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VCE CHEMISTRY UNITS 3 AND 4

• Marilyn Schell •Margaret Hogan •

All rights reserved. No part of this publication may be reproduced, stored in a retrieval system, or transmitted in any form or by any means, electronic, mechanical, photocopying, recording or otherwise, without the prior permission of Science Press. ABN 98 000 073 861

© Science Press 2018First published 2018

Science PressPrivate Bag 7023 Marrickville NSW 1475 AustraliaTel: +61 2 9516 1122 Fax: +61 2 9550 [email protected] www.sciencepress.com.au

Contents

Words to Watch iv

Introduction v

Dot Points

Unit 3 How Can Chemical Processes Be Designed to Optimise Efficiency? vi

Unit 4 How Are Organic Compounds Categorised, Analysed and Used? vii

Questions

Unit 3 How Can Chemical Processes Be Designed to Optimise Efficiency? 1

Unit 4 How Are Organic Compounds Categorised, Analysed and Used? 107

Answers

Unit 3 How Can Chemical Processes Be Designed to Optimise Efficiency?

Unit 4 How Are Organic Compounds Categorised, Analysed and Used?

Appendix

Answers

Data Sheet

Periodic Table

Index

Dot Point VCE Chemistry Units 3 and 4 iii

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Contents

account, account for State reasons for, report on,

give an account of, narrate a series of events or

transactions.

analyse Interpret data to reach conclusions.

annotate Add brief notes to a diagram or graph.

apply Put to use in a particular situation.

assess Make a judgement about the value of

something.

calculate Find a numerical answer.

clarify Make clear or plain.

classify Arrange into classes, groups or categories.

comment Give a judgement based on a given

statement or result of a calculation.

compare Estimate, measure or note how things are

similar or different.

construct Represent or develop in graphical form.

contrast Show how things are different or opposite.

create Originate or bring into existence.

deduce Reach a conclusion from given information.

define Give the precise meaning of a word, phrase or

physical quantity.

demonstrate Show by example.

derive Manipulate a mathematical relationship(s) to

give a new equation or relationship.

describe Give a detailed account.

design Produce a plan, simulation or model.

determine Find the only possible answer.

discuss Talk or write about a topic, taking into account

different issues or ideas.

distinguish Give differences between two or more

different items.

draw Represent by means of pencil lines.

estimate Find an approximate value for an unknown

quantity.

evaluate Assess the implications and limitations.

examine Inquire into.

explain Make something clear or easy to understand.

extract Choose relevant and/or appropriate details.

extrapolate Infer from what is known.

hypothesise Suggest an explanation for a group of facts or phenomena.

identify Recognise and name.

interpret Draw meaning from.

investigate Plan, inquire into and draw conclusions about.

justify Support an argument or conclusion.

label Add labels to a diagram.

list Give a sequence of names or other brief answers.

measure Find a value for a quantity.

outline Give a brief account or summary.

plan Use strategies to develop a series of steps or processes.

predict Give an expected result.

propose Put forward a plan or suggestion for consideration or action.

recall Present remembered ideas, facts or experiences.

relate Tell or report about happenings, events or circumstances.

represent Use words, images or symbols to convey meaning.

select Choose in preference to another or others.

sequence Arrange in order.

show Give the steps in a calculation or derivation.

sketch Make a quick, rough drawing of something.

solve Work out the answer to a problem.

state Give a specific name, value or other brief answer.

suggest Put forward an idea for consideration.

summarise Give a brief statement of the main points.

synthesise Combine various elements to make a whole.

Words to Watch

Dot Point VCE Chemistry Units 3 and 4iv

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Words to Watch

What the book includes

This book provides questions and answers for each dot point in the Victorian Certificate of Education Study Design for each core topic in the Year 12 Chemistry syllabus:


• Area of Study 1 What Are the Options for Energy Production?

• Area of Study 2 How Can the Yield of a Chemical Product Be Optimised?


• Area of Study 1 How Can the Diversity of Carbon Compounds Be Explained and Categorised?

• Area of Study 2 What Is the Chemistry of Food?

