EQUILIBRIUM

Topic 9: HSC course

This topic involves both qualitative and quantitative study of reaction equilibria. The student should understand the concept of dynamic equilibrium and be able to apply it to a wide range of situations.

1. Concept of dynamic equilibrium

Systems involving simultaneous reversible reactions will come to a state of dynamic equilibrium in which:

(a) no further macro changes take place

(b) there is continual interchange between reactants and products so that: rate of forward reaction = rate of reverse reaction

the equilibrium state for a reaction can be approached from either direction

chemical systems vary in the time it takes them to reach equilibrium

an ionic compound establishes equilibrium with its ions in solution.

2. Disturbing an equilibrium

if an equilibrium is upset, it will automatically re-establish itself

concentration, volume, pressure and temperature changes can affect the position of an equilibrium

catalysts alter the time taken to reach equilibrium, but do not alter the equilibrium position

3. Equilibrium constant

is expressed as a function of the molar concentrations of the reactants and products at equilibrium

(a) write the K expression given the relevant reactants and products

(b) write the Ka expression for a given acid

ONLY calculations determining the value of K given the equilibrium concentrations or a combination of initial and equilibrium concentrations.

Note: i) the units of K are not required;

ii) calculations determining the equilibrium concentrations given a

K value are NOT required.

the equilibrium constant remains the same (at constant temperature) even if the equilibrium concentrations of reactants or products are varied

equilibrium constants may increase or decrease with a change in temperature

the equilibrium constant is a useful basis for predictions about equilibrium

(a) this has practical implications in getting the maximum yield from reactions

(b) the Ka value can be interpreted in terms of acid strength.

4. Mandatory experiences

a qualitative study of a system in which equilibrium is involved, establishing the presence of all the reactants and products of a reaction

measurement of an equilibrium constant.

5. Suggested experiences

investigation of the dependence of an equilibrium constant on temperature

qualitative experiments in which equilibrium is approached from either direction

investigation of the effect of a catalyst on the time taken to reach equilibrium.

EQUILIBRIUM

Many chemical reactions do not go to completion but are reversible.

Reversibility is shown by the double arrow: .

Equilibrium occurs when the rate of forward reaction equals the rate of reverse reaction. Concentrations all chemical species are then constant.

Factors which affect equilibrium position

• Temperature
• Concentration
• Pressure
• Volume



• Only Temperature affects K.

Equilibrium only occurs in closed systems.

Phase changes that reach equilibrium:

E.g. 1: CO2 gas in a soft drink bottle.

CO2(aq) CO2(g)

E.g. 2: Iodine

I2(s) I2(g)

Left = reactants

Right = products

Characteristics of Equilibrium systems

Constant macroscopic properties indicate attainment of equilibrium. Macroscopic properties include measurable quantities such as:

• colour
• gas pressure
• gas volume
• electrical conductivity.

Chemical Equilibria are dynamic. At equilibrium the rate of forward reaction equals the rate of reverse reaction.

Changing the position of equilibrium

A variety of factors can alter the position for various reactions:

• Concentration: When the concentration of a reactant or product is increased, the number of particle collisions will increase. According to Le Chatelier's principle the reaction will try to compensate by using up some of these added products or reactants, and therefore reestablish equilibrium.

• Gas pressure: For gasses the gas pressure is a measure of the concentration.



• Surface area: Changing the surface area of one or more of the solid components usually means that the equilibrium is achieved at a faster rate, but the position of the equilibrium is unchanged.



• Temperature: If heat is a reactant, the equilibrium is endothermic. If heat is a product, the equilibrium is exothermic.
• Generally, increasing the temperature shifts an endothermic equilibrium to the right (product side).
• e.g. A + B + heat -> C.
• Increasing the temperature shifts an exothermic equilibrium to the left (reactants side).
• e.g. X + Z -> B + heat
• See page 306.



