Chemical Equilibrium-8.pptx

Chapter 8 Equilibria
8.1 Reversible reactions and equilibrium
8.2 Changing the position of equilibrium
8.3 Equilibrium expressions and the equilibrium constant, Kc
8.4 Equilibria in gas reactions: the equilibrium constant, Kp
8.5 Equilibria and the chemical industry
8.6 Acid–base equilibria
8.7 Indicators and acid–base titrations
Table of Contents

• Tell about some processes that they go to completion.
Warm up
• Describe what reversible means.
• Find a synonym for reversible.
Equilibria

Reactions
Irreversible Reversible
The one goes completion.
Explosions, burning
Processes and decay of
leaves, etc.
The one that can occur in both
the forward and the reverse
directions.
Formation of NH3 , HCl,H2SO4.
Chemical Equilibrium

Reversible reactions
The reaction that can occur in both the forward and the
backward directions are called reversible reactions.
Some reactions can be reversed.
For example,
when blue, hydrated copper(II) sulfate is heated, it loses its
water of crystallization and changes to white anhydrous
copper(II) sulfate as shown below;
This is called the forward reaction.
When water is added to anhydrous copper(II) sulfate, the
reaction is reversed.
This is called the backward reaction.
These 2 reactions in the same equation by using two half
arrows: ⇌

Reversible reactions

1. Physical Equilibrium
• An equilibrium established by changing states of matter is
called physical equilibrium.
H2O(l)  H2O(g)

Physical Equilibrium
H2O(l)  H2O(g)
Vaporization and condensation of water in a closed
container at constant temperature is a physical equilibrium.
When rate of vaporization and rate of condensation become
equal an equilibrium state is established.
At equilibrium the amounts of water vapor and liquid water
remain constant. But evaporation and condensation
continues on microscopic scale.
This is called dynamic equilibrium.

2. Chemical equilibrium
A state of balance in which the rate of a forward reaction
equals the rate of the reverse reaction and the concentrations
of products and reactants remain unchanged is called
chemical equilibrium.
For example;
The reaction between hydrogen and iodine carried out in a
sealed glass tube at 400 °C:
2 2
H ( ) I ( ) 2HI( )


 

g g g
Molecules of HI are breaking down to H2 & I2 at the same
rate as hydrogen and iodine molecules are reacting together
to form hydrogen iodide (see Figure 8.3).

2. Chemical Equilibrium
Rate Comparison for H2(g) + I2(g)  2HI(g)
When the forward rate and the reverse rate are equal,
the system is at chemical equilibrium.

Rate Comparison for H2(g) + I2(g)  2HI(g)
Concentration, M
Time
H2 or I2
HI
teq
When the reaction reaches
equilibrium concentration of
substances become constant.

Characteristics of equilibrium
1. It is dynamic
2. Forward and reverse reactions occur at the same rate
3. It must be in a closed system.
4. The concentrations of reactants and products remain
constant at equilibrium
Example 1
a. H2(g) + Cl2(g)  2HCl(g) (homogeneous)
b. Mg(s) + 2HCl(aq)  MgCl2(aq) + H2(g) (heterogeneous)

Characteristics of equilibrium
1) Dynamic
The dynamic equilibrium means that the molecules or ions of
reactants and products are continuously reacting. Reactants are
continuously being changed to products and products are
continuously being changed back to reactants.
2)The forward and backward rxns occur at the same rate
At equilibrium, the rate of the forward reaction equals the rate of
the backward reaction. Molecules or ions of reactants are
becoming products, and those in the products are becoming
reactants at the same rate

3) The concentrations of reactants and products remain
constant at equilibrium
The concentrations remain constant because, at equilibrium,
the rates of the forward and backward reactions are equal. The
equilibrium can be approached from two directions. For
example, in the reaction

4) Equilibrium requires a closed system
A closed system is one in which none of the reactants or
products escapes from the reaction mixture. In an open system
some matter is lost to the surroundings.
Given figure shows the difference between a closed system and
an open system when calcium carbonate is heated at a high
temperature in a strong container.
Many chemical reactions can be studied without placing them in
closed containers. They can reach equilibrium in open flasks if
the reaction takes place entirely in solution and no gas is lost.

