D, E and F. Justify your answer.
1 Mark
Hexanoic acid is a carboxylic acid. Predict what colour it will turn universal indicator solution.
1 Mark
Monomer X is an alkene. State the type of polymerisation it undergoes and draw a short section of the polymer.
2 Marks
Match the monomer to the polymer.
2 Marks
D
E
F
Addition of bromine water
Remains orange
Remains orange
Turns colourless
Addition of aqueous sodium carbonate
Fizzes
No reaction
No reaction
Addition of sodium
Fizzes
Fizzes
No reaction
Most demanding 19
Compare the processes of condensation and addition polymerisation.
20
A hydrocarbon has the formula C 4H6. It decolourises bromine water.
glucose
polyester
a Identify the functional group present.
amino acid
cellulose
b Draw a structure for C 4H6. 21
acid and alcohol
protein
Going further 8
Write the word equation for the addition of hydrogen to butene.
9
Which of the following is an alcohol? a CH3CH2CH2OH b CH3CH2COOH
1 Mark
4 Marks
2 Marks
2 Marks
Substance D is a hydrocarbon containing six carbons. When heated with a catalyst it forms E which has two carbons and one other compound, F. E reacts to form G. When sodium is added to G, the gas evolved pops when a lighted splint is inserted. a Identify D, E, F and G, writing equations where relevant. b Describe any additional reactions that E undergoes.
6 Marks
Total: 40 Marks
c CH3CH2CH3 d CH3CHCH2
10
Explain the trend of boiling points of the alkane series.
11
Cracking is carried out on fractions produced from crude oil.
1 Mark 2 Marks
a Explain the purpose of cracking. b Write a balanced equation for the cracking of heptane, C7H16 . 12
2 Marks
Butene, C 4H8, is an alkene. a Describe how butene could be converted to an alcohol.
13
b Explain how you could make butene saturated.
2 Marks
The viscosity of three unknown alkanes was compared: D: 1, E: 1.2 F: 1.6. The higher the number the greater the viscosity. Explain which has the highest boiling point.
2 Marks
More challenging
260
14
Work out the formula for an alkane with 20 carbons.
1 Mark
15
Explain the type of polymerisation possible between a carboxylic acid and an alcohol.
1 Mark
AQA GCSE Chemistry: Student Book
End of chapter questions
261
radon astatine
Ds
111 109 108 107 106 105 89
Ac*
104
rutherfordium actinium
88
Ra
radium
87
Fr
francium
55 [223]
Ba Cs
cesium
Rf
72 [261] 57 [227] 56 [226]
Db
Sg
Bh
bohrium seaborgium dubnium
Hs
hassium
Mt
meitnerium
110
Rg
darmstadtium roentgenium
78 [271] 77 [268] 76 [277] 74 [266] 73 [262]
Re W Ta
75 [264]
tungsten tantalum
La*
Hf
hafnium lanthanum barium
40 178 39 139 38 137 37 133
Y
rhenium
43 186 42 184 41 181
Tc Mo Nb
niobium
Zr
zirconium yttrium
Sr
strontium
Rb
rubidium
* The Lanthanides (atomic numbers 58–71) and the Actinides (atomic numbers 90–103) have been omitted. Relative atomic masses for Cu and Cl have not been rounded to the nearest whole number.