• Effect of catalysts: A catalyst increases the rate of forward reaction to the same extent as the reverse reaction, so there is no change in the position of equilibrium when a catalyst is added.

Le Chatelier's principle

Le Chatelier's principle states:

If a system is at equilibrium and a change is made which upsets the equilibrium, then the system alters in such a way as to counteract the change and a new equilibrium is established.

Effects of Changes in Pressure

Consider the equation:

2NO_2(g) N₂O_2(g)

brown colourless

An increase in pressure would cause a shift to the right. The shift to the right reduces the number of molecules in the system thereby counteracting the increase in pressure. The colour would become lighter.

QUESTION: Predict what changes would occur for the equilibrium:

N_2(g) + 3H_2(g) 2NH_3(g)

a) Increase pressure.

b) Increase volume.

c) Add more NH3.

d) Add more H2.

a) Shift to the right. Less molecules

b) Shift to the left. More space for more molecules.

c) Shift to the left. Make more reactants

d) Shift to the right.

Representing Equilibrium Changes Graphically

See graph on page 318 or in workbook.

Shifts in Equilibrium systems

The maintenance of blood acidity via equilibrium is known as buffering. Buffers are important in the body in maintaining states of balance.

Ionic Equilibria and the Common ion effect

When NaCl dissolved in water it will continue to dissolve until the solution is saturated. The saturated solutions represents an equilibrium system in which the rate of dissolving of the solid crystal (dissociation) is exactly balanced by the rate of reverse reaction in which the ions combine to form the solid crystal (association).

Dissociation ->

NaCl_(s) ^+heat -> Na⁺_(aq) + Cl^-_(aq)

<- Association

If Na+ or Cl- ions are added this increases the rate of reverse (association) reaction. When extra ions are added which are already present in the solution, this causes the equilibrium to shift to try and use up the added ions. This is called the common ion effect.

EQUILIBRIUM CONSTANT

aA + bB -> yY + zZ

K = [Y]^y [Z]^z

[A]^a [B]^b

square brackets = concentration per mol L^-1

Value of k and position of Equilibrium

• If K>1, the reaction lies to the right, (products).
• If K<1, the reaction lies to the left, (reactants).

K (Equilibrium Constant)

• K is only constant so long as the temperature remains constant.
• K is unaffected by changes in concentrations of reactants and products (stays in same proportion).
• Solids are not included in the value of K as their concentration does not change.
• Liquids (like water) are also not included in equilibrium calculations
• K tells us the extent of the reaction (how far), but not about the rate of the reaction (how fast).

Calculation using the Equilibrium constant

• If the chemical equation is written the opposite way round then K is the reciprocal.

E.g.

• aA + bB -> yY + zZ K = 40
• yY + zZ -> aA + bB K = 1/40 = 0.025



• K is not affected by changes in concentration, volume or pressure.

Method for Calculating K

• Write a balanced equation for the equilibrium
• Write a mathematical expression for K
• Omit any species present as solids
• Calculate and list concentrations (mol L^-1) of species in equilibrium.
• Substitute concentration values into expression for K.
• Work out units.

Prac: Measurement of K

CH₃COOH_(l) + H₂O H3O⁺_(aq) + CH3COO^-_(aq)

¹pH = 3, therefore [H+] = 0.001 M

K_a = [ H3O⁺_(aq)] [CH3COO^-]

[CH₃COOH_(l)]

Note: H2O is left out of expression because there is so much the actual concentration does not really change.

K_a = [0.001] [0.001]

[0.999]

= 1 x 10^-6 mol L^-1

K_a < 1: Lies to the left (reactants) only weakly ionised.

pH:

pH = -log₁₀ [H₃O⁺]

pH = 3

3 = -log₁₀ [H₃O⁺]

-3 = log₁₀ [H₃O⁺]

10^-3 = [H₃O⁺]

[H₃O⁺] = 0.001 mol L^-1

Acid Equilibrium constant Ka

Weak acids partially ionise and are in equilibrium with their ions.