Format of the book

The book has been formatted in the following way:

1.1 Subtopic from syllabus.

1.1.1 Assessment statement from syllabus.

1.1.1.1 First question for this assessment statement.

1.1.1.2 Second question for this assessment statement.

The number of lines provided for each answer gives an indication of how many marks the question might be worth in an examination. As a rough rule, every two lines of answer might be worth 1 mark.

How to use the book

Completing all questions will provide you with a summary of all the work you need to know from the syllabus. You may have done work in addition to this with your teacher as extension work. Obviously this is not covered, but you may need to know this additional work for your school exams.

When working through the questions, write the answers you have to look up in a different colour to those you know without having to research the work. This will provide you with a quick reference for work needing further revision.

Introduction

Dot Point VCE Chemistry Units 3 and 4 v

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Introduction

Area of Study 1 What Are the Options for Energy Production?

1.1 Obtaining energy from fuels 5

1.1.1 The definition of a fuel, including 5 the distinction between fossil fuels and biofuels with reference to origin and renewability (ability of a resource to be replaced by natural processes within a relatively short period of time).

1.1.2 Combustion of fuels as exothermic 9 reactions with reference to the use of the joule as the SI unit of energy, energy transformations and their efficiencies and measurement of enthalpy change including symbol (ΔH) and common units (kJ mol−1, kJ g−1, MJ/tonne).

1.1.3 The writing of balanced 13 thermochemical equations, including states, for the complete and incomplete combustion of hydrocarbons, methanol and ethanol, using experimental data and data tables.

1.1.4 The definition of gas pressure 15 including units, the universal gas equation and standard laboratory conditions (SLC) at 25°C and 100 kPa.

1.1.5 Calculations related to the combustion 19 of fuels including use of mass-mass, mass-volume and volume-volume stoichiometry in calculations of enthalpy change (excluding solution stoichiometry) to determine heat energy released, reactant and product amounts and net volume of greenhouse gases at a given temperature and pressure (or net mass) released per MJ of energy obtained.

1.1.6 The use of specific heat capacity 21 of water to determine the approximate amount of heat energy released in the combustion of a fuel.

1.2 Fuel choices 26

1.2.1 The comparison of fossil fuels 26 (coal, crude oil, petroleum gas, coal seam gas) and biofuels (biogas, bioethanol, biodiesel) with reference to energy content, renewability and environmental impacts related to sourcing and combustion.

1.2.2 The comparison of the suitability 31 of petrodiesel and biodiesel as transport fuels with reference to sources, chemical structures, combustion products, flow along fuel lines (implications of hygroscopic properties and impact of outside temperature on viscosity) and the environmental impacts associated with their extraction and production.

1.3 Galvanic cells as a source of energy 33

1.3.1 Redox reactions with reference to 33 electron transfer, reduction and oxidation reactions, reducing and oxidising agents, and use of oxidation numbers to identify conjugate reducing and oxidising agents.

1.3.2 The writing of balanced half 38 equations for oxidation and reduction reactions and balanced ionic equations, including states, for overall redox reactions.

1.3.3 Galvanic cells as primary cells 41 and as portable or fixed chemical energy storage devices that can produce electricity (details of specific cells not required) including common design features (anode, cathode, electrolytes, salt bridge and separation of half-cells) and chemical processes (electron and ion flows, half equations and overall equations).


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1.3.4 The comparison of the energy 45 transformations occurring in spontaneous exothermic redox reactions involving direct contact between reactants (transformation of chemical energy to heat energy) compared with those occurring when the reactants are separated in galvanic cells (transformation of chemical energy to electrical energy).

1.3.5 The use of the electrochemical 46 series in designing and constructing galvanic cells and as a tool for predicting the products of redox reactions, deducing overall equations from redox half equations and determining maximum cell voltage under standard conditions.

1.4 Fuel cells as a source of energy 50

1.4.1 The common design features 50 of fuel cells including use of porous electrodes for gaseous reactants to increase cell efficiency (details of specific cells not required).