84 83 82 81
acid rain rain which has been made more acidic by pollutant gases
80 79 [272]
Po
Elements with atomic numbers 112–116 have been reported but not fully authenticated
85
At
polonium
Bi
bismuth lead
Pb Tl
thallium
Hg
mercury gold
Au Os
osmium
Ir
Pt
platinum iridium
47 197 46 195 45 192
acids dissolve in water to produce solutions with a pH of less than 7
technetium
44 190
Rh
A
molybdenum
Ru
ruthenium
Glossary
86
54 [222] 53 [210] 52 [209] 51 209 49 204 48 201
50 207
I Te
tellurium
Sb
antimony
Pd
palladium rhodium
silver
Ag
Cd
In
indium cadmium
tin
Sn
Rn
xenon iodine
Xe
36 131 35 127 34 128 33 122 31 115 30 112 29 108 28 106 27 103 26 101 25 [98] 24 96 23 93 22 91 21 89 20 88 19 85
V Ti
titanium
K
potassium
12 40
Ca
calcium
Sc
scandium
48 45
Na
11 39
magnesium
Mg
AQA GCSE Chemistry: Student Book
sodium
4 24 3 23
Be
beryllium
Li
lithium
9 7
2 1
354
32 119
Ga
cobalt iron
Fe Cr
atomic (proton) number
name
atomic symbol
relative atomic mass
Key
55
Mn
manganese chromium vanadium
56 52 51
1
1
H
hydrogen
59
Co
59
Ni
nickel
Cu
copper
zinc
Zn
gallium
Ge
germanium
As
arsenic
Se
selenium
Br
Kr
krypton bromine
18 84 17 80 16 79 15 75 14 73 63.5
65
13 70
Cl
Ar
argon chlorine
S
sulfur
P
phosphorus
Si
silicon
Al
aluminum
10
40 35.5 32 31 28 27
neon
9 8 7 6
F B
11
boron
C
12
carbon
N
nitrogen
O
oxygen
16 14
6 5 4 3
5
fluorine
Ne
2 20 19
helium
4
He
7
0
Chemistry
Avogadro’s constant the number of atoms, molecules or ions in one mole of a given substance, and is 6.02 × 1023 per mole
B
activation energy the energy needed for a chemical reaction to happen
balanced symbol equation chemical equation written in chemical symbols showing the number of atoms on each side of the equation balance
addition polymer a very long molecule resulting from polymerisation, e.g. polythene
barium chloride chemical used to test for sulfates in aqueous solutions
aggregate gravel added to cement and sand to make concrete
base reacts with an acid to form a salt
alcohols family of organic compounds with the functional group –OH alkali metals the metals in Group 1 of the periodic table alkalis compounds which produce hydroxide ions in water alkanes a family of hydrocarbons with all single carbon-carbon covalent bonds and general formula CnH2n+2 alkenes a family of hydrocarbons with one double carbon-carbon bond and general formula CnH2n
battery two or more electrochemical cells joined together bioleaching process that uses bacteria to leach metal compounds from rocks biological catalyst molecules in cells of living organisms that speed up chemical reactions boiling point temperature at which the bulk of a liquid turns to vapour buckminsterfullerene a very stable spherical structure of 60 carbon atoms joined by covalent bonds (an allotrope of carbon)
allotropes different forms of the same element
C
alloy a mixture of a metal with one or more other metals or non-metals to change the properties of the metal
carbon an element present in all living things and forms a huge range of compounds with other elements
alpha particles radioactive particles which are helium nuclei – helium atoms without the electrons (they have a positive charge) amino acids small molecules from which proteins are built ammeter meter used in an electric circuit for measuring current anion ion with a negative charge; they move to the anode during electrolysis anode electrode in electrolysis with a positive charge aquifer underground layer of permeable rock or loose materials (gravel or silt) where groundwater is stored atom the basic ‘building block’ of an element, the smallest part of an element that can take part in a chemical reaction atom economy a measure of the amount of starting materials that become useful products atomic number the number of protons in the nucleus of an atom
carbon-14 radioactive isotope of carbon carbon dioxide (CO2) a greenhouse gas which is emitted into the atmosphere as a product of combustion carbon footprint the total amount of carbon dioxide and other greenhouse gases emitted over the full life cycle of a product, service or event. carboxylic acids family of organic compounds with the functional group –COOH catalyst a chemical that speeds up a reaction but is not used up by the reaction cathode the negative electrode in electrolysis cell an electrochemical cell is a unit that uses a chemical reaction to provide electricity charge(s) a property of matter, charge exists in two forms, positive and negative, which attract each other chemical properties the characteristic chemical reactions of substances chlorination addition of chlorine to water supplies to kill micro-organisms
Glossary
355