trinity academy · at trinity academy, we follow the ocr a course. this course has been long seen...

Trinity Academy

GCSE to A-Level Transition

Summer Work

A-Level Chemistry

(OCR Chemistry 7402)

Name: _____________________

Teachers: ___________________

Welcome to A-level Chemistry! I hope you will enjoy the challenges offered by this fascinating subject and

come to find it a rewarding and worthwhile experience. A-Level Chemistry is a fantastic course which

opens many doors for the future. Almost all higher education courses look favourably at Chemistry as a

subject, with it being a prerequisite for many STEM (Science, Technology, Engineering & Mathematics) and

medical course.

A-Level Chemistry at Trinity Academy

At Trinity Academy, we follow the OCR A course. This course has been long seen as an excellent all-round

course which gets a good balance of theory and practical, as well as introducing real life applications of

further level Chemistry. Below is an outline of the modules of the course and how they will be assessed.

In the first year of the course, you will study modules 1-4. These will be assessed internally in the form of

“Annual Assessment”. No externally examinations will take place at the end of the first year of teaching. The

second year of teaching will include modules 5-6 and will be assessed by three papers at the end of the

course. Further information can be found by looking at the course specification (found on the OCR website).

As an A-Level student at Trinity Academy you will also complete the “A-Level Practical Endorsement”. This

stand-alone qualification is reported separately to the A-Level grade and works on a pass/fail criteria. It is

not formally assessed; however students must demonstrate competency in a minimum of 12 different

practical activities. Many of these will be carried out in school; however, some will be completed at external

provides (e.g. university laboratories).

Chemistry A-Level is a difficult but prestigious course. It requires good mathematical skills, logic, strong work

ethic and an ability to understand abstract concepts. A-Level Chemistry does require significant amounts of

independent learning and reading and is not an A-Level you should be taking if you are unsure of whether

you want to study it. The point of this transition booklet and the taster sessions is to help you decide whether

this is the right subject for you.

Mr J Taylor (Head of Chemistry)

Summer Work Outline

This summer work has been designed to bridge the gap between GCSE Chemistry/Combined Science

(Chemistry) and AS level. This jump in demand regularly catches out students early in the course who still

continue to work with a KS4 mentality. For this reason, I have designed this summer work to stretch you

but also to give you a deeper understanding of the material you have studied at GCSE.

There is a lot of content that needs to be completed as part of this summer work ready for the first lesson

of teaching. It is important that you do not leave all the work until the last minute and that work is

staggered throughout your summer break. It is recommended that you follow the timeline below as a

guide. Each week, you should look to spend 1-2 hours working through this booklet.

Week Commencing Areas for learning

2nd July 2018 Section 1 – The structure of the Atom

9th July 2018 Section 2 – Formation of Ions

Section 3 – Intermolecular Bonding

16th July 2018 Section 4 – Bonding and Properties

23rd July 2018 Section 5 – Chemical Equations

30th July 2018 Section 6 – Inorganic Chemistry

6th August 2018 Section 7 – Organic Chemistry

13th August 2018 Section 8 – Chemical Reactions

Section 9 – Rates of Reactions

20th August 2018 Section 10 – Equilibria

Section 11 – Calculations

27th August 2018 Section 12 – Enthalpy

Section 13 – Investigating and Interpreting

You have been provided with a copy of the CGP “Head Start to A-Level Chemistry” book. This book will help

you work through the worksheets you are required to complete. It recaps all the crucial topics you have

studied as part of your GCSE course and introduces the basics of new concepts which are introduced at the

start of A-Level. The book uses study notes and examples, plus practice questions to test your

understanding. It’s the perfect way to hit the ground running at the start of the course, particularly for the

OCR exam board content.

Your workbook has been split into the same sections that the book has been split into for your benefit. This

will allow you to easily locate the appropriate section in the book that will assist you with the worksheets

set. It is important that you complete all the worksheets as this will allow you bridge the gap between

GCSE and A-Level.

As well as completing the worksheets provided, you are expected to complete a mind map for each section

of the book. You will have probably used mind maps during your GCSE revision, so this provides good

practice for you to continue your development. Your mind maps should be completed in time for the first

Chemistry lesson back in September. You mind maps can include colour, pictures, diagrams and any other

relevant content.

It is important that all the content in this booklet is completed in time for the new academic year.

Workbooks will be checked during the first week of the term. Poor/none completion of the material could

result in your place on the course being rescinded or conditions but onto your entry.

Helpful Links

https://chemrevise.org/

http://www.alevelchemistryrevision.co.uk/

http://www.physicsandmathstutor.com/

https://www.chemguide.co.uk/

https://www.savemyexams.co.uk/

https://revisionscience.com/

An additional tool for help is the Trinity Academy Science Twitter page (@TAScienceDept). By becoming a

follower of this page, you will be able to contact myself for additional help if needed. This page will also

provide you will interesting articles to read and study tips.

https://chemrevise.org/

http://www.alevelchemistryrevision.co.uk/

http://www.physicsandmathstutor.com/

https://www.chemguide.co.uk/

https://www.savemyexams.co.uk/

https://revisionscience.com/

Section 1

The Structure of the Atom

Work Title Completed

Spider Diagram

Electrons and Orbitals

Development of theories about atomic

structure

Isotopic Abundance

Isoelectric species

Spider Diagram: Section 1 – The Structure of the Atom

Atomic structure 2.3.