1.4.2 The comparison of the use of 52 fuel cells and combustion of fuels to supply energy with reference to their energy efficiencies (qualitative), safety, fuel supply (including the storage of hydrogen), production of greenhouse gases and applications.

1.4.3 The comparison of fuel cells and 54 galvanic cells with reference to their definitions, functions, design features, energy transformations, energy efficiencies (qualitative) and applications.

Area of Study 2 How Can the Yield of a Chemical Product Be Optimised?

2.1 Rate of chemical reactions 57

2.1.1 Chemical reactions with reference 57 to collision theory, including qualitative interpretation of Maxwell-Boltzmann distribution curves.

2.1.2 The comparison of exothermic 62 and endothermic reactions including their enthalpy changes and representations in energy profile diagrams.

2.1.3 Factors affecting the rate of a 66 chemical reaction including temperature, surface area concentration of solutions, gas pressures and presence of a catalyst.

2.1.4 The role of catalysts in changing 71 the rate of chemical reactions with reference to alternative reaction pathways and their representation in energy profile diagrams.

2.2 Extent of chemical reactions 73

2.2.1 The distinction between reversible 73 and irreversible reactions, and between rate and extent of a reaction.

2.2.2 Homogenous equilibria involving 75 aqueous solutions or gases with reference to collision theory and representation by balanced chemical or thermochemical equations (including states) and by concentration-time graphs.



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2.2.3 Calculations involving equilibrium 80 expressions and equilibrium constants (Kc only) for a closed homogeneous equilibrium system including dependence of value of equilibrium constant, and its units, on the equation used to represent the reaction and on the temperature.

2.2.4 Le Châtelier’s principle: 86 identification of factors that favour the yield of a chemical reaction, representation of equilibrium system changes using concentration-time graphs and applications, including competing equilibria involved in the occurrence and treatment of carbon monoxide poisoning resulting from incomplete combustion of fuels.

2.3 Production of chemicals by 91 electrolysis

2.3.1 Electrolysis of molten liquids 91 and aqueous solutions using different electrodes.

2.3.2 The general operating principles 94 of commercial electrolytic cells, including basic structural features and selection of suitable electrolyte (molten or aqueous) and electrode (inert or reactive) materials to obtain desired products (no specific cell is required).

2.3.3 The use of the electrochemical 98 series to explain or predict the products of an electrolysis, including identification of species that are preferentially discharged, balanced half equations, a balanced ionic equation for the overall cell reaction, and states.

2.3.4 The comparison of an electrolytic 101 cell with a galvanic cell with reference to the energy transformations involved and basic structural features and processes.

2.3.5 The application of stoichiometry 104 and Faraday’s laws to determine amounts of product, current or time for a particular electrolytic process.

2.4 Rechargeable batteries 106

2.4.1 The operation of rechargeable 106 batteries (secondary cells) with reference to discharging as a galvanic cell and recharging as an electrolytic cell, including the redox principles (redox reactions and polarity of electrodes) and the factors affecting battery life with reference to components and temperature (no specific battery is required).

Answers to How Can Chemical Processes Be Designed to Optimise Efficiency?



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Area of Study 1 How Can the Diversity of Carbon Compounds Be Explained and Categorised?

1.1 Structure and nomenclature of organic compounds

1.1.1 The carbon atom with reference to valence number, bond strength, stability of carbon bonds with other elements and the formation of isomers (structural and stereoisomers) to explain carbon compound diversity, including identification of chiral centres in optical isomers of simple organic compounds and distinction between cis and trans isomers in simple geometric isomers.

1.1.2 Structures including molecular, structural and semi-structural formulas of alkanes (including cyclohexane), alkenes, alkynes, benzene, haloalkanes, primary amines, primary amides, alcohols (primary, secondary, tertiary), aldehydes, ketones, carboxylic acids and non-branched esters.

1.1.3 IUPAC systematic naming of organic compounds up to C8 with no more than two functional groups for a molecule, limited to non-cyclic hydrocarbons, haloalkanes, primary amines, alcohols (primary, secondary, tertiary), carboxylic acids and non-branched esters.