E.g.:

HA + H₂0 H3O⁺ + A^-

Ka = [H3O] [A-]

[HA]

H2O is left out as it is constant.

E.g.: 2

CH₃COOH_(l) + H₂O H3O⁺_(aq) + CH3COO^-

K_a = [ H3O⁺_(aq)] [CH3COO^-]

[CH₃COOH_(l)]

Ka values are useful in determining the extent of the ionisation.

The smaller the value of Ka the weaker the acid (or less ionised).

Ka for CH3OOH = 1.8 x 10^-5.

Since Ka << than 1 only a very small percentage of the molecules are ionised.

Ka for H2SO3 = 1.7 x 10^-2.

H2SO3 is ionised to a greater extent.

Finding degree of ionisation

concentration reactant

concentration initial product

e.g.:1

0.100 M acid

0.0043 mol L^-1 H+ ions in final product

0.0043 M

0.1 M

= 4.3 %

K involving solids

Concentration of solids is constant.

e.g.

CaCO3(s) CaO(s) + CO2

K = [CaO] [CO2]

[CaCO3]

K = [CO2]

ACIDS IN EQUILIBRIUM

Changes in pH

0 Acid

7 Neutral

14 Base

What effect would an increase the pH have on a system in equilibrium?

E.g. H2SO3 + H2O H3O⁺ + HSO3^-

Lowering the pH is the same as making the system more acidic, so lowering the pH is like increasing [H3O+].

This would shift the equilibrium to the left to counteract the change.

Increasing the pH is the same as making the system more alkaline so increasing pH is like decreasing the [H3O+] (or adding OH- which would react with the H3O+). This would shift the equilibrium to the right to counteract the change.

Ionic Solids in equilibrium with their Hydrated ions

CaCl2(s) Ca2+(aq) + 2Cl-

Ksp = [Ca2+] [Cl-]

Ksp: solubility product constant.

Ksp tells us the position of the equilibrium, i.e. how far to the left or right the equilibrium lies relative to other ionic solids (provided units same).

Ammonium Synthesis

Hydrogen: Water

Nitrogen: air

Fe3O4 catalyst plus other oxides

High pressure and temperature

N2 + 3H2 2NH3 _+energy H = -92 kJ

As the temperature is increases the speed increases but the yield decreases.

High pressure increases yield: shift right.

(a) Acetic acid, in solution with its ions.

(b) CH3COOH(aq) + H2O(1) CH3COO

(d) 1. Make a standard solution of acetic acid from glacial acetic acid.

2. Measure the pH of the solution using a pH meter.

3. Calculate the H3O + concentration, which is equal to the CH3COO - concentration.

4. Subtract the H3O + concentration from the original CH3COOH concentration to give the CH3COOH concentration at equilibrium.

a qualitative study of a system in which equilibrium is involved, establishing the presence of all the reactants and products of a reaction.

• Used Copper Nitrate (Cu(NO3)2) and Sodium Iodide (NaI)
• Cu²⁺ + 4I^- --> CuI₂ + I₂
• Test for Cu2+ ions: ammonia (NH3) solution turns deep blue.
• Test for I- ions: Lead Nitrate (Pb(NO3)2) gives yellow precipitate in the presence of I- ions. PbI₂.
• Test for I2: starch turns black/blue

Part 2

2CrO₄^2- + H⁺ -> Cr₂O₇^2- + OH^-

yellow orange

Part 3

The reversible reaction

Fe³⁺ + SCN^- --> FeSCN²⁺ + heat

Dark red

measurement of an equilibrium constant.

Acid dissociation constant

Using a weak acid eithanoic acid.

• Pour 1.0 M ethanoic acid into test tube and meaure pH using pH meter.
• pH = 3 therefore H+ and CH3COO- concentrations = 0.001
• 1 - 0.001 = CH3COOH concentration = 0.999
• Therefore Ka = 0.001²/0.999 = 1 x 10^-6.