2.3. Electrons and orbitals

Aufbau’s principle states that “electrons fill orbitals starting with the lowest energy orbital first”

Hund’s rule states that “when filling a set of orbitals of identical energy, electrons are added with parallel spins to different orbitals rather than pairing two electrons in the same orbital”

For each of the elements, draw the electrons in the atom as you would have represented them at GCSE

level (1 mark) followed by an A-level representation (1 mark) and a short hand form of the electronic

configuration (1 mark);

e.g. silicon, Si;

1. oxygen, O;

2. calcium, Ca;

3. iron, Fe;

Give one limitation of the way you were taught to draw electrons in atoms at GCSE level (1 mark)

..............................................................................................................................................................................

1s

2s

3s

4s

3p

4p

↑↓ ↑↓

↑↓ ↑

2p

3d

↑↓ ↑↓

↑ ×

× × ×

× ×

× ×

× ×

×

×

× ×

Electronic configuration; [Ne] 3s2 3p2

↑↓

Electronic configuration; .......................

1s

2s

3s

4s

3p

4p

2p

3d

1s

2s

3s

4s

3p

4p

2p

3d

1s

2s

3s

4s

3p

4p

2p

3d

Electronic configuration; .......................

Electronic configuration; ......................


2.1. Development of theories about atomic structure

Our current understanding of atomic structure is a result of the discoveries of several scientists over

many years, each scientist adding to the model.

Complete the table below by adding the name of the scientist and the discovery made. Choose from the lists

below the table. (9 marks)

Approx. year of discovery

Scientist Addition made to our current understanding of atomic structure

1803 John Dalton Proposed that all matter is made up of tiny particles called atoms

1897

1911

1915

1924

1932

Scientists: Ernest Rutherford; Wolfgang Pauli; J. J. Thomson; James Chadwick; Niels Bohr

Discoveries; Proposed that the electrons orbit around the nucleus in orbits with a set size and energy

Discovered that atoms contain neutral particles called neutrons in their nucleus

Realised that atoms are divisible and contain very tiny, negatively charged particles called electrons

Discovered that an atom is made up of a nucleus and an extra-nuclear part. The central nucleus is

positively charged and the negative electrons revolve around this central nucleus.

Proposed the concept of electron spin

BONUS MARK: Which of the scientists listed above was a famous football goalkeeper in his country?

Analysis 10.1.2.

10.1.2. Isotopic abundance

1. The ratio of the different isotopes of certain elements can be used to identify objects from outer

space. By comparing the isotope patterns with samples known to originate on earth the scientists can

make recommendations as to the origins of unknown objects.

The mass spectrum opposite is of a

sample of chromium extracted from a

rock recently found in the Nevada

desert. Scientists believe it may be from

a meteor.

Use the mass spectrum to determine

the relative atomic mass of the

chromium in the rock. Based on your

result, make a recommendation as to

the origin of the rock sample.

(3 marks)

2. The data below gives the m/z ratio and relative abundance of different isotopes of an element X.

Determine the relative atomic mass of the element X to 1 d.p.

Suggest an identity for X. (4 marks)

3. The element magnesium (relative atomic mass 24.3) has three naturally occurring isotopes, 24.0

Mg, 25.0

Mg and 26.0

Mg. If the percentage of the heaviest isotope is 11.0%, what is the percentage of the

lightest isotope present? (3 marks)

m/z 204 205 206 207 208

Abundance 2.7 0.0 46.0 42.2 100.0

4.3% 8.2%

2.4%

85.1%


2.2. Isoelectronic species

For each of the species below, write out the full electronic configuration and then identify an anion and a

cation which is isoelectronic with the initial species.

e.g. neon, Ne; 1s2 2s2 2p6

Isoelectronic anion; F ¯ Isoelectronic cation; Mg2+

(1 mark for each correct electronic configuration, 1 mark for each correct isoelectronic anion and cation)

1. helium, He; ...............................................................................................................................

Isoelectronic anion; ……………..........................

Isoelectronic cation; ……….................................

2. krypton, Kr; ...............................................................................................................................



3. calcium ion, Ca2+

; .....................................................................................................................



BONUS 10th mark Identify a pair of common transition metal ions that are isoelectronic;

.................................................................................................................................................

Section 2

Formation of Ions


Spider Diagram

Trends in Ionisation Energy

Oxidation Numbers

Spider Diagram: Section 2 – Formation of Ions


2.4. Trends in ionisation energy

An atom’s ionisation energy is defined as;

‘The amount of energy required to remove one mole of electrons from one mole of atoms in the gaseous state’

The first ionisation energy is the energy required to remove the first electron [X(g) → X+(g) + 1 e−].

2.

(c) Explain any anomalies from the general trend using your understanding of electronic structure.

(2 marks)

(d) Which anomaly provides evidence for Hund’s rule?

(1 mark)

1st i

onis

atio

n e

nerg

y /

kJ m

ol−

1

Be Mg Ba Sr Ca

1st i

onis

atio

n e

nerg

y /

kJ m

ol−

1

Na Mg Al Si P S Cl Ar

1. (a) Sketch a plot of the first ionisation

energies of the elements of group 2

(1 mark)

(b) Explain the general trend shown using

your understanding of atomic

structure and electron configurations

(2 marks)

(a) Sketch a plot of the first ionisation

energies of the elements across

period 3

(2 marks)

(b) Explain the general trend shown

using your understanding of

electronic structure

(2 marks)

Redox 9.1.1

9.1.1 Oxidation numbers Work out the oxidation numbers for the bold elements in the compounds and perform the calculation.

Present your answer to your teacher. There is 1 mark for each correct oxidation number you have deduced.

Note: You are calculating the oxidation number of the element, not it’s overall contribution to the compound

eg, CaCl2 you would give the answer for Cl as -1, not as …Cl

2 (-2).

=

V3+

Cr2O72-

NaI

NaNO3

H2SO4

HClO

NaH

KMnO4

VO2+

MgO

Section 3

Intermolecular Bonding


Spider Diagram

Intermolecular Forces

Electronegativity and polarity

Spider Diagram: Section 3 – Intermolecular Bonding

Bonding 3.2.3.

