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AS Chemistry. Enthalpy Changes. Learning Objectives Candidates should be able to: Explain that some chemical reactions are accompanied by energy changes, principally in the form of heat energy.

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slide1

AS Chemistry

Enthalpy Changes

slide2

Learning Objectives

  • Candidates should be able to:
  • Explain that some chemical reactions are accompanied by energy changes, principally in the form of heat energy.
  • Construct (and interpret) a reaction pathway diagram, in terms of the enthalpy change of the reaction.
  • Calculate enthalpy changes from appropriate experimental results, including the use of the relationship: q=mcT.
  • Explain and use the terms enthalpy change of reaction and standard conditions, with particular reference to formation and combustion.
slide3

Starter activity

Can you complete tasks 1 and 2 on your notes.

slide10

AS Chemistry

Hess’ Law and Enthalpy Cycles

slide11

Learning Objectives

  • Candidates should be able to:
  • apply Hess’ Law to construct simple energy cycles
  • carry out calculations involving such cycles
slide12

Starter activity

Can you write out definitions for ΔHθc and ΔHθf ?

slide13

Standard enthalpy change of combustion, ΔHθc, 298 is the enthalpy change when 1 mole of a substance is burned completely in oxygen under standard conditions (100kPa and 298K), all reactants and products being in their standard states.

Standard enthalpy change of formation, ΔHθf, 298 is the enthalpy change when 1 mole of a compound is formed from its elements under standard conditions (100kPa and 298K), all reactants and products being in their standard states.

slide14

Hess’ Law

The enthalpy change of a reaction depends only on the initial and final state of the reaction and is independent of the route by which the reaction may occur.

slide15

‘CRAP rule’

H = (Hc reactants) - (Hc products)

H = (Hf products) - (Hf reactants)

slide19

AS Chemistry

Bond Enthalpies

slide20

Learning Objectives

Candidates should be able to apply Hess’ Law to construct simple energy cycles, and carry out calculations involving such cycles and relevant energy terms, with particular reference to average bond energies.

slide21

Starter activity

Can you write equations for the ΔHθc and ΔHθfof glucose (C6H12O6)?

slide22

Bond breaking and bond making

Chemical reactions involve bond breaking and bond making.

slide23

Bond energy

The quantity of energy needed to break a particular bond in a molecule is called the bond dissociation enthalpy (Hdiss), or bond enthalpy for short. It refers to the enthalpy change when one mole of bonds of the same type are broken in gaseous molecules under standard conditions.

slide24

Bond energy

H – H (g) H(g) + H(g)

slide25

Mean Bond Enthalpy

The mean bond enthalpy is the amount of energy needed to break a covalent bond.

They are average values taken from many different molecules

slide27

Bond breaking:

Total endothermic value = (+347 x 1) + (+413 x 5) + (+358 x 1) + (+464 x 1) + (+498 x 3) = +4728 kJ

Bond making:

Total exothermic value = (-464 x 6) + (-805 x 4) = -6004 kJ

Sum total of bond breaking and bond making:

DHc = +4728 + - 6004 = -1276 kJ mol-1

slide28

AS Chemistry

Kinetics

slide29

Learning Objectives

  • Candidates should be able to:
    • Explain and use the terms rate of reaction and activation energy.
    • Show understanding, including reference to the Boltzmann distribution, of what is meant by the term activation energy.
slide30

Starter activity

Working in groups of 3, complete task 1.

slide31

Different rates of reaction

Seconds

Minutes

Hours

Days

Weeks

Months

Years

Decades

Centuries

Millennia

dynamite exploding

magnesium and acid

cake baking

fruit ripening

plants growing

rusting of iron

erosion of rock

crude oil forming

slide32

Rate of Reaction

The rate of a reaction is found by measuring the amount in moles of a reactant which is used up, or the amount of product produced, in a given time.

