YEAR 12 CHEMISTRY
Term 2 - 2012
REVERSIBLE REACTIONS
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Some chemical reactions are reversible.
Water can become liquid, solid or gas depending
on the circumstances.
We saw in the previous unit that the reaction to
form an ester could either move forward to
produce the ester and water or in the right
conditions move backward to produce the alcohol
and the carboxylic acid.
This reaction is known as a reversible reaction.
REVERSIBLE REACTIONS
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Without a something push it to completion this
reaction will reach a point where it contains both
products and reactants.
When the reaction reaches this point both
reactions are occurring at equal rates. This
means there is no overall change in the amount
of each item present. This is called dynamic
equilibrium.
We write this reaction as:
Reactant ↔ Product
EQUILIBRIUM
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Equilibrium is the term given to an object in a
state of balance.
Static equilibrium is like a tug of war where both
sides are equal strength. The rope does not move
because the force on both sides is equal
Dynamic equilibrium is like a sports game, there
is always a certain number of players on the field
but the players themselves can change.
REVERSIBLE REACTIONS
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Chapter 1 Page 203
REVERSIBLE REACTIONS
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For reversible reactions we have:
Forward reactions - the reaction going as written
from left to right.
Reverse or backward reactions - the reaction
going opposite to the way it is written from right
to left.
What is the product of the reverse reaction for
the following reversible reaction?
Pb(NO3)2(s) ↔ Pb2+ + 2NO3-(aq)
REVERSIBLE REACTIONS
Reversible Reaction - A reaction which can go
forward or backward depending on the
circumstances.
 Dynamic Equilibrium - The state where the
concentration of products and reactants in a
reaction remains stable, the forward and reverse
reactions are occurring at the same rate.
 Forward Reaction - A reversible reaction
which is occurring from left to right.
 Backward / Reverse Reaction - A reversible
reaction which is occurring from right to left.
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TYPES OF REVERSIBLE REACTIONS
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All precipitation reactions are reversible
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Potassium chromate ↔ Potassium dichromate
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Esters
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Gases
HOW DOES EQUILIBRIUM OCCUR?
Ag+(aq) + Fe2+(aq) ↔ Fe3+ + Ag(s)
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When the reaction above begins we have large
concentrations of Ag+ and Fe2+.
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Reactants decrease as products increase
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As products increase they begin to react
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Eventually reach state of equilibrium where both
reactions are occurring at the same rate.
Br2 (g) + 2NO(g) ↔ 2NOBr (g)
CATALYSTS
How do catalysts affect equilibrium?
 Draw graph:
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EXERCISES
EXERCISES
EXERCISES
EXERCISES
EXERCISES
CARBON DIOXIDE EQUILIBRIUM
Read page 207- 208.
Write down equation of dissolution of CO2 in H2O.
What is the relationship between solubility and
pressure for gases in liquids.
Does the reaction lie to the right or left?
SOLUBILITY AND PRESSURE FOR GASES IN
LIQUIDS
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As the pressure increases ↑ solubility increases ↑
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As pressure decreases ↓ solubility decreases ↓
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As solubility increases ↑ temperature decreases ↓
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As solubility decreases ↓ temperature increases ↑
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This is the opposite of solids and liquids in which
temperature normally increases ↑ as solubility
increases ↑.
EQUILIBRIUM POSITION
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The extent to which the reaction has gone in the
forward or reverse direction.
Lies to the left
Reverse reaction is favoured
Most of reactant is still resent
Small amount of product
 Lies to the right
Forward reaction is favoured
Most of the reactants has converted to product

LE CHATELIER’S PRINCIPLE
If a system in
equilibrium is
disturbed, the system
adjusts itself so as to
minimize the
disturbance
LE CHATELIER’S PRINCIPLE
Disturbance to a system may include:
Concentration
 Pressure
 Temperature
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Note a system at equilibrium is not disturbed by
adding more solid to it. This is because the
concentration of ions is not changed when a solid
is added.