3.2.3. Intermolecular forces

Molecules are attracted to each other by weak intermolecular forces. There are three types of intermolecular

force;

• Van der Waal’s forces

• Dipole-dipole forces

• Hydrogen bonding

For each group of molecules below, identify the strongest type of intermolecular force present in each

molecule (1 mark) and then use this information to order the molecules according to their boiling point, from

lowest to highest (1 mark).

1. CH4 SiH

4 SnH

4

2. NH3 PH

3 AsH

3

3. HF HCl HBr

4. CH3F CH

3Cl CH

4

5. HF H2O NH

3

Bonding 3.2.2.

3.2.2. Electronegativity and polarity

A polar bond is a bond in which the electrons between the atoms that are bonded together covalently are

shared unequally. The unequal share of electrons is usually shown by a + and a – sign. If a molecule

contains more than one polar bond, the effect of the polarity of all the bonds in the molecule may result in the

molecule having a dipole moment.

Use the table of the Pauling electronegativity of different

elements to identify any polar bonds in the molecules below.

Then use these polar bonds to decide if the molecule has a

dipole moment (this can be shown by an arrow with a line

through it ; the head of the arrow points towards the

negative end.)

(2 marks per molecule)

e.g. H2O

1. HCl

2. CO2

3. CH3Cl

4. CCl4

5. NH3

H

2.1

He

Li

1.0

Be

1.5

B

2.0

C

2.5

N

3.0

O

3.5

F

4.0

Ne

Na

0.9

Mg

1.2

Al

1.5

Si

1.8

P

2.1

S

2.5

Cl

3.0

Ar

Section 4

Bonding and Properties


Spider Diagram

Ionic Dot and Cross

Covalent Dot and Cross

Which type of Chemical Bond

Bonding Summary

Period 3 melting point

Period 3 Ionisation energy

Atomic Radius across Period 3

Properties and Bonding

Spider Diagram: Section 4 – Bonding and Properties

Bonding 3.1.2.

3.1.2. Ionic dot and cross

Draw dot and cross diagrams to illustrate the bonding in the following ionic compounds.

(2 marks for each correct diagram)

1. Lithium fluoride, LiF

2. Magnesium chloride, MgCl2

3. Magnesium oxide, MgO

4. Lithium hydroxide, LiOH

5. Sodium cyanide, NaCN

Bonding 3.1.1.

3.1.1. Covalent dot and cross

Draw dot and cross diagrams to illustrate the bonding in the following covalent compounds. If you wish you

need only draw the outer shell electrons;

(2 marks for each correct diagram)

1. Water, H2O

2. Carbon dioxide, CO2

3. Ethyne, C2H

2

4. Phosphoryl chloride, POCl3

5. Sulfuric acid, H2SO

4

Bonding 3.1.3.

3.1.3. Which type of chemical bond

There are three types of strong chemical bonds; ionic, covalent and metallic.

1. Sort the compounds below into groups within the circles below according to their chemical bonding;

sodium chloride, NaCl magnesium, Mg magnesium oxide, MgO

methane, CH4 oxygen, O2 barium iodide, BaI2

aluminium, Al ammonia, NH3 caesium, Cs

2. For each of the types of compound, indicate if you would expect them to;

(a) have a high or a low melting point

(b) conduct electricity

IONIC COVALENT

METALLIC

Bonding 3.1.4.

3.1.4. Bonding summary

A student has written the revision cards below to help her prepare for the exam. However she has made a

number of mistakes. Can you correct her cards to make sure she has accurate information to revise from;

(1 mark for each correct correction made)

Ionic bonding

Between a metal and a non-metallic atom, e.g. NaCl

Electrons are shared between the atoms

The molecules have high melting points owing to the strong electrostatic attraction between the ions

Ionic compounds do not conduct electricity at all as the ions that carry the current are held in a fixed position in the lattice structure

Metallic bonding

In metallic bonding, the outer electrons from the metal atoms merge to produce a lattice of negative metal ions in a sea of delocalised electrons

The strength of the metal depends on two things; - the charge on the metal ion - the size of the metal ion Therefore sodium is stronger than magnesium

Metals have low melting points because of the repulsive forces between the negative electrons which need little energy to be broken

Metals conduct electricity because of the sea of delocalised electrons which can move through the structure to carry the charge

Covalent bonding

Between two non-metallic atoms, e.g. CO2

Electrons are transferred between the atoms

Covalent molecules have high melting points because of the strong covalent bonds which must be broken

Covalent compounds do not conduct electricity at all as there are no free electrons

Periodic Trends 4.1.1

4.1.1 Period 3 melting points

1. Fill in the table above to show how melting point changes across Period 3 according to bonding type.

(6 marks)

2. Explain the differences in melting point between the following pairs of elements

(a) Magnesium and aluminium (2 marks)

(b) Phosphorus and sulfur (2 marks)

Element Na Mg Al Si P S Cl Ar

Bonding Metallic

Na+

Covalent molecular

Cl2

Intermolecular bonding

- - - -

Melting point (K)

1156 1380 2740 2628 553 718 238 87


4.1.2 Period 3 ionisation energy

1. The diagram shows the trend in 1

st ionisation energy across the Period 3 elements. Complete the

diagram giving the explanations for the trends seen. (7 marks)

2. Define the term 1st ionisation enthalpy and illustrate with an equation. (3 marks)

General trend1st ionisation energy increases acrossthe Period because...