1.2 Categories, properties and reactions of organic compounds

1.2.1 An explanation of trends in physical properties (boiling point, viscosity) and flash point with reference to structure and bonding.

1.2.2 Organic reactions, including appropriate equations and reagents, for the oxidation of primary and secondary alcohols, substitution reactions of haloalkanes, addition reactions of alkenes, hydrolysis reactions of esters, the condensation reaction between an amine and a carboxylic acid, and the esterification reaction between an alcohol and a carboxylic acid.

1.2.3 The pathways used to synthesise primary haloalkanes, primary alcohols, primary amines, carboxylic acids and esters, including calculations of atom economy and percentage yield of single-step or overall pathway reactions.

1.3 Analysis of organic compounds

1.3.1 The principles and applications of mass spectroscopy (excluding features of instrumentation and operation) and interpretation of qualitative and quantitative data, including identification of molecular ion peak, determination of molecular mass and identification of simple fragments.

1.3.2 The principles and applications of infra-red spectroscopy (IR) (excluding features of instrumentation and operation) and interpretation of qualitative and quantitative data including use of characteristic absorption bands to identify bonds.



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1.3.3 The principles (including spin energy levels) and applications of proton and carbon-13 nuclear magnetic resonance spectroscopy (NMR) (excluding features of instrumentation and operation); analysis of carbon-13 NMR spectra and use of chemical shifts to determine number and nature of different carbon environments in a simple organic compound; and analysis of high resolution proton NMR spectra to determine the structure of a simple organic compound using chemical shifts, areas under peak and peak splitting patterns (excluding coupling constants) and application of the n + 1 rule.

1.3.4 Determination of the structures of simple organic compounds using a combination of mass spectrometry (MS), infra-red spectroscopy (IR) and proton and carbon-13 nuclear magnetic resonance spectroscopy (NMR) (limited to data analysis).

1.3.5 The principles of chromatography including use of high performance liquid chromatography (HPLC) and construction and use of a calibration curve to determine the concentration of an organic compound in a solution.

1.3.6 Determination of the concentration of an organic compound by volumetric analysis, including the principles of direct acid-base and redox titrations (excluding back titrations).

Area of Study 2 What Is the Chemistry of Food?

2.1 Key food molecules

2.1.1 Proteins: formation of dipeptides and polypeptides as condensation polymers of 2-amino acids; primary (including peptide links), secondary, tertiary and quaternary structure and bonding; distinction between essential and non-essential amino acids as dietary components.

2.1.2 Carbohydrates: formation of disaccharides from monosaccharides, and of complex carbohydrates (specifically starch and cellulose) as condensation polymers of monosaccharides; glycosidic links; storage of excess glucose in the body as glycogen; comparison of glucose, fructose, sucrose and the artificial sweetener aspartame with reference to their structures and energy content.

2.1.3 Fats and oils (triglycerides): common structural features including ester links; distinction between fats and oils with reference to melting points; explanation of different melting points of triglycerides with reference to the structures of their fatty acid tails and the strength of intermolecular forces; chemical structures of saturated and unsaturated (monounsaturated and polyunsaturated) fatty acids; distinction between essential and nonessential fatty acids; and structural differences between omega-3 fatty acids and omega-6 fatty acids.



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2.1.4 Vitamins: inability of humans to synthesise most vitamins (except vitamin D) making them essential dietary requirements; comparison of structural features of vitamin C (illustrative of a water soluble vitamin) and vitamin D (illustrative of a fat soluble vitamin) that determine their solubility in water or oil.

2.2 Metabolism of food in the human body

2.2.1 Metabolism of food as a source of energy and raw materials: general principles of metabolism of food involving enzyme catalysed chemical reactions with reference to the breakdown of large biomolecules in food by hydrolytic reactions to produce smaller molecules, and the subsequent synthesis of large biologically important molecules by condensation reactions of smaller molecules.