 Concentration not amount effects equilibrium
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CARBON DIOXIDE AND CARBONIC ACID
Read page 210 – 211.
Predict the direction of the equilibrium
if:
a) The plunger was pulled creating
more volume
b) The temperature of the system
decreased
c)
Adding sodium hydroxide to the mix.
FORCING REACTIONS TO COMPLETION
Water liquid and gas equilibrium:
FORCING REACTIONS TO COMPLETION
H2O(l)  H2O(g)
Water in a terrarium reaches
equilibrium because it is a closed
system water evaporates while vapour
condenses.
Water in wet clothes attempts to
establish equilibrium however as liquid
evaporates, it is carried by wind and
diffuses into the atmosphere and so
dryness will eventually occur.
FORCING REACTIONS TO COMPLETION
When synthesizing chemicals, chemists
may wish to push a reversible reaction
to completion to obtain the maximum
amount of a product.
A common way to force reactions to
completion is to remove a product as it
is produced. Addition of an excess of
cheap or common reactant is another
way.
EXERCISES PAGE 212
EXERCISES PAGE 212
EXERCISES PAGE 212
EXERCISES PAGE 212
EXERCISES PAGE 212
EXERCISES PAGE 212
EXERCISES PAGE 212
EXERCISES PAGE 212
EXERCISES PAGE 212
EXERCISES PAGE 212
CONDITION FOR EQUILIBRIUM
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There is a quantitative relationship between
reactants and products at equilibrium.
Using
When the reaction is at equilibrium at a constant
temperature the expression [I3-] / [I2] [I-] has a
constant value.
CONDITION FOR EQUILIBRIUM
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See table 2.1 page 240.
CONDITIONS FOR EQUILIBRIUM
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Table shows a clear relationship between the
reactants and products in an equilibrium
reaction.
For any equilibrium reaction there is a function
of the concentration of the species which has a
constant value at equilibrium
This constant is given the symbol K and known
as the equilibrium constant.
EQUILIBRIUM CONSTANT
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In an equilibrium reaction where
aA + bB ↔ cC + dD
When the reaction is at equilibrium the
expression:
[𝐶]𝑐 [𝐷]𝑑
[𝐴]𝑎 [𝐵]𝑏
Has a constant value, regardless of the starting
concentrations of the substances involved.

This is called the equilibrium constant K
EQUILIBRIUM CONSTANT VS REACTION QUOTIENT
[𝐶]𝑐 [𝐷]𝑑
=K
𝑎
𝑏
[𝐴] [𝐵]
The above equation at constant temperature is
known as the equilibrium expression
[𝐶]𝑐 [𝐷]𝑑
Q=
[𝐴]𝑎 [𝐵]𝑏
Alternatively the above equation at constant
temperature is known as the reaction quotient
EQUILIBRIUM CONSTANT VS REACTION QUOTIENT
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The equilibrium constant [K] is the relationship
between the products and reactants at
equilibrium.
The reaction quotient [Q] is the relationship
between the products and reactants at any given
point during the reaction.
When Q = K the reaction is at equilibrium.
 If Q ≠ K the reaction is not at equilibrium.
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EQUILIBRIUM CONSTANT VS REACTION QUOTIENT
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It is possible to use Q to determine which direction a
reaction if occurring.
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If Q < K the reaction goes from left to right until Q = K
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If Q > K the reaction goes from right to left until Q = K
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If Q = K the reaction is at equilibrium.
REACTION QUOTIENT
RULES FOR WRITING EQUILIBRIUM
EXPRESSIONS
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Both Q and K are always written with the
products in the numerator and the reactants in
the denominator.
K or Q =
[𝑝𝑟𝑜𝑑𝑢𝑐𝑡𝑠]
[𝑟𝑒𝑎𝑐𝑡𝑎𝑛𝑡𝑠]
K or Q =
RULES FOR WRITING EQUILIBRIUM
EXPRESSIONS
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Coefficients of products or reactants are written
as powers for that product or reactant.
eg:
UNITS FOR EQUILIBRIUM CONSTANTS
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Reaction quotients have units.