The units are often mol dm-3 s-1.

changing the rate of a reaction
Changing the rate of a reaction

There are five factors which can affect the rate of a reaction:

  • Surface Area
  • Concentration
  • Temperature
  • Use of a catalyst
  • Intensity of light
slide34

Collision Theory

Reactions occur when the particles of reactants collide, provided they collide with a certain minimum amount of kinetic energy (and in the correct orientation).

slide36

A

B

Reactants

Enthalpy

Enthalpy

Enthalpy

Enthalpy

Products

Reactants

Products

Progress of Reaction

Progress of Reaction

Progress of Reaction

Progress of Reaction

C

D

Products

Reactants

Products

Reactants

slide39

AS Chemistry

Effect of Temperature

slide40

Learning Objectives

  • Candidates should be able to:
    • Explain qualitatively, in terms of both of the Boltzmann distribution and of collision frequency, the effect of temperature change on the rate of reaction.
slide41

Starter activity

Can you complete task 1?

slide47

AS Chemistry

Catalysis

slide48

Learning Objectives

  • Candidates should be able to:
    • explain that, in the presence of a catalyst, a reaction has a different mechanism, i.e. one of lower activation energy, and interpret this catalytic effect in terms of the Boltzmann distribution.
    • describe enzymes as biological catalysts (proteins) which may have specific activity.
slide49

Starter activity

Answer past paper question.

slide52

Effect of a catalyst

"A catalyst provides an alternative route for the reaction with a lower activation energy."

slide55

Types of catalysis

Heterogeneous catalysis:where the reactants and catalyst are in different physical states. Common in industrial processes.

Homogeneous catalysis: where the reactants and catalyst are in the same physical state. E.g. enzyme-catalysed reactions in cells.

slide59

Homogeneous catalysis

Peroxodisulphate and iodide ions

slide60

Enzymes

Enzymes are proteins that act as biological catalysts. Without them the reactions that make life possible would be too slow for life to exist.

slide64

AS Chemistry

Effect of concentration

slide65

Learning Objectives

  • Candidates should be able to:
  • explain qualitatively, in terms of collisions, the effect of concentration changes on the rate of a reaction.
slide66

Starter activity

In your pairs, can you complete the activity ‘Kinetics starter’?

slide69

Concentration – time graph

Amount of reactant

Time

slide73

Collecting gas

e.g. Zn(s) + 2HCl(aq) ZnCl2(aq) + H2(g)

slide74

Colorimeter

e.g. Na2S2O3(aq) + 2HCl(aq) 2NaCl(aq) + SO2(g) + S(s) + H2O(l)

slide78

Learning Objectives

  • Candidates should be able to:
    • describe and explain redox processes in terms of electron transfer and/or of changes in oxidation number (oxidation state)
slide79

Starter activity

In groups can you select the most appropriate term/s to describe the following chemical reactions?

slide87

h.

CH4 + Cl2 CH3Cl + HCl

slide88

Learning Objectives

Zn + CuOZnO + Cu

Definitions

  • Oxidation is gain of oxygen.
  • Reduction is loss of oxygen.
slide89

Blast Furnace

An oxidising agent is a substance which oxidises something else. It gains electrons and is reduced.

A reducing agent reduces something else. It loses electrons and is oxidised.

slide90

Organic reactions

  • Oxidation is loss of hydrogen.
  • Reduction is gain of hydrogen.
slide91

Universal definition

Consider the following reactions:

SO2 + H2O + HgO H2SO4 + Hg

SO2 + 2H2O + Cl2 H2SO4 + HCl

Clearly in both reactions there is an oxidation of SO2(g) to SO3(g) i.e. H2SO4.

Yet the second reaction does not involve oxygen!

We clearly need a more universal definition of oxidation and reduction.

slide92

Electron transfer

Definitions

  • Oxidation is loss of electrons.
  • Reduction is gain of electrons.

Mg + ZnCl2MgCl2 + Zn

Mg + CuCl2 MgCl2 + Cu

slide93

The oxidation state of an uncombined element is zero.

