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.
•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.
Starter activity
Can you complete tasks 1 and 2 on your notes.
Exothermic
Endothermic
Simple calorimeter – flammable
liquids
Bomb calorimeter
Simple calorimeter – chemicals
in solution
AS Chemistry
Hess’ Law and
Enthalpy Cycles
Learning Objectives
Candidates should be able to:
•
apply Hess’ Law to construct simple energy cycles
•
carry out calculations involving such cycles
Starter activity
Can you write out definitions for ΔHθc and
ΔHθf ?
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.
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.
‘CRAP rule’
H = (Hc reactants) - (Hc products)
H = (Hf products) - (Hf reactants)
AS Chemistry
Bond Enthalpies
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.
Starter activity
Can you write equations for the ΔHθc and ΔHθf
of glucose (C6H12O6)?
Bond breaking and bond making
Chemical reactions involve bond breaking and bond
making.
Bond energy
The quantity of energy needed to break a particular bond
in a molecule is called the bond dissociation enthalpy
(H diss), 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.
Bond energy
H – H (g)

H(g)
+
H(g)
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
Bond
Number
broken
Number
formed
Average Molar
Bond Enthalpy
kJ mol-1
C-C
1
0
+347
C-H
5
0
+413
C-O
1
0
+358
O-H
1
6
+464
O=O
3
0
+498
C=O
0
4
+805
Here bond enthalpies are defined endothermically.
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:
Hc = +4728 + - 6004 = -1276 kJ mol-1
AS Chemistry
Kinetics
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.
Starter activity
Working in groups of 3, complete task 1.
Different rates of reaction
Seconds
dynamite exploding
magnesium and acid
Minutes
cake baking
Hours
fruit ripening
Days
Weeks
Months
Years
Decades
Centuries
Millennia
plants growing
rusting of iron
erosion of rock
crude oil forming
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.
There are five factors which can affect
the rate of a reaction:
1.
2.
3.
4.
5.
Surface Area
Concentration
Temperature
Use of a catalyst
Intensity of light
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).
Activation Energy -
the minimum
energy required for a reaction to occur.
A
B
Enthalpy
Enthalpy
Reactants
Products
Reactants
Products
Progress of Reaction
Progress of Reaction
C
D
Enthalpy
Enthalpy
Products
Reactants
Products
Progress of Reaction
Reactants
Progress of Reaction
Maxwell-Boltzmann Distribution
AS Chemistry
Effect of
Temperature
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.
Starter activity
Can you complete task 1?
Activation Energy
Increasing temperature
Increasing temperature
Temperature and Activation Energy
Temperature and Activation Energy
AS Chemistry
Catalysis
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.
Starter activity
Answer past paper question.
Activation Energy
Effect of a catalyst
Effect of a catalyst
"A
catalyst
provides
an
alternative
route for the
reaction with a lower activation
energy."
Alternative Pathway
Alternative Pathway
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.
Heterogeneous catalysis
Homogeneous catalysis
Peroxodisulphate and iodide ions
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.
‘Lock and Key’ mechanism
Enzymes are proteins
Active sites
AS Chemistry
Effect of
concentration
Learning Objectives
Candidates should be able to:
explain qualitatively, in terms of collisions, the
effect of concentration changes on the rate
of a reaction.
Starter activity
In your pairs, can you complete the activity
‘Kinetics starter’?
2H2O2 (aq)  2H2O (l) +
O2(g)
120
Volume of CO 2 /cm 3
Amount of product
140
100
80
60
40
20
0
Time
Amount of
product
Concentration – time graph
Amount of reactant
Concentration – time graph
Time
Effect of concentration
Zero-order reaction
Effect of
Pressure
Collecting gas
e.g. Zn(s) + 2HCl(aq)  ZnCl2 (aq) + H2(g)
Colorimeter
e.g. Na2S2O3(aq) + 2HCl(aq)  2NaCl(aq) + SO2(g) + S(s) + H2O(l)
Precipitation or ‘clock experiment’
Loss in mass
AS Chemistry
Redox
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)
Starter activity
In groups can you select the most appropriate
term/s to describe the following chemical
reactions?
Learning Objectives
Zn + CuO  ZnO + Cu
Definitions
•Oxidation is gain of oxygen.
•Reduction is loss of oxygen.
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.
Organic reactions
•Oxidation is loss of hydrogen.
•Reduction is gain of hydrogen.
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.
Electron transfer
Definitions
•Oxidation is loss of electrons.
•Reduction is gain of electrons.
Mg + ZnCl2  MgCl2 + Zn
Mg + CuCl2  MgCl2 + Cu
•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:
usual oxidation
exceptions
state
element
Group
metals
1
Group
metals
2
always +1
always +2
usually -2
except in peroxides and F2O (see
below)
Hydrogen usually +1
except in metal hydrides where it is
-1 (see below)
Oxygen
Fluorine
always -1
Chlorine
usually -1
except in compounds with O or F
(see below
AS Chemistry
Redox reactions
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).
Starter activity
Can you work out the oxidation states of the
transition metal elements in the following
compounds?
•KMnO4
•K2Cr2O7
Naming compounds
•SnO
•NO2-
•SnO2
•NO3-
•FeCl2
•SO32-
•FeCl3
•SO42-
•PbCl4
•MnO4-
•Cu2O
•CrO42-
•Mn(OH)2
•VO3-
What is oxidised, what’s reduced?
1.
2ClO3-
2.
2Br-
3.
8I-
4.
I2