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mol/L
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To calculate the value of the unit:
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Cancel units above and below the division
symbol. If the remainder is below the division
symbol the units will be in mol/L-1 etc. If the
remainder is above the division symbol units will
be in mol/L1 etc.
EXERCISES
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Page 244
Questions 1 - 4
TEMPERATURE AND EQUILIBRIUM
CONSTANT
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The equilibrium constant K only remains stable if
temperature is stable. If temperature changes then
the value of the equilibrium constant will change.
If reaction is exothermic and temperature increases:
Reaction will move from right to left
K value will decrease.
If reaction is endothermic and temperature increases:
Reaction will move from left to right
K value will increase.
TEMPERATURE AND EQUILIBRIUM
CONSTANT
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For an exothermic reaction:
K decreases as temperature increases.
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For an endothermic reaction:
K increases at temperature increases.
EXERCISES
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Questions 16 and 17
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Page 252.
NOTES FOR EQUILIBRIUM
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Always use coefficients as written.
2H2O + O2 ↔ 2H2O2
Will have a difference equilibrium constant to:
1H2O + ½O2 ↔ 1H2O2
If concentrations are not given they must be in
mol/L
EQUILIBRIUM IN GASES
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Recall: In gas at constant temperatures pressure
is proportional to concentration.
PV = nRT
 Where P = pressure
V = volume
n = number of moles
R = The gas constant
T = temperature.
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EQUILIBRIUM IN GASES
PV = nRT
 This can be rearranged:
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P=
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𝑛
𝑉
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
𝑛
RT
𝑉
= concentration of the gas.
If we were to compress gas in a system to ½ the
volume then we would have double the pressure
and double the concentration.
i.e. same moles of gas in half the volume
EQUILIBRIUM IN GASES
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To increase the pressure in a gas we can:
Adding more gas to a given volume or
 Keeping amount of gas stable and decreasing
volume.
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Therefore:
As pressure increases, concentration increases
and volume decreases.
 As pressure decreases, concentration decreases
and volume increases
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GAS EQUILIBRIUM AND LE CHATELIER
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If pressure increases then equilibrium is
disturbed and the reaction will go in the direction
which minimises the disturbance.
If there is a decrease in volume the reaction will
go in the direction which produces less moles of
gas. e.g.
If pressure is increased which direction will the
following reaction go?
PCl3(g) + Cl2(g) ↔ PCl5(g)
GAS EQUILIBRIUM AND LE CHATELIER
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If moles of gas of products and reactants are
equal, then changing pressure will not disrupt
equilibrium.
Concentration of products and reactants will
remain the same.
CO(g) + NO2(g) ↔ CO2(g) + NO(g)
GAS EQUILIBRIUM AND LE CHATELIER
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If moles of gas of products and reactants are
equal, then changing pressure will not disrupt
equilibrium.
Concentration of products and reactants will
remain the same.
CO(g) + NO2(g) ↔ CO2(g) + NO(g)
CALCULATIONS USING K
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It is difficult to measure equilibrium constant.
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Why?
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If we remove substances the reaction quickly
moves to minimise the disturbance because of
LeChatelier.
To counter this chemists often use the absorption
of light or pH meters to measure Q and K values
as these measure without disturbing the system.
CALCULATIONS USING K
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Do we need to know the concentrations of every
species to be able to calculate K?
MEASUREMENT OF EQUILIBRIUM
CONSTANTS
Example 1
2SO2 (g) + O2 (g)  2SO3 (g)
At equilibrium at 900K concentration of SO3 was 9
times the concentration of SO2 when the equilibrium
concentration of oxygen was 0.068 mol/L. Calculate
the equilibrium constant.
Set out the information that you know from the
problem.