  • The sum of the oxidation states of all the atoms or ions in a neutral compound is zero.
  • The sum of the oxidation states of all the atoms in an ion is equal to the charge on the ion.
  • The more electronegative element in a substance is given a negative oxidation state. The less electronegative one is given a positive oxidation state.
  • Some elements almost always have the same oxidation states in their compounds:
slide95

AS Chemistry

Redox reactions

slide96

Learning Objectives

  • Candidates should be able to:
    • describe and explain redox processes in terms of electron transfer and/or of changes in oxidation number (oxidation state).
slide97

Starter activity

  • Can you work out the oxidation states of the transition metal elements in the following compounds?
    • KMnO4
    • K2Cr2O7
slide98

Naming compounds

  • SnO
  • SnO2
  • FeCl2
  • FeCl3
  • PbCl4
  • Cu2O
  • Mn(OH)2
  • NO2-
  • NO3-
  • SO32-
  • SO42-
  • MnO4-
  • CrO42-
  • VO3-
slide99

What is oxidised, what’s reduced?

  • 2ClO3-2Cl- + 3O2
  • 2Br- + 2H+ + H2SO4 Br2 + SO2 + 2H2O
  • 8I- + 8H+ + H2SO4 4I2 + H2S + 4H2O
  • I2 + SO3- + H2O 2I- + SO42- + 2H+
slide101

Writing ionic equations

In a reaction chlorine gas oxidises iron(II) ions to iron(III) ions. In the process, the chlorine is reduced to chloride ions.

Write a balanced equation for this redox process.

slide102

Writing ionic equations

Manganate(VII) ions, MnO4-, oxidise hydrogen peroxide, H2O2, to oxygen gas. The reaction is done with potassium manganate(VII) solution and hydrogen peroxide solution acidified with dilute sulphuric acid.

The manganate(VII) is reduced to Mn2+.

  • Write a balanced equation for this redox process.
slide103

Writing ionic equations

This technique can be used just as well in examples involving organic chemicals.

Potassium dichromate(VI) solution acidified with dilute sulphuric acid is used to oxidise ethanol, CH3CH2OH, to ethanoic acid, CH3COOH.

  • The Cr2O72-is reduced to Cr3+.
  • Write a balanced equation for this redox process.
slide104

AS Chemistry

Electrolysis

learning objectives
Learning Objectives
  • Candidates should be able to explain, including the electrode reactions, the industrial process of:
    • the electrolysis of brine, using a diaphragm cell;
    • the extraction of aluminium from molten aluminium oxide/cryolite; and
    • the electrolytic purification of copper.
starter activity
Starter activity

Complete task 1 on your worksheet.

slide110

Environmental concerns

Fort William

Anglesey

slide111

Purification of Copper

Printed circuit board

slide115

AS Chemistry

Dynamic Equilibrium

learning objectives1
Learning Objectives

Candidates should be able explain, in terms of rates of the forward and reverse reactions, what is meant by a reversible reaction and dynamic equilibrium.

starter activity1
Starter activity

In pairs, consider the reaction given below. How many facts about this reaction can you write down? Try to use the correct scientific terminology.

N2(g) + 3H2(g) 2NH3(g)

H –ve

dynamic equilibrium1
Dynamic Equilibrium

2 HI(g)H2(g) + I2(g)

slide129

AS Chemistry

Le Chatelier’s Principle

learning objectives2
Learning Objectives

Candidates should be able to state Le Chatelier’s Principle and apply it to deduce qualitatively (from appropriate information) the effects of changes in temperature, concentration or pressure, on a system at equilibrium

starter activity2
Starter activity

Question 1 from worksheet ‘Problems for 7.1’

slide132

Le Chatelier’s Principle

Put simply, Le Chatelier’s Principle states that:

If a system is at equilibrium, and a change is made in any of the conditions, then the system responds to counteract the change as much as possible.

slide133

Effect of concentration

Suppose you have an equilibrium established between four substances A, B, C and D.