8.2 Changing the position of equilibrium
Position of equilibrium
The position of equilibrium refer to the relative amounts of
products and reactants present in an equilibrium mixture.
If a system in equilibrium is disturbed (e.g. by a change in
temperature) and the concentration of products increases
relative to the reactants, we say that the position of equilibrium
has shifted to the right.
If the concentration of products decreases relative to the
reactants, we say that the position of equilibrium has shifted to
the left.

Factors Affecting Equilibrium
(Le Chatelier’s Principle)

Le Chatelier’s Principle
Changes in both concentration and
temperature affect the position of equilibrium.
When any of the reactants or products are
gases, changes in pressure may also affect
the position of equilibrium.
French chemist Henri Le Chatelier's (1850–1936) observed how
these factors affect the position of equilibrium.
He put forward a general rule, known as Le Chatelier’s principle.
If one or more factors that affect a dynamic
equilibrium is changed, the position of
equilibrium moves to minimise this change.

How does change in concentration affect the
position of equilibrium?
1) When the concentration of one or more of the reactants is
increased.
o the system is no longer in equilibrium
o the position of equilibrium moves to the right to reduce the
effect of the increase in concentration of reactant
o more products are formed until equilibrium is restored.
2) When the concentration of one or more of the products is
increased.
o the system is no longer in equilibrium
o the position of equilibrium moves to the left to reduce the
effect of the increase in concentration of product
o more reactants are formed until equilibrium is restored.

For example, look at the reaction.
What happens when we add more ethanol?
The concentration of ethanol is increased.
According to Le Chatelier’s principle, some of the ethanol must
be removed to reduce the concentration of the added ethanol.
The position of equilibrium shifts to the right.
More ethanol reacts with ethanoic acid and more ethyl ethanoate
and water are formed

What happens when we add more water?
The concentration of water is increased.
According to Le Chatelier’s principle,
some of the water must be removed to reduce the concentration
of the added water.
The position of equilibrium shifts to the left.
So more water reacts with ethyl ethanoate and more ethanoic
acid and ethanol are formed.
A natural effect of how change in concentration affects the
position of equilibrium is shown by the formation of stalactites
and stalagmites (Figure 8.7).

A natural effect of how change in concentration affects the
position of equilibrium is shown by the formation of stalactites
and stalagmites.
Figure 8.7: Stalactites and stalagmites are
formed as a result of water passing through
rocks containing calcium carbonate.
The solution running through these rocks
Contains water, dissolved carbon dioxide and
calcium Hydrogen carbonate.
CaCO3(s) + H2O(l) + CO2(aq) ⇌ Ca(HCO3)2(aq)
When droplets of this mixture are formed on the roof of the cave,
some of the carbon dioxide in the droplets escapes into the air. The
position of equilibrium shifts to the left and calcium carbonate is
deposited.

The effect of pressure on the position of equilibrium
Change in pressure only affects reactions where gases are
reactants or products. The molecules or ions in solids and
liquids are packed closely together and cannot be compressed
very easily. In gases, the molecules are far apart (Figure 8.8).
The pressure of a gas is caused by the molecules hitting the
walls of the container. Each molecule in a mixture of gases
contributes towards the total pressure. So, at constant
temperature, the more gas molecules there are in a given
volume, the higher the pressure.