2.2.2 Enzymes as protein catalysts: active site; modelling of process by which enzymes control specific biochemical reactions (lock and key and induced fit models); consequences of variation in enzyme-substrate interaction (lock and key mechanism) due to the behaviour of a particular optical isomer; explanation of effects of changes in pH (formation of zwitterions and denaturation), increased temperature (denaturation) and decreased temperature (reduction in activity) on enzyme activity with reference to structure and bonding; action of enzymes in narrow pH ranges; and use of reaction rates to measure enzyme activity.

2.2.3 The distinction between denaturation of a protein and hydrolysis of its primary structure.

2.2.4 Hydrolysis of starch in the body: explanation of the ability of all humans to hydrolyse starch but not cellulose, and of differential ability in humans to hydrolyse lactose; glycaemic index (GI) of foods as a ranking of carbohydrates based on the hydrolysis of starches (varying proportions of amylose and amylopectin) to produce glucose in the body.

2.2.5 Hydrolysis of fats and oils from foods to produce glycerol and fatty acids; oxidative rancidity with reference to chemical reactions and processes, and the role of antioxidants in slowing rate of oxidative rancidity.

2.2.6 The principles of the action of coenzymes (often derived from vitamins) as organic molecules that bind to the active site of an enzyme during catalysis, thereby changing the surface shape and hence the binding properties of the active site to enable function as intermediate carriers of electrons and/or groups of atoms (no specific cases required).

2.3 Energy content of food

2.3.1 The comparison of energy values of carbohydrates, proteins and fats and oils.

2.3.2 Glucose as the primary energy source, including a balanced thermochemical equation for cellular respiration.

2.3.3 The principles of calorimetry; solution and bomb calorimetry, including determination of calibration factor and consideration of the effects of heat loss; and analysis of temperature-time graphs obtained from solution calorimetry.



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DOT POINT

1 How Can Chemical Processes Be Designed to Optimise Efficiency?

Unit 3How Can Chemical Processes Be Designed to Optimise Efficiency?

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Dot Point VCE Chemistry Units 3 and 4

What Are the Options for Energy Production?

DOT POINT

AREA OF STUDY 1


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1.1 Obtaining energy from fuels.

1.1.1 The definition of a fuel, including the distinction between fossil fuels and biofuels with reference to origin and renewability (ability of a resource to be replaced by natural processes within a relatively short period of time).

1.1.1.1

(a) What is meant by the term fossil fuel?

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(b) Name three fossil fuels.

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(c) For one of these fossil fuels, explain the photosynthetic origin of its energy.

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1.1.1.2

(a) Explain why carbon dioxide is produced when fossil fuels are burnt to produce energy.

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(b) Define the term ‘non-renewable’.

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(c) Justify the use of the term non-renewable being applied to fossil fuels.

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1.1.1.3

(a) Define ‘biofuel’ and name two examples.

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(b) Outline three concerns that have led to the development of biofuels.

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1.1.1.4 Ethanol can be used as a renewable alternative to fossil fuels.

(a) Justify the classification of ethanol as a renewable fuel.

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(b) One argument against using ethanol as fuel is that the crops required, and the land needed to grow these crops, could be better used in providing food for starving people. Evaluate this argument.

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6How Can Chemical Processes Be Designed to Optimise Efficiency?

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1.1.1.5 Biofuels are claimed to be ‘carbon neutral’.

(a) What is meant by this term?

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(b) Is it correct to describe biofuels as carbon neutral?

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1.1.1.6 Outline the processes involved in the industrial production of ethanol from sugar cane.

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1.1.1.7 Justify the classification of ethanol as a renewable resource.

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1.1.1.8 Distinguish between the terms renew, reuse and recycle.

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1.1.1.9 Which alternative most accurately presents an advantage of using ethanol as an alternative fuel?

(A) Ethanol does not produce any greenhouse gases during combustion.

(B) Ethanol uses up greenhouse gases when it burns.

(C) During combustion, ethanol produces more greenhouse gases than other fuels.

(D) The net production of greenhouse gases is lower for ethanol than for petrol.

Explain your response.

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1.1.1.10 A fuel is a substance that can be used to release energy.

(a) What is the unit of energy?

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(b) Distinguish between fossil fuels and biofuels in terms of how they are formed.

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(c) Distinguish between fossil fuels and biofuels in terms of their renewability.