What would happen if you changed the conditions by increasing the concentration of A?

slide134

Effect of pressure

What would happen if you changed the conditions by increasing the pressure?

slide135

Effect of temperature

What would happen if you changed the conditions by increasing the temperature?

slide136

AS Chemistry

Equilibrium constants

learning objectives3
Learning Objectives

Candidates should be able to

  • deduce expressions for equilibrium constants in terms of concentrations, Kc, and partial pressures, Kp.
  • deduce whether changes in concentration, pressure or temperature or the presence of a catalyst affect the value of the equilibrium constant for a reaction.
  • calculate the values of equilibrium constants in terms of concentrations or partial pressures from appropriate data.
  • calculate the quantities present at equilibrium, given appropriate data.
starter activity3
Starter activity

In pairs, consider the reaction given below. If you wanted to make as much ammonia as possible what conditions would you use?

N2(g) + 3H2(g) 2NH3(g)

H –ve

slide139

K – the equilibrium constant

Equilibrium constant

  • to provide a quantitative measure of the extent of a reaction;
  • to determine the position of equilibrium.
slide140

Kc

a A + b Bc C + d D

slide141

Calculating Kc

CH3CH2OH + CH3COOH CH3COOCH2CH3 + H2O

n(start) 1.0 1.0 0 0

n(equil.) 0.34 0.34 0.66 0.66

[ ] 3.4 3.4 6.6 6.6

Kc = 6.6 x 6.6 / 3.4 x 3.4 = 3.7 (no units)

slide142

Kp

Partial pressure

  • The total pressure exerted by a mixture of gases is the sum of the partial pressure of the gases.
slide143

Kp

Partial pressure and mole fraction

pA = xA x ptot

Partial pressure terms are expressed in SI units as Pa or kPa.

slide144

The equilibrium constant Kp

a A + b Bc C + d D

Kp =

slide145

Calculating Kp

PCl5(g) PCl3(g) + Cl2

n(start) 2.0 0 0

n(equil.) 0.8 1.2 1.2

x 0.25 0.375 0.375

P 166 249 249

Kp = 249 x 249 / 166 = 374 kPa

slide147

AS Chemistry

Equilibria of importance

learning objectives4
Learning Objectives

Candidates should be able to describe and explain the conditions used in the Haber process and the Contact process, as examples of the importance of an understanding of chemical equilibrium in the chemical industry.

starter activity4
Starter activity

The gases SO2, O2 and SO3 are allowed to reach equilibrium. The partial pressures of the gases are pSO2 = 0.050 atm, pO2 = 0.025 atm, pSO3 = 1.00 atm.

Find the values of Kp for the equilibria

a) SO2(g) + ½O2(g)SO3(g)

2SO2(g) + O2(g) 2SO3(g)

Comment on your results!

slide152

AS Chemistry

Acid and Base Equilibria

slide153

Starter Activity

A white solid is formed at X. Can you explain what is happening in this reaction? What words would you use to describe it?

slide154

Learning objectives:

  • Candidates should be able to:
    • show understanding of, and use the Bronsted-Lowry theory of acids and bases.
    • explain qualitatively the differences in behaviour between strong and weak acids and bases and the pH values of their aqueous solutions in terms of the extent of dissociation.
slide155

Properties of Acids

Tastesour

Have a pH < 7

Turn litmus red

Neutralise alkalis

Produce H+ in solution

React with metals to produce H2

React with carbonates to produce CO2

slide156

Early theory of acid behaviour

In 1884 Arrhenius stated that:

  • Acids are substances which produce hydrogen ions in solution.
  • Bases are substances which produce hydroxide ions in solution.
  • Allowed an explanation for neutralisation:
slide157

Limitations of Arrhenius

Easy to explain:

More of a challenge:

slide158

Limitations of Arrhenius

There is no solution!!!

slide159

Bronsted-Lowry

Acids are PROTON DONORS and bases are PROTON ACCEPTORS.

slide160

Conjugate acid-base pairs

A conjugate acid-base pair are related by the transfer of a proton.

slide161

Water

H2O +H2O

H3O+ +OH-

Amphoteric behaviour

slide162

Strong or weak acids

A strong acid is one which is virtually 100% ionised in solution.

A weak acid is only partially ionised in solution.