At equilibrium:
[O2] = 0.068 mol/L
[SO3]/[SO2] = 9 so
[SO3] = 9
[SO2] = 1
EQUILIBRIUM EXPRESSION:
K=
2
[ SO3 ]
2
[ SO2 ] [O2 ]
2
K=
[9]
2
[1] [0.068]
K = 1.2 x 103 (mol/L)-1
Example 2
PCl5 (g)   PCl3 (g) Cl2 (g)
0.0100 mole phosphorus pentachloride was placed in a
1.00L flask at 523K. At equilibrium the concentration
of chlorine was 0.0083 mol/L. Calculate the
equilibrium constant for the reaction.
Take the information you know from the equation:
The ratio for the equation is 1:1:1. At completion 1
mole of PCl5 forms 1 mole of PCl3 and 1 mole of Cl2. At
equilibrium if we have 0.0083 mole of Cl2 then we also
have 0.0083 mole of PCl3 .
Cl2 = PCl3 = 0.0083 mole
The PCl3 and Cl2 came from the PCl5 therefore 0.0083
mole of PCl5 must have been used. If we initially had
0.0100 mole of PCl5 then at equilibrium we have:
0.0100 – 0.0083 = 0.0017 mole
SETTING OUT
PCl5 (g)   PCl3 (g) Cl2 (g)
Initially:
0.0100 mol/L
0
At equilibrium: 0.0100 – 0.0083
0.0017
0
0.0083
0.0083
0.0083
0.0083
K=
[ PCl3 ][Cl2 ]
[ PCl5 ]
K=
[0.0083][0.0083]
[0.0017]
K = 0.041 mol/L
USE OF K
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K tells us the position of equilibrium and the
concentration of species at equilibrium.
If K is large e.g. > 103 equilibrium lies to the
right and the reaction favours the products.
If K is small e.g. < 10-3 equilibrium lies to the left
and the reaction favours the reactants.
If K is in the centre there are similar amounts of
products and reactants.
USING Q AND K
N2O4(g) ↔ 2NO2
The equilibrium constant is 0.48 at 100°C.
0.1mol N2O4 and 0.25 mol NO2 were placed in a
1.0L flask at 100°C. Is the mixture at equilibrium?
If not which direction will the reaction proceed.
USING Q AND K
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Just like any mathematical expression the
expression for K can be rearranged.
If we know K and some other concentration in
the expression we can rearrange it to determine
the unknown values.
USING Q AND K
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At 1000K the equilibrium constant for the
decomposition of phosgene(COCl2) into carbon
monoxide and chlorine is 0.40.
A sample of phosgene was placed in an evacuated
container and heated to 1000K. When
equilibrium was reached, the concentration of
carbon monoxide was 0.24 mol/L. Calculate the
equilibrium concentration of phosgene.
USING Q AND K
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At 1000K the equilibrium constant for the
decomposition of phosgene(COCl2) into carbon
monoxide and chlorine is 0.40.
In a second experiment another sample of
phosgene was brought to equilibrium at 1000K,
the equilibrium concentration of phosgene was
0.18mol/L. Calculate the equilibrium
concentration of chlorine.
EXERCISES
Questions 6 - 14
Page 249-250
SOLUBILITY EQUILIBRIA
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Remember when a solid ionic compound is in
solution with that solid equilibrium occurs.
If the solid has low solubility there is a
relationship between the concentration of ions at
equilibrium.
e.g. PbSO4(s) ↔ Pb2+ + SO42-(aq)
At equilibrium:
[Pb2+] [SO42-] = constant
SOLUBILITY EQUILIBRIA
[Pb2+] [SO42-] = constant
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Note that there is no term for the solid as the
concentration for a solid is constant.
When a pure substance, solid or liquid is used its
concentration its concentration is not included in
the equilibrium constant expression for the
reaction.