Electron more tightly heldbecause

Easier to remove because

Slightly easier to removebecause


4.1.3 Atomic Radius across Period 3

1. State and explain the general trend in atomic radius across Period 3 (excluding Argon). (4 marks)

2. Atomic radius is a general term. Measurements are taken of metallic radii for metals and

covalent radii for molecules. Draw a diagram to show how you could calculate the atomic

radius of 2 covalently bonded atoms. (3 marks)

3. Why does Argon not follow the trend? (2 marks)

4. State the effect of atomic radius on the first ionisation energy of an element. (1 mark)

Bonding 3.3.

3.3. Properties and bonding

Match the compound on the left to its correct structure from the middle bank of statements and one or more

statements from the column on the right. Aluminium has been done for you

Compound Structure

aluminium

diamond

iodine

chlorine

potassium fluoride

molecular crystal

macromolecular crystal

simple covalent

molecule

ionic compound

Conducts electricity

when solid


when molten

High melting point

Low melting point

A covalent bond must

be broken to melt it

Weak intermolecular

forces are broken

when it boils

An ionic bond is broken

when it melts


when in solution

BONUS MARK Sketch the arrangement of molecules in a crystal of iodine

metal

Section 5

Chemical Equations


Spider Diagram

Balancing Equations

Writing Equations for Text

Spider Diagram: Section 5 – Chemical Equations

0.1.1. Balancing equations

© Royal Society of Chemistry, registered charity number 207890. This resource is shared under

a Creative Commons Attribution-NonCommercial-NoDerivatives 4.0 International licence. To view

a copy of the licence, visit https://creativecommons.org. Images © Shutterstock.

Balance the equations below.

1. …..C + …..O2 …..CO

2. …..Ba + …..H2O …..Ba(OH)2 + …..H2

3. …..C2H6 + …..O2 …..CO2 + …..H2O

4. …..HCl + …..Mg(OH)2 …..MgCl2 + H2O

5. …..N2 + …..O2 …..NO

6. …..Fe2O3 + …..C …..Fe + …..CO2

7. …..CH3CH2OH + …..[O] …..CH3COOH + …..H2O

8. …..HNO3 + …..CuO …..Cu(NO3)2 + H2O

9. …..Al3+ + …..e– …..Al

10. …..[Fe(H2O)6]3+

+ …..CO3

2– …..Fe(OH)3(H2O)3 + …..CO2 + …..H2O

(10 marks)

https://creativecommons.org/

0.1.3. Writing equations from text

© Royal Society of Chemistry, registered charity number 207890. This resource is shared under

a Creative Commons Attribution-NonCommercial-NoDerivatives 4.0 International licence. To view

a copy of the licence, visit https://creativecommons.org. Images © Shutterstock.

The following questions contain a written description of a reaction. In some cases the products

may be missing as you will be expected to predict the product using your prior knowledge.

For more advanced equations you may be given some of the formulae you need.

For each one, write a balanced symbol equation for the process. (10 marks)

1. The reaction between silicon and nitrogen to form silicon nitride Si3N4.

...............................................................................................................................................

2. The neutralisation of sulfuric acid with sodium hydroxide.

...............................................................................................................................................

3. The preparation of boron trichloride from its elements.

...............................................................................................................................................

4. The reaction of nitrogen and oxygen to form nitrogen monoxide.

...............................................................................................................................................

5. The combustion of ethanol (C2H5OH) to form carbon dioxide and water only.

...............................................................................................................................................

6. The formation of silicon tetrachloride (SiCl4) from SiO2 using chlorine gas and carbon.

...............................................................................................................................................

7. The extraction of iron from iron(III) oxide (Fe2O3) using carbon monoxide.

...............................................................................................................................................

8. The complete combustion of methane.

...............................................................................................................................................

9. The formation of one molecule of ClF3 from chlorine and fluorine molecules.

...............................................................................................................................................

10. The reaction of nitrogen dioxide with water and oxygen to form nitric acid.

...............................................................................................................................................

https://creativecommons.org/

Section 6

Inorganic Chemistry


Spider Diagram

Group 2 Trends

Group 7 Trends

Spider Diagram: Section 6 – Inorganic Chemistry

Periodic Trends 4.2

4.2 Group 2 trends

1. Complete the diagram showing the general trends in Group 2 by choosing properties from the grey

boxes to annotate the arrows. (7 marks)

2. The trend in solubility of the sulfates is useful as it provides a test for the sulfate anion. Describe how

this test is carried out, what is observed when the test is positive for sulfate ions and write an equation

including state symbols for this test. (3 marks)

Periodic Trends 4.3

4.3 Group 7 trends

F2

Cl2

Br2

I2

Increasedown

the group

Decreasedown

the group

Atomic radiusBoiling points

Reactivity(oxidising ability)

of halogen

Electronegativity

Reactivity(reducing ability)

of halide ion

1. Complete the diagram showing the general trends in Group 7 by choosing properties from the grey

boxes to annotate the arrows. (5 marks)

2. The trend in the reducing ability of the halide ions can be illustrated by the reaction of sodium halides

with sulfuric acid. Illustrate this trend using the equations for NaF and NaI and the O.S numbers for

sulfur. Write a conclusion stating what your equations show. (5 marks)

Section 7

Organic Chemistry


Spider Diagram

Functional Groups

Nomenclature

Formulae

Alkanes

Polymers from Alkenes

Alcohols

Spider Diagram: Section 7 – Organic Chemistry

Organic Chemistry 5.1.1

5.1.1 Functional groups

In each of the speech bubbles write the general name for the functional group ringed. (In the structures

below the rings are shown in a ’skeletal’ form. Where you can see a corner then there is a carbon with the

appropriate number of hydrogens)

8. Considering the limonene molecule given above.

(a) Draw the displayed formula (1 mark)

(b) Calculate the molecular formula (1 mark)

(c) Deduce the empirical formula (1 mark)

1.