Figure 8.9 shows what happens when we increase the pressure
on the reaction represented by

For example,
consider the reaction
2SO2(g) + O2(g) ⇌ 2SO3(g)
There are three moles of gas molecules on the left of the
equation and two on the right.
What happens when we increase the pressure?
The molecules are closer together, because the pressure is
higher.
According to Le Chatelier’s principle, the reaction must shift
in the direction that reduces the number of molecules of
gas.
The position of equilibrium shifts to the right.
So more SO2 reacts with O2 to form SO3

The effect of temperature on the position of
equilibrium
The decomposition of hydrogen iodide is an endothermic
reaction
2HI(g) ⇌ H2(g) + I2(g) ΔHr = +9.6 kJ mol−1The
effect of temperature on the equilibrium concentration of
hydrogen iodide and hydrogen at equilibrium for the forward
reaction is shown in Table 8.3.
You can see from Table 8.3 that, as the temperature
increases, the concentration of product increases. The
position of equilibrium shifts to the right. We can explain this
using Le Chatelier’s principle:

o An increase in temperature increases the energy of the
surroundings
o According to Le Chatelier’s principle, the reaction will go
in the direction that resist the increase in energy
o The reaction will go in the direction in which energy is
absorbed, which is the endothermic reaction
o The position of equilibrium shifts to the right, producing
more H2 and I2

If an increase of temperature favours an endothermic
reaction, a decrease of temperature favours an exothermic
reaction. This means that:
o A decrease in temperature decreases the energy of the
surrounding.
o according to Le Chatelier’s principle, the reaction will go
in the direction that opposes the decrease in energy.
o so the reaction will go in the direction in which energy is
released, which is the exothermic reaction.

Table 8.4 summarises the effect of temperature changes on
the position of equilibrium for endothermic and exothermic
reactions.
HOME WORK

Do catalysts have any effect on the position of
equilibrium
A catalyst is a substance that increases the rate of a
chemical reaction. Catalysts reduce the time taken to reach
equilibrium, but they have no effect on the position of
equilibrium once this is reached. This is because they
increase the rate of the forward and reverse reactions
equally.

8.3 Equilibrium expressions and the equilibrium
constant, Kc
Important
• Remember square brackets, e.g. [Cl2], are used here to
show the concentration of the substance inside the
brackets.
• When writing equilibrium expressions we assume that
[X][Y] means [X] × [Y]

constant, Kc
Equilibrium expression
When hydrogen reacts with iodine in a closed tube at 500 K,
the following equilibrium is set up.
Table 8.5 shows the relationship between the equilibrium
concentrations of H2, I2 and HI. The square brackets in the
last column refer to the concentration, in mol dm−3, of the
substance inside the brackets. The results are obtained as
follows:

Equilibrium Constant Expression
aA + bB cC + dD
Ratef = kf[A]a
[B]b
Rater = kr[C]c
[D]d
kf[A]a
[B]b
= kr[C]c
[D]d
at equilibrium,
Ratef = Rater
=
kf
kr
[C]c
[D]d
[A]a
[B]b
Kc =
[C]c
[D]d
[A]a
[B]b
Kc =
[products]
[reactants]
In general,

• Solid and liquid substances are not included in
Kc expressions.
Example
Write Kc expressions of the following
equilibrium reactions.
a. N2H4(g)  2NO2(g)
b. 2Mg(s) + O2(g)  2MgO(s)
c. 2NH3(g)  N2(g) + 3H2(g)
d. Ag+(aq) + Cl-(aq)  AgCl(s)
Kc =
[NO2]2
[N2H4]
Kc =
1
[O2]
Kc =
[N2][H2]3
[NH3]2
Kc =
[Ag+
] [Cl-
]
1
Solution

constant, Kc
Equilibrium constant, Kc or Kp
A constant calculated from the equilibrium expression for a
reaction. It can be in term concentrations, Kc or partial
pressure, Kp.
What are the units of Kc?
The unit of equilibrium constant Kc is mol/liter.
Partial pressure
The pressure exerted by an individual gas is called partial
pressure.
The equilibrium constant Kp may be expressed as;
Kp = ( pc)c (pD)d
(pA)a (pB)b

constant, Kp
Partial pressure
The pressure exerted by an individual gas is called partial
pressure.
The total pressure of a gas equals the sum of the partial
pressures of the individual gases.
ptotal = pA + pB + pC …where pA, pB, pC are the partial
pressures of the individual gases in the mixture

constant, Kp
In general;
The equilibrium constant Kp may be expressed as;
Kp = ( pc)c (pD)d
(pA)a (pB)b
For example, the equilibrium expression for the reaction

