Chemical Equilibrium & Rate of Reaction Belfast Lecture

Chemical Equilibrium & Rate of Reaction Belfast Lecture
© DGMcC

To understand how reaction conditions affect reactions you
should understand: CHEMICAL EQUILIBRIUM - reversible processes and equilibrium - at equilibrium the rates of the forward and reverse reactions are equal - equilibrium constant - qualitative effect of changes in reactant and product concentration on the position of equilibrium - qualitative effects of changing the temperature and the ……...total pressure to a reaction vessel (including those used ……...on an industrial scale) on the position of equilibrium; © DGMcC

Irreversible reactions
Most Chemical reactions are considered irreversible in that products are not readily changed back into reactants. Mg + 2HCl  MgCl2 + H2 When magnesium reacts with acid it is not easy to unreact it and get back the magnesium. Wood reacting with oxygen When wood burns it is pretty difficult to un-burn it back into wood again! © DGMcC

Reversible Reactions Cotton Wool Cotton Wool Soaked in Conc. Soaked in
Hydrochloric Acid Cotton Wool Soaked in Conc. Ammonia HCl(g) NH3(g) Ring of “White Smoke” NH3(g) + HCl(g)  NH4Cl(s) © DGMcC

Reversible Reactions Moist pH Paper Solid Ammonium chloride
NH4Cl(s)  NH3(g) + HCl(g) Heat Strongly © DGMcC

Reversible Reactions NH3(g) + HCl(g)  NH4Cl(s)
NH4Cl(s)  NH3(g) + HCl(g) NH3(g) + HCl(g) NH4Cl(s) © DGMcC

Reversible Reactions CuSO4.5H2O  CuSO4 + 5H2O
© DGMcC

Reversible reactions are not uncommon
A + B ⇄ C + D On mixing A & B there will be no C & D ∴ the rate of the forward reaction is high and the rate of the reverse reaction will be zero. As the reaction proceeds there is less A & B ∴ the rate of the forward reaction decreases. As C & D are formed in increasing amounts what will happen to the rate of the reverse reaction ? It will increase. Eventually what will happen to the rates? They will become equal. © DGMcC

When this happens we say the system is in chemical equilibrium.
A + B ⇄ C + D When this happens we say the system is in chemical equilibrium. This is a dynamic equilibrium. The reaction starts then seems to stop. The forward and reverse reactions are both proceeding at the same rate. At equilibrium what happens to the amount of A? For every molecule of A reacting there is one formed ∴ the amount stays constant. This is true of the other reactants and products also. © DGMcC

Dynamic equilibrium Equilibrium – because of the unchanging amounts
Dynamic – because reaction is still occurring It is rather like the situation where a person is running on a treadmill. Neither have stopped but the person could remain in the same place for ever! The symbol  is used to mean dynamic equilibrium. The person stays in the same place! © DGMcC

Which of these is true about a dynamic equilibrium?
All the product molecules are used up. All the reactants molecules are used up. The reaction has stopped both in the forward and backward directions. The composition of the reaction mixture remains the same. © DGMcC

Which of these is a reversible process? Reacting acid with alkali.
Heating hydrated (blue) copper sulphate. Burning coal. Dissolving magnesium in acid. © DGMcC

Equilibrium Constants A + B ⇄ C + D
The concentration of each substance at equilibrium will therefore remain constant We can write an equilibrium constant for this reaction in terms of the concentrations of the reactants and products Kc = [C] x [D] [A] x [B] where [ ] = equilibrium concentration in moles /dm3 Remember it is always [products] [reactants] © DGMcC

A + 2B ⇄ 3C + 4D Kc = [C]3x [D] 4 [A] x [B] 2 Kp = P3C x P4D PA x P2B
where [ ] = equilibrium concentration in moles /dm3 Kp = P3C x P4D PA x P2B where P is the equilibrium partial pressure of a gas. © DGMcC

(The partial pressure of a gas in a mixture is the pressure that that gas would exert if it alone occupied the total volume of the mixture. The sum of the partial pressures of the gases in a mixture is equal to the total pressure of the mixture.) The equilibrium constants Kc & Kp are numerical constants for a given reaction at a given temperature. If Kc or Kp are large the amount of products will be large and reactants small, ∴ we say the equilibrium position is over to the right hand side (RHS). If Kc or Kp are small the amount of products will be small and reactants large, ∴ we say the equilibrium position is over to the left hand side (LHS). © DGMcC

State symbols are very important .
N.B. Only dissolved substances appear in Kc no gases or solids. Likewise only gases appear in Kp no solids or solutions. State symbols are very important . (s), (l), (g), (aq). e.g. 1. CaCO3(s) ⇄ CaO(s) + CO2(g) Kp = PCO2 e.g Fe 2+(aq) Ag+ (aq) ⇄ Fe 3+(aq) + Ag (s) Kc = [Fe 3+(aq)] [Fe 2+(aq)] x [Ag+ (aq)] © DGMcC

Le Chatelier The French chemist
Le Chatelier worked all this lot out!!! In a dynamic equilibrium the position of the equilibrium will shift in order to reverse any changes you introduce. © DGMcC

Le Chatelier’s Principle:
When a change is forced upon a system in chemical equilibrium the position of the equilibrium moves so that the change is opposed. A + B ⇄ C + D Predict how the above equilibrium position be changed by removing D adding C removing B Move to RHS Move to LHS Move to LHS How would the equilibrium constant be changed by removing D adding C removing B No change Only Temperature changes K © DGMcC

Changes