2.

3.

4. 5.

Ether

6. 7.

Ester


5.1.2 Nomenclature

1. Group the following functional groups into prefixes (where the functional group goes before the naming

stem) and suffixes (goes after the naming stem). (4 marks)

2. Name the following compounds (6 marks)


5.1.3 Formulae

1. Define the term empirical formula (1)

2. An algebraic formula that can describe any member of a family of compounds is a way of describing a

type of formula in organic chemistry.

(a) Which formula does this definition refer to? (1 mark)

(b) What would this formula be for the family of –

i. Alkanes (1 mark)

ii Alkenes (1 mark)

3. The empirical formula of the compound 3-methylpentan-2,2-diol is C3H7O

(a) Deduce its molecular formula (1 mark)

(b) Write the structural formula for this compound (1 mark)

(c) Show the displayed formula of this compound (1 mark)

4. A hydrocarbon is shown to contain 92.3% carbon and 7.7% hydrogen by mass

(a) Calculate the empirical formula of the hydrocarbon (1 mark)

(b) The relative molecular mass of the hydrocarbon is 78. What is its molecular formula? (1 mark)

5. Cyclohexane and hex-1-ene have the same molecular formula. What is it? (1 mark)


5.2.4 Alkanes

1. Give 3 features of a homologous series (3 marks)

2. Name the method used to separate liquid hydrocarbons (1 mark)

3. The boiling point of alkanes increases as chain length increases. Explain this trend. (2 marks)

4. n-pentane (CH3CH

2CH

2CH

2CH

3) is the straight chain isomer of pentane.

(a) Draw the other 2 isomers (1 mark)

(b) State the difference in boiling point between these isomers and n-pentane. (1 mark)

(c) Explain this difference (2 marks)


5.2.11 Polymers from alkenes

1. Complete the following table to show the

alkene and the polymer it makes.

(4 marks)

2. In recent years the sustainability of plastics derived form crude oil has become a focus for chemists.

There are now many polymers derived from plant sources. Using your knowledge of reactions in AS

chemistry, suggest a route to poly(ethane) that is derived from a plant source. Include reagents and

conditions where appropriate. (6 marks)


5.3.3 Alcohols

Consider the following alcohols

Give the letters of the alcohols which…..

1. Are tertiary alcohols

2. Are oxidised to carboxylic acids

3. Show a colour change orange to green when treated with acidified potassium dichromate solution

Section 8

Chemical Reactions


Spider Diagram

Spider Diagram: Section 8 – Chemical Reactions

Section 9

Rates of Reactions


Spider Diagram

Collision Theory

The Importance of Maxwell-Boltzmann

Spider Diagram: Section 9 – Rates of Reactions

Kinetics 7.1.

7.1. Collision theory

This question is all about the reaction between zinc metal and hydrochloric acid to produce zinc chloride and

hydrogen gas.

1. Write a balanced symbol equation for the reaction that occurs. (1 mark)

......................................................................................................................................................................

2. The reaction flasks below show the same reaction but under different conditions. The acid is in

excess in all five flasks.

(a) In which flask is the reaction rate the slowest? ........................................................ (1 mark)

(b) The graph below shows how the volume of hydrogen given off changed with time for the reaction

that occurred in flask A. Sketch on the same set of axes, the curves you would expect to get if you

repeated the measurements for flasks B, C, D and E.

(8 marks)

1 g granular zinc,

100 cm3 1 mol dm

-3

HCl

1 g granular zinc

100 cm3 2 mol dm

-3

HCl at 35 °C

0.5 g powdered zinc

100 cm3 1 mol dm

-3

HCl at 55 °C

0.5 g granular zinc

100 cm3 0.5 mol dm

-3

HCl at 35°C

2 g powdered zinc

100 cm3 1 mol dm

-3

HCl at 35 °C

A

D C

B

E

Volume of

H2 / cm3

Flask A

Time / min

Kinetics 7.2.2.

7.2.2. The importance of Maxwell-Boltzmann

The distribution of energy amongst the particles in a gas is represented by the Maxwell Boltzmann

distribution. The key characteristics are;

1. No particles have zero energy

2. Most particles have intermediate energies

3. A few particles have very high energies indeed

4. The average energy is not the same as the most probable energy

1. The sketch opposite shows a typical

Maxwell Boltzmann distribution.

Indicate where on the curve each of the

above characteristics 1 - 4 is shown.

(4 marks)

2. Catalytic converters in cars reduce pollution by removing toxic gases

from exhaust fumes. The gases pass over a ceramic honeycomb coated

with platinum and rhodium metals. As the car warms up, the ceramic

honeycomb reaches its operating temperature and catalyses the reaction

between the gases.

(a) The distribution of energies of particles of gas at the start of the car

journey is shown on the graph below. Draw a second line on the

graph to indicate how the distribution will have changed 30 minutes

into the journey when the catalyst has reached its operating

temperature. (4 marks)

Energy, E

Energy, E

No. of particles with energy

Ea

(b) Why is it very important that

the catalytic converter reaches

its operating temperature as

quickly as possible?

(2 marks)

No. of particles with energy E

Section 10

Equilibria


Spider Diagram

Le Chatelier’s Principle

Equilibria and Industry

Spider Diagram: Section 10 – Equilibria

Equilibria 8.2.

8.2. Le Châtelier’s principle

Le Châtelier’s principle states that if a system at equilibrium is disturbed, the equilibrium moves in the

direction that tends to minimise the disturbance.

Use Châtelier’s principle to suggest two disturbances that can be made to each of the equilibria below to

bring about the desired changes; (2 marks for each question)

1. Cl2(aq) + H2O(l) ⇌ HClO(aq) + HCl(aq)

Two disturbances which would result in a decrease in the concentration of chlorine are;

......................................................................................................................................................................