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(d) Name two fossil fuels and two biofuels.

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1.1.2 Combustion of fuels as exothermic reactions with reference to the use of the joule as the SI unit of energy, energy transformations and their efficiencies and measurement of enthalpy change including symbol (ΔH) and common units (kJ mol−1, kJ g−1, MJ/tonne).

1.1.2.1

(a) State the law of conservation of energy.

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(b) Outline an example of the change of energy from one form to another.

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(c) When we burn substances such as wood, coal or methane gas, heat energy is given out. If energy is not being created, identify the source of this heat energy.

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1.1.2.2 Distinguish between endothermic and exothermic reactions and identify two examples of each.

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1.1.2.3 Outline the meaning of enthalpy (H).

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1.1.2.4 Outline what is meant by ΔH and how it is measured.

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1.1.2.5 The following equation represents the combustion of octane, a major component of petrol. This is an exothermic chemical reaction.

2C8H18(l) + 25O2(g) → 16CO2(g) + 18H2O(l)

Predict whether the total bond energy of the reactant chemicals C8H18 and O2 would be greater, equal to or less than the total bond energy of the products CO2 and H2O.

Explain your answer.

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1.1.2.6 Energy profile diagrams can be used to show information about exothermic and endothermic reactions.

The following illustration shows an energy profile diagram for an exothermic reaction system.

Ene

rgy

Reaction path

H

EA

ER

EP

Where:

ER = energy of reactants

EP = energy of products

EA = activation energy (the energy needed to start the reaction)

DH = change in enthalpy

The enthalpy change (DH) is said to be negative because during the reaction the energy level within

the system has �� (decreased/increased).

1.1.2.7 Which alternative below provides a correct definition for the molar heat of combustion of a fuel?

(A) The heat energy, in joules or kilojoules, released by the combustion of 1 mole of a fuel.

(B) The heat energy, in joules or kilojoules, used by the combustion of 1 mole of a fuel.

(C) The heat energy, in kilojoules, released by the combustion of 1 kilogram of a fuel.

(D) The heat energy, in kilojoules, used to burn 1 kilogram of a fuel.

Define combustion.

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1.1.2.8 The table below compares the molar heat of combustion for four different alcohols.

Alcohol Molar mass (g) Heat of combustion (kJ mol–1)

A 60 2016

B 74 2677

C 32 727

D 46 1367

Which fuel would provide the most energy for each gram burned?

(A) A (B) B (C) C (D) D

Show your calculations for the correct answer.

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1.1.2.9

(a) Describe a method that could be used in a school laboratory to determine the experimental heat of combustion of a liquid fuel such as ethanol. Use procedural text type. Indicate the measurements that would be needed.

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(b) Suggest reasons why the value obtained is likely to be much less than the theoretical value.

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(c) Write balanced equations for the complete and incomplete combustion of octane. Indicate which equation is for complete combustion and which is for incomplete combustion.

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1.1.2.10 A group of year 12 students performed a first-hand investigation of the heat of combustion of ethanol. They burned ethanol in a spirit burner, and used it to heat 100 mL of water, as shown in the diagram below.

Container

Thermometer

Water

Wick

Fuel, e.g. ethanolin spirit burner

Lid

The results they obtained were:

Initial temperature of 100 mL water = 22.6°C

Final temperature of 100 mL water = 35.9°C

Initial mass of spirit burner + ethanol = 235.56 g

Final mass of spirit burner + ethanol = 234.23 g

Specific heat of water = 4.18 × 103 J kg–1 K–1

Use these results to calculate the experimental molar heat of combustion of ethanol.

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DOT POINT

Answers

221 Answers


1.1.1.1 (a) A fossil fuel is a substance which is burned to provide energy and which has been produced from once-living organisms being buried by sediment and compressed for millions of years.

(b) Coal, natural gas, petroleum.

(c) Various. Coal – plants are buried by sediments and compressed for millions of years. Originally, when the plants were alive they obtained energy from the Sun by using solar energy for the process of photosynthesis in (the chloroplasts of ) their cells.