What are the units of Kp?
The units of pressure are Pascals, Pa. The units of Kp depend
on the form of the equilibrium expression.

constant, Kp
Partial pressure and mole fractions
The number of moles of gas is proportional to the volume of
the gas at constant temperature.
So it follows that the partial pressure of a gas is proportional
to its concentration.
The mole fraction of a gas is given by the relationship:

An understanding of equilibrium is important in the chemical
industry. Equilibrium reactions are involved in some of the
stages in the large-scale production of ammonia, sulfuric
acid and many other chemicals.
Equilibrium and ammonia production
The synthesis of ammonia is carried out by the Haber
process. The equilibrium reaction involved is
We can use Le Chatelier’s principle to show how to get the
best yield of ammonia. At high temperatures, when the
reaction is faster, the position of equilibrium is to the left
because the reaction is exothermic (ΔH is negative).

What happens if we remove ammonia by condensing it to a
liquid?
We can do this because ammonia has a much higher
boiling point than hydrogen and nitrogen.
the position of equilibrium shifts to the right to replace the
ammonia that has been removed
More ammonia is formed from hydrogen and nitrogen to
keep the value of Kp constant

pH values
We can use universal indicator or a pH meter to deduce
whether a substance is acidic, alkaline or neutral.
• acids have pH values below pH 7
• alkalis have pH values above pH 7
• a neutral solution has a pH value of exactly 7
Neutralisation
Reaction of acid with base(alkali) to form salt and water
is called a neutralisation reaction.

Some simple definitions of acids and base
Acid: a substance that neutralises a base.
A salt and water are formed.
A better definition of an acid is a substance that releases
hydrogen ions when it dissolves in water.
For example:
HCl(g) + aq → H+(aq) + Cl−(aq)

Base:
A base is a compound that contains oxide or hydroxide ions
and reacts with an acid to form a salt and water.
Alkalis are bases which are soluble in water.
A base that is soluble in water is called an alkali.
For example: NaOH(s) + aq → Na+(aq) + OH−(aq)

HOME WORK PAGE 192

Bronsted – Lowry Acid-Base Theory
An acid any chemical species that is able to lose, or "donate" a
hydrogen ion (proton),
a base is a species with the ability to gain or "accept" a
hydrogen ion (proton).
NH3(g) + H2O(l) NH4
+
(aq) + OH-
(aq)
base 1 acid 1 base 2
acid 2
conjugate acid-base pair
conjugate acid-base pair

Bronsted – Lowry Acid-Base Theory
HOME WORK

Strength of Acids and Bases
Acids and bases when dissolved in water dissociates into
electrically charged ions.
Strong acids and bases are 100% ionized
Whereas weak acids and bases ionized partially.

Distinguishing a weak acid from a strong acid
We can distinguish between a strong and weak acid by their pH
values, electrical conductivity and reactivity.
When making these comparisons, we must use dilute solutions
of strong and weak acids of the same concentration.
pH values
Dilute solutions of a strong acids have lower pH values than
those of weak acids of the same concentration.
This is because the concentration of hydrogen ions is greater in
strong acids

Electrical conductivity
Dilute solutions of a strong acids have greater electrical
conductivity than those of weak acids of the same concentration.
This is because the concentration of hydrogen ions (and other
ions) is greater in strong acids.
We can determine electrical conductivity by dipping a
conductivity electrode into a solution of the acid:
a conductivity meter connected to the electrode shows the
conductivity.
Reaction with reactive metals
Acid react with some metals to produce salt and hydrogen gas.

What is titration?
A process in which a solution of known concentration is used
to determine the concentration of an unknown solution is
called titration.

Introducing indicators
An acid–base indicator is a dye or mixture of dyes that changes
colour over a specific pH range (Figures 8.15 and 8.16). In
simple terms, many indicators can be considered as weak acids
in which the acid (HIn) and its conjugate base (In−) have
different colours.

Introducing indicators
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Chemical Equilibrium-8.pptx

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