......................................................................................................................................................................

2. 4 HCl + O2 ⇌ 2 Cl2 + 2 H2O ∆H –ive

Two disturbances which would result in an increase in the concentration of chlorine are;

......................................................................................................................................................................

......................................................................................................................................................................

3. PCl5(g) ⇌ PCl3(g) + Cl2(g) ∆H +ive

Two disturbances which could be made without changing the amount of reagents or products in the

system which would result in a shift of the equilibrium to the right are;

......................................................................................................................................................................

......................................................................................................................................................................

4. CH2=CH2(g) + H2O(g) ⇌ CH3CH2OH(g) ∆H –46 kJ mol–1

Two disturbances which would result in an increase in the percentage yield of ethanol are;

......................................................................................................................................................................

......................................................................................................................................................................

5. HCOOH + CH3OH ⇌ HCOOCH3 + H2O ∆H 0 kJ mol–1

Two disturbances which would result in no change in the position of the equilibrium are;

......................................................................................................................................................................

......................................................................................................................................................................

Equilibria 8.3.

8.3. Equilibria and industry

A number of industrial processes involve reversible reactions. In these cases, Le Châtelier’s principle can be

used to help find the best conditions for obtaining the maximum reaction yield.

1. Decide which set of conditions A – C would result in the highest yield of the desired product for each of

the equilibria (a) – (c) below; (2 marks)

(a) Production of hydrogen iodide A: low temperature H2(g) + I2(g) ⇌ 2 HI(g) ∆H +53 kJ mol

–1 high pressure

(b) Making hydrogen B: high temperature

CH4(g) + H2O(g) ⇌ 3 H2(g) + CO(g) ∆H +206 kJ mol–1

low pressure

(c) Production of methanol C: high temperature

CO(g) + 2 H2(g) ⇌ CH3OH(g) ∆H –91 kJ mol–1

pressure has no effect

2. Another industrial process involving a reversible reaction is the production of sulphuric acid in the

Contact Process. The first stage of the process is shown below;

2 SO2(g) + O2(g) ⇌ 2 SO3(g) ∆H –196 kJ mol–1

(a) i. Use Le Châtelier’s principle to explain why, at a given pressure, the percentage yield of

sulfur trioxide increases with a lowering of the overall temperature. (3 marks)

.......................................................................................................................................................

.......................................................................................................................................................

.......................................................................................................................................................

ii. To increase the rate of the reaction, a vanadium pentoxide catalyst is used. Explain what

effect this has on the overall percentage yield of sulfur trioxide. (2 marks)

.......................................................................................................................................................

.......................................................................................................................................................

(b) The reaction is run at pressures close to atmospheric pressure. Use Le Châtelier’s principle to

explain why this choice of pressure is unexpected and give a possible explanation for why it is

chosen. (3 marks)

...............................................................................................................................................................

...............................................................................................................................................................

...............................................................................................................................................................

Section 11

Calculations


Spider Diagram

Moles and Mass

Moles Summary

Empirical and Molecular Formulae

Percentage Yield and Atom Economy

Spider Diagram: Section 11 – Calculations

Quantitative Chemistry 1.1.1.

1.1.1. Moles and mass

Work out the answers to the following simple calculations (1 t = 1 tonne = 1,000 kg);

1. No. of moles in 10.0 g of O2 the mass in g of 2.41 moles of H

2O

(2 marks)

2. Mass in g of 0.2 moles of K2CO

3 mass in g of 0.5 moles of MgCO

3

(2 marks)

3. No. of moles in 12.4 t of NaNO3 ÷ no. of moles in 12.4 t of NaCl

(2 marks)

4. No. of moles in 25.9 g of sodium – no. of moles in 25.9 g of sodium chloride

(2 marks)

5. ? × molar mass of in g mol1 of calcium carbonate no. of moles in 4.2 kg of SiCl

4

(2 marks)

Quantitative Chemistry 1.1.4.

1.1.4. Moles summary

Mark the student’s answers to the questions below (shown to the right). Mark all 10 correctly to get the full

10 marks.

1. Magnesium reacts with acid as shown; Mg + 2 HCl → MgCl2 + H2

(a) How many moles of Mg reacts with 1 mole of HCl 1 mole

(b) How many moles of Mg must be reacted to produce 1 mole of H2 1 mole

2. Potassium reacts with water to produce potassium hydroxide and hydrogen gas.

(a) Write a balanced equation for the reaction K + 2 H2O → K(OH)2 + H2

(b) How many moles of potassium must be reacted with an excess of water to

produce 0.075 moles of potassium hydroxide? 0.075 moles

3. The dehydration of hydrated copper sulphate is a reversible reaction;

CuSO4.5H

20 ⇌ CuSO

4 + 5 H

2O

(a) What mass water is produced when 0.25 moles of hydrated copper sulphate is heated? 22.5 g

(b) What mass of hydrated copper sulphate must be heated to produce 18 g of H2O? 249.6 g

4. The equation for the complete combustion of methane is; CH4 + 2 O

2 → CO

2 + 2 H

2O

(a) How many moles of carbon dioxide would be produced by the complete combustion

of 8 g of CH4? 0.5 moles

(b) What mass of oxygen is needed for the complete combustion of 32 g of methane? 64 g

5. In an acid / base titration between ethanoic acid and sodium hydroxide the equation for the reaction is;

CH3COOH + NaOH → CH

3COONa

+ + H

2O

(a) How many moles of NaOH is needed to neutralise 50 cm3 of 0.1 mol dm

-3 CH

3COOH?