Carbon dioxide + water solar energy→ glucose + oxygen

6CO2(g) + 6H2O(l) solar energy→ C6H12O6(aq) + 6O2(g)

The solar energy absorbed was converted into the chemical energy in glucose which became part of the structure of the plant. Compression for millions of years converted the buried plant material to peat, then brown coal and black coal.

1.1.1.2 (a) All fossil fuels are made from once-living things which are carbon based, and so fossil fuels contain carbon or its compounds. These will produce carbon dioxide when they undergo combustion.

Petroleum contains hydrocarbons which undergo combustion to produce carbon dioxide and water, e.g.

Octane + oxygen → carbon dioxide + water

C8 H18 (l) + 25 ___ 2

O2(g) → 8CO2(g) + 9H2O(l)

Natural gas contains mainly methane which also burns to produce carbon dioxide and water.

CH4(g) + 2O2(g) → CO2(g) + 2H2O(g)

Coal is mainly carbon which burns to produce carbon dioxide.

C(s) + O2(g) → CO2(g)

(b) Non-renewable means unable to produce more of a substance after existing supplies have been used up.

(c) This applies to all fossil fuels because they are made by the burial of once-living organisms and their compression for millions of years. We are using huge amounts of fossil fuels and nowhere are new supplies being generated. Even if they were, it would take too long (millions of years) to convert the organic matter into fuels. Thus we cannot replace the fuels we are using.

1.1.1.3 (a) A biofuel is a liquid fuel produced from living or recently living materials such as waste plant and animal matter, e.g. bioethanol and biodiesel.

(b) Various. Biofuels have been developed because of concerns about:

• The pollution caused by the combustion of fossil fuels, especially by transport vehicles which are one of the top emitters of greenhouse gases.

• Fears that supplies of fossil fuels are running out as they are a finite resource and non-renewable.

• The lack of availability of fuels due to political unrest.

• Instability in world markets – the cost of importing fuel and its effect on inflation. In Australia we import petrol.

1.1.1.4 (a) Ethanol can be made from crops such as sugar cane and corn or their waste products. As the ethanol is used up, more crops can be grown to make more ethanol. Its ability to be replaced when used means it is renewable.

(b) If food crops are used the argument has validity. However, today there is no need to use edible food; with better production techniques now available, ethanol can be produced from bagasse, the wastes left after the consumable part of the crop (e.g. corn or sugar) has been harvested. This also prevents the burning off, after harvesting, of waste crops, e.g. the stalks, a process which causes considerable pollution. Also, weather-damaged crops and forestry wastes that are unsuitable for human consumption can be used to make ethanol, although this is currently more expensive.

1.1.1.5 (a) Carbon neutral means that the fuel has no net output of carbon dioxide – the amount of carbon dioxide released by combustion of the fuel is the same as the amount of carbon dioxide taken out of the atmosphere for photosynthesis by the growing plant.

(b) Biofuels are not really carbon neutral, although they can be closer to carbon neutral than petrol. Carbon dioxide is taken out of the atmosphere during the growth of biomass crops. However, not only is carbon dioxide released to the atmosphere by combustion of the fuel, but more is released when producing fertiliser for the crops, and producing energy to make the fuel from biomass, and transport the chemicals. Thus more carbon dioxide is released to the atmosphere than is taken out.

1.1.1.6 Ethanol can be produced by fermenting sugar in soluble forms such as sucrose and molasses from sugar cane, and fructose from corn plants. If cellulose is present, it needs to be first hydrolysed to glucose.

The sugars are fermented by enzymes produced by fungi such as the yeast Saccharomyces cerevisiae.

About 30% of the sugar produced from plant matter is in the form of xylose, a sugar that cannot be fermented by fungi. To overcome this problem, genetically engineered E. coli bacteria are now being used instead of fungi as they can ferment both glucose and xylose.

After fermentation the ethanol must be separated from the reaction mixture by distillation.

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Dot Point VCE Chemistry Units 3 and 4222Answers

1.1.1.7 Ethanol is a renewable resource because it is mainly produced by the fermentation of plant matter, such as the residues from the production of corn and sugar cane. More crops can be grown to make more ethanol and replace that which is used.