An increase in the amount of pure solid does not
effect equilibrium
SOLIDS IN KSP
If
𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛
𝑠𝑜𝑙𝑖𝑑
= Ksp
And the concentration of the solid is constant than
𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛
𝑐𝑜𝑛𝑠𝑡𝑎𝑛𝑡
= Ksp and
Solution = [Ksp][solid]
SOLUBILITY EQUILIBRIA
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The equilibrium constant in this case is called
the solubility product and written as Ksp
[Pb2+] [SO42-] = Ksp
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The reaction for Ksp is always written with the
solid on the left and the solution on the right.
Why?
Solid ↔ Solution
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As with any equilibrium expression coefficients
are written as powers for the relevant terms.
SOLUBILITY EQUILIBRIA
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Solubility expression
MxAy(s) ↔ xMb+(aq) + yAc-(aq)
[Mb+]x [Ac-]y = Ksp
Ionic product expression
[Mb+]x [Ac-]y = Q
If IP < Ksp solid dissolves until =
 If IP > Ksp solid precipitates until =
 If IP = Ksp sys is at equilibrium
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EXERCISES
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Page 259
Exercises 26 - 29
REVISION
Concepts we have covered so far:
 Dynamic equilibrium
 Reversible reactions
 Catalysts
 Kc
 Ksp
Q
 Le Chatelier
 Temperature and Equilibrium
 Equilibrium in Gases
INTRODUCTION TO ACIDS AND BASES
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Acid is a substance which in solution produces H+
ions or H3O+ ions (this is the more correct
terminology)
Base is a substance which contains the O2- ion or
OH- ion or in solution produces the OH- ion.
Soluble base is called an alkali
PROPERTIES OF ACIDS
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Acids have a sour taste
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Acids sting or burn the skin
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In solution acids conduct electricity
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Acids turn blue litmus red
PROPERTIES OF BASES
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Bases have a soapy feel
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Bases have a bitter taste
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In solution bases are good conductors of
electricity
Bases turn red litmus blue
ACID BASE REACTIONS
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Acids react with bases in neutralisation reactions
to form salts.
A salt is an ionic compound formed when a base
reacts with an acid. They are not limited to NaCl
A neutralisation reaction occurs when the H+
ions in the acid react with the OH- ions in the
base to produce H2O
ACID-BASE REACTIONS
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We can write an acid base reaction in 3 ways:
Neutral species equation:
HCl + NaOH  H2O + NaCl
Complete ionic equation:
H+ + Cl- + Na+ + OH-  H2O + Na+ + Cl-
Net ionic equation:
H+ + OH-  H2O
ACID-BASE REACTIONS
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
Because the Na+ and the Cl- do not actually take
part in the reaction we call them spectator ions
and do not always have to include them in the
reaction when writing it.
In general in a neutralisation reaction
Acid + base  salt + water
There are some exceptions to this rule.
NAMING SALTS
When naming salts there are some general rules:
Cation is named first  anion then comes from
the name of the acid.
 Hydrohalic acids e.g. HCl become halide salts eg
NaCl
 Oxyacids e.g. H2SO3 (carbonic acid) the ic at the
end of the name becomes ate carbonic acid 
calcium carbonate CaCO3
 Ous acids e.g. nitrous HNO2 become ites e.g.
sodium nitrite NaNO2
 Anions formed from oxyacids are oxy anions.
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TABLE OF ACIDS AND ANIONS PG 217
QUESTIONS
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Page 217
Questions 8 - 13
ACIDIC AND BASIC OXIDES
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Oxides are compounds that can show either
acidic or basic properties.
An acidic oxide either:
 Reacts with water to form an acid or
 Reacts with bases to form salts or
 Both
A basic oxide:
 Reacts with acids to form salts
 Does not react with alkali solutions (NaOH or
KOH)
ACIDIC AND BASIC OXIDES


Acidic oxides are oxides of non-metals. They are
covalent compounds generally found at the top
right of the periodic table.
Basic oxides are oxides of metals. They are ionic
compounds generally found at the left of the
periodic table.
QUESTIONS
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Page 219
Questions 14 and 15