+
+
+
2Cl-
+
3O2
2H+
+
H2SO4

Br2
+
SO2
+
2H2O
8H+
+
H2SO4

4I2
+
H2S
+
4H2O
SO3-
+
H2 O

2I-
+
SO42-
+
2H+
What is oxidised, what’s reduced?
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.
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.
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.
AS Chemistry
Electrolysis
Manufacturing Chlorine using a
Diaphragm Cell
•
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.
Complete task 1 on your worksheet.
Extraction of Aluminium
Environmental
concerns
Fort William
Anglesey
Purification of Copper
Printed circuit
board
Effect of impurities on conductivity
AS Chemistry
Dynamic
Equilibrium
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.
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
2 HI(g)  H2(g) + I2(g)
AS Chemistry
Le Chatelier’s
Principle
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
Question 1 from worksheet ‘Problems for 7.1’
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.
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?
Effect of pressure
What would happen if you changed the conditions
by increasing the pressure?
Effect of temperature
What would happen if you changed the conditions by
increasing the temperature?
AS Chemistry
Equilibrium
constants
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.
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
K – the equilibrium constant
Equilibrium constant
•to provide a quantitative measure of the extent of
a reaction;
•to determine the position of equilibrium.
Kc
aA +
bB
cC +
c
d
[C ] [ D]
Kc 
a
b
[ A] [ B]
dD
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)
Kp
Partial pressure
The total pressure exerted by a mixture of gases is the
sum of the partial pressure of the gases.
Kp
Partial pressure and mole fraction
pA = xA x ptot
Partial pressure terms are expressed in SI units as
Pa or kPa.
The equilibrium constant Kp
aA +
Kp =
bB
cC +
c
d
p(C ) p( D)
p( A) a p( B) b
dD
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
Change in temperature
∆H for
reaction
Change in
Temp.
Shift of
Equilibrium
Yield of
Product

Decrease
Equilibrium
constant
Exo
Increase
Exo
Decrease

Increase
Increase
Endo
Increase

Increase
Increase
Endo
Decrease

Decrease
Decrease
Decrease
AS Chemistry
Equilibria of
importance
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.
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)
b) 2SO2(g) + O2(g)
SO3(g)
2SO3(g)
Comment on your results!
Haber process
Process
Main product
Contact
Sulphuric acid
Haber
Ammonia
Paints, detergents and Fertilisers,
explosives,
soaps, fertilisers and dyes. nitric acid, polymers.
Balanced equation for 1.S + O2  SO2
3H2 + N2
2NH3
2. 2SO2 + O2
2SO3
main reaction/s
Main uses of product
Catalyst
3. SO3 + H2SO4  H2S2O7
4. H2S2O7 + H2O  2H2SO4
Vanadium(V) oxide
Is
the
equilibrium Exothermic
reaction exothermic or
endothermic?
Porous iron (with metal
oxide promoter)
Exothermic
Optimum conditions
Low temperature
High pressure
for highest yield
400 – 600oC
Actual conditions used
1-2 atm
Low temperature
High pressure
400 – 500oC
200 atm
Why are these
conditions chosen?
Catalyst ineffective at low T.
High pressure uneconomical.
Reaction slow at low T.
High pressure too costly.
Points of interest
Reaction mixture cooled
after each exothermic
stage.
Vast excess of air.
Ammonia removed as
produced.
N2 and H2 recycled.
Waste heat re-used.
AS Chemistry
Acid and Base
Equilibria
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?
Learning objectives:
Candidates should be able to:
•show understanding of, and use the BronstedLowry 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.
Properties of Acids
Taste sour
Turn litmus
red
Have a pH
<7
Neutralise
alkalis
React with
metals to
produce H2
React with
carbonates to
produce CO2
Produce H+
in solution
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:
Limitations of Arrhenius
Easy to explain:
More of a challenge:
Limitations of Arrhenius
There is no solution!!!
Bronsted-Lowry
Acids are PROTON DONORS and bases are
PROTON ACCEPTORS.
Conjugate acid-base pairs
A conjugate acid-base pair are related by the
transfer of a proton.
Water
H2O
+ H 2O
Amphoteric behaviour
H3O+ + OH-
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.