1.1.1.8 Renew – able to grow more and replace that which was used.

Reuse and recycle – use a substance again, e.g. melt down aluminium cans and recast them to use again. Note that you cannot reuse or recycle a fuel.

1.1.1.9 D Combustion of ethanol does produce greenhouse gases, e.g. CO2, but some CO2 is also used up during its production, so the net CO2 volume produced is lowered.

1.1.1.10 (a) The joule (J).

(b) Fossil fuels are formed by geological processes such as heat and pressure, compressing and changing layers of the remains of once living plants and animals. Biofuels are produced using living, or recently living materials such as waste plant and animal matter. Biofuels are produced over a short time span but fossil fuels can only be produced over millions of years.

(c) Fossil fuels cannot be renewed, biofuels can be renewed.

(d) Various, e.g. Fossil fuels – coal, petroleum; biofuels – bioethanol, biodiesel.

1.1.2.1 (a) Energy can be changed from one form to another, but it cannot be created or destroyed. The energy of the Universe is conserved (stays constant).

(b) Various, e.g. As a ball rolls down a hill its potential energy decreases (lower height) and its kinetic energy increases (it gets faster). So energy has been changed from potential to kinetic energy.

(c) Chemical potential energy stored in the bonds of the substance being burnt is changing into heat energy.

1.1.2.2 Endothermic – a reaction in which heat is absorbed into the system.

Examples – various, e.g. dissolving sodium hydroxide or sodium thiosulfate in water; photosynthesis.

Exothermic – a reaction in which heat is released into the environment.

Examples – various, e.g. any combustion reaction; acid on active metals, acids on carbonates, neutralisation reactions, respiration. Give specific examples, e.g. the combustion of hydrogen and the neutralisation of sodium hydroxide by hydrochloric acid.

1.1.2.3 Enthalpy is the internal energy of a system. The symbol for enthalpy is H. Enthalpy cannot be measured, but the change in enthalpy (ΔH) can be measured.

1.1.2.4 ΔH is the change in enthalpy as bonds are broken and formed during a chemical reaction. This can be measured by the change in heat energy of the system – the increase or decrease in temperature as the reaction proceeds.

1.1.2.5 Greater. This is an exothermic reaction so energy is released to the environment. Energy is needed to break the C-H and O=O bonds within the octane and oxygen molecules. But more energy is released when carbon dioxide and water molecules form. The bonds in the products contain less potential chemical energy than the bonds in the reactants. The energy difference is released to the environment.

1.1.2.6 Decreased.

1.1.2.7 A Combustion – a chemical reaction involving combination with oxygen and the release of energy to the surroundings.

1.1.2.8 B 2677 ______ 74

= 36 kJ/g. (A) = 34, (B) = 36, (C) = 23, (D) = 30, making (B) the highest.

1.1.2.9 (a) Put a small amount of the fuel in a spirit burner and weigh the spirit burner + fuel.

Measure 100 mL water into a conical flask (or similar container) and record the temperature.

Suspend the conical flask over the spirit burner so as much heat as possible will go into the water.

Light the wick of the spirit burner. Measure and record the temperature of the water frequently, stirring as it is heated until the temperature increases about 10°C.

Put out the flame. Measure and record the maximum temperature reached by the water.

Reweigh the spirit burner + fuel and calculate fuel used.

Calculate the heat of combustion.

(b) Heat is lost to the surrounding air and to the container.

Also combustion of the fuel may not be complete. The theoretical value is for the complete combustion of ethanol and less energy is obtained if the combustion is incomplete.

(c) Complete combustion: C8H18(l) + 25 ___ 2

O2(g) → 8CO2(g) + 9H2O(l)

Incomplete combustion: Various, e.g. C8H18(l) + 17 ___ 2

O2(g) → 2CO2(g) + 4CO(g) + 2C(s) + 9H2O(l)

1.1.2.10 193 kJ mol–1

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Dot Point VCE Chemistry Units 3 and 4 223 Answers

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