5 x 10-3 moles

(b) What volume of 0.1 mol dm-3

ethanoic acid is needed to neutralise 75 cm3 of 0.125 mol dm

-3 NaOH?

93.8 cm3

Quantitative Chemistry 1.4.

1.4. Empirical and molecular formulae

The technicians at the University have discovered a number of bottles containing amino acids which have

lost their labels. In order to identify them, they carried out elemental analyses. Use the information provided

to match the compound to its label;

(1 mark for each correct empirical formula, 1 mark for each correct match)

Alanine

C3H

7NO

2

Aspartic acid

C4H

8N

2O

3

Lysine

C6H

14N

2O

2

Threonine

C4H

9NO

3

Glutamine

C5H

10N

2O

3

Amino Acid E

C 40.4 %; H 7.9 %; N 15.7 %;

O 36.0 %

Amino Acid C

C 1.6 g; H 0.27 g; N 0.93 g;

O 1.6 g

Amino Acid B

C 36 g; H 7 g; N 14 g; O 16 g

Amino Acid D

C 40.3%; H 7.6%; N 11.8%;

O 40.3%

Amino Acid A

C 0.60 g; H 0.10 g; N 0.28 g;

O 0.48 g

Quantitative Chemistry 1.5.

1.5. Percentage yield and Atom economy

Percentage yield and atom economy are two numbers which help us gauge how efficient a reaction is for

making a specific chemical. The atom economy tells us in theory how many atoms must be wasted in a

reaction. The percentage yield tells us about the efficiency of the process.

1. Oxygen can be produced by a number of processes. Two possible processes are shown below;

Electrolysis of water; 2 H2O → 2 H

2 + O

2

Catalytic decomposition of hydrogen peroxide; 2 H2O

2 → 2 H

2O + O

2

By calculating the percentage atom economy of each process, decide which process is better for

producing oxygen. (3 marks)

2. Two students complete the synthesis of paracetamol from 4-aminophenol as shown by the equation

below;

4-aminophenol + ethanoyl chloride → paracetamol + hydrogen chloride

HOC6H

4NH

2 + CH

3COCl → HOC

6H

4NHCOCH

3 + HCl

Both students react 2 moles of 4-aminophenol with excess ethanoyl chloride.

Student 1 makes 1.5 moles of paracetamol.

Student 2 makes 220 g of paracetamol.

Which student has the better percentage yield? (4 marks)

3. Copper can be made by either roasting copper sulphide or by the reduction of copper carbonate with

carbon. The equations for the two processes are shown below.

CuS + O2 → Cu + SO

2

0.24 moles 0.18 moles

2 CuCO3 + C → 2 Cu + 3 CO

2

0.56 moles 0.36 moles

By comparing the percentage atom economy and the percentage yields of the processes as shown,

evaluate which is the better method from an industrial viewpoint.

(3 marks)

Section 12

Enthalpy


Spider Diagram

Definitions

Using bond enthalpies

Spider Diagram: Section 12 – Enthalpy

Thermodynamics 6.1.

6.1. Definitions

Complete the gaps in the boxes below;

Standard molar enthalpy change of combustion, ………..

Definition; The enthalpy change when one mole of a compound is completely burned in excess

oxygen under standard conditions, all reactants and products in their standard states.

e.g. Hc (C4H10); ………...…...............................................................................................................

(3 marks)

Standard enthalpy change, H

Definition; The heat energy change at ………...…................................................................................................ under

standard conditions (pressure ………...…...................... ; temperature ................................... ).

(2 marks)

Standard molar enthalpy change of formation, Hf

Definition; The enthalpy change when one mole………...…........................................................................................

………...…...........................................................................................................................................................................................................

………...…...........................................................................................................................................................................................................

e.g. Hf (NH3); 1/2 N2(g) + 3/2 H2(g) NH3(g) (3 marks)

Mean bond energy

Definition; The ………...….....................................................................................................................................................................

………...…...........................................................................................................................................................................................................

………...…...........................................................................................................................................................................................................

(2 marks)

Thermodynamics 6.4.

6.4. Using bond enthalpies

1. A student is carrying out a project to compare the theoretical

and experimental value for the enthalpy change of

combustion of ethanol. Using the data in the table, calculate

a theoretical value for Hc [CH3CH2OH(l)].

(HINT Remember to fully balance any equations before

starting your calculations) (4 marks)

2. When the student shows his calculation to his teacher, she points out that mean bond enthalpies are

only applicable for molecules in the gas state. Therefore the student must take into account the

enthalpy change of vaporisation of ethanol [CH3CH2OH(l) CH3CH2OH(g), Hvap +39 kJ mol1).

Use this value to correct your answer to Q1 (You may assume that the water formed from the

combustion is in the gas state). (1 mark)

3. The student now wishes to determine an experimental value for the enthalpy of

combustion of ethanol. He intends to burn approximately 1 g of fuel and measure the

heat energy produced by heating up a known volume of water in a copper calorimeter

(using the equipment shown).

Using your answer to question 2, suggest a suitable volume of water for the copper

calorimeter if he is aiming for a temperature rise of no more than 40 C?

(Specific heat capacity of water = 4.2 J K1 g1) (4 marks)

4. The experimental value obtained by the student is considerably lower than the theoretical value

calculated. Suggest one reason for this (other than experimental error). (1 mark)

Bond Mean bond

enthalpy / kJ mol1

C―C 347

C―H 413

C―O 358

O―H 464

O=O 498

C=O 805

Section 13

Investigating and Interpreting


Spider Diagram

Equipment

Treatment of Errors

Observation Exercise

Inferences

Spider Diagram: Section 13 – Investigating and Interpreting

Experimental skills 11.1.

11.1. Equipment

You have just started in a new research lab and have the following equipment available to you. Choose which piece(s) of equipment you would use for each of the following tasks and name each piece you use.

You can assume you have plenty of bungs, tubing and adaptors plus access to a safe source of heat;

1. To measure the volume of carbon dioxide produced by the following reaction;

ethanoic acid + sodium hydrogen carbonate → sodium ethanoate + water + carbon dioxide

2. To obtain pure water from a sample of salt water

3. To oxidise a sample of propan-1-ol to propanoic acid by heating with an excess of an acidified aqueous

solution of sodium dichromate under reflux conditions

4. To make up an accurate 0.05 mol dm–3

solution of sodium hydroxide

5. To recrystallise a sample of paracetamol contaminated with an insoluble solid impurity and collect the

solid product

(2 marks for each correct set of equipment correctly named)

Anti-bumping granules

1

2 3 4 5 6 7 8

10 11

12 13 14

Top pan balance

9

15


11.2. Treatment of errors

A student submitted the following report on his recent experiment to determine the concentration of a

solution of hydrogen peroxide;

1. The accuracy of each piece of equipment used is shown below.

(a) Using the individual equipment errors listed below, calculate the total percentage error for the

reaction. (4 marks)

100 cm3 measuring cylinder ± 0.5 cm

3 100 cm

3 gas syringe ± 0.5 cm

3

100 cm3 conical flask ± 5 cm

3 Top pan balance ± 0.005 g

(b) Suggest how this percentage error could have been reduced using the same equipment. (1 mark)

2. The teacher wants to give the student some feedback on how he could have improved the accuracy of

his experiment. Suggest two pieces of feedback she could give him. (2 marks)

3. Comment on the student’s use of significant figures in his analysis. From your comments, correct the

student’s final concentration of the hydrogen peroxide. (3 marks)

Method 10 cm3 of the hydrogen peroxide was measured using a 100 cm3 measuring cylinder and placed in a 100 cm3 conical flask. A bung and tubing was attached and the other end connected to a 100 cm3 gas syringe. 0.2 g of the MnO2 catalyst was weighed out. The catalyst was added to the hydrogen peroxide solution and the bung quickly replaced. The total volume of oxygen produced was measured.

Results

Experiment 1 2 3

Mass of MnO2 / g 0.21 0.24 0.20

Volume of gas produced / cm3 15 13 13

Analysis The equation for the reaction is: 2 H2O2 → 2 H2O + O2

Average volume of gas produced = 13.7 cm3

Assuming room temperature and pressure, no. of moles in 13.7 cm3 of gas = 13.7 cm3

÷ 24,000 cm3 mol–1 = 5.694 x 10-4 moles

2 moles of H2O2 produce 1 mole of O2 so the no. of moles of H2O2 in 10 cm3 = 1.139 x 10–3 moles

∴concentration of H2O2 solution = 1.139 x 10-3 mol ÷ 0.010 dm3 = 0.1139 mol dm–3


11.4. Observation exercises

A student has made the following observations for the reactions described. Unfortunately, the observations

are not described accurately enough for the student to get the marks. In each case write a more accurate

observation for the reaction that occurred.

Reaction Student observation Corrected / improved observation

A small piece of magnesium

was added to a test tube

containing hydrochloric acid;

Mg + 2 HCl → MgCl2 + H

2

A clear gas was produced

which burnt with a pop

A solution of

hex-1-ene was added dropwise

to bromine water

Br2 + CH

2=CHC

4H

9

CH2BrCHBrC

4H

9

The solution went clear

A solution of silver nitrate was

added dropwise to a solution

containing chloride ions until no

further change was observed.

Dilute ammonia was then added

dropwise.;

Ag+ (aq) + Cl¯ (aq)

AgCl (s)

[Ag(NH3)2]+ (aq)

The solution went cloudy

then clear

An excess of zinc powder was

added to a blue solution of

copper sulfate

Zn (s) + CuSO4 (aq)

ZnSO4 (aq) + Cu (s)

The solution turned clear

and an orange precipitate

formed

10 drops of an aldehyde was

added to a small quantity of

Tollens’ reagent in a test tube

and warmed

The solution turned silver

NH3 solution


11.5. Inferences

Use the student descriptions of some simple test tube reactions to identify each of the salts A to E;

(2 marks for the correct identification of the cation and anion present in each salt)

Salt A

• A lilac flame was produced with the flame test • A small quantity of the solid was dissolved in water and dilute nitric acid added followed by a few

drops of silver nitrate solution. A yellow precipitate formed which could not be dissolved by the addition of concentrated ammonia solution.

Salt B

A small sample of the salt was dissolved in water to make an aqueous solution. Dropwise addition of

NaOH(aq) to the aqueous solution produced a green precipitate which slowly turned brown on standing.

A separate sample of the aqueous solution was acidified by the addition of hydrochoric acid. On

addition of an aqueous solution of barium chloride a white precipitate formed.

Salt C

Addition of concentrated sulphuric acid to a small sample of the solid salt produced steamy fumes and a brown vapour. Carrying out the flame test on a small sample produced a red flame.

Salt D

Solution Test one1 Test two2

Aqueous solution of salt D

A white precipitate formed. A white precipitate formed which

redissolved on addition of dilute

ammonia solution to give a

colourless solution

1 Test one: Dropwise addition of a solution of sodium sulphate

2Test two: Addition of nitric acid followed by silver nitrate solution

Salt E

A small sample of the salt was dissolved in water to produce a colourless solution. The solution was split between two test tubes.

Addition of sodium hydroxide followed by aluminium powder to the first test tube produced ammonia gas.

Addition of an aqueous solution of potassium iodide to the second test tube yielded a bright yellow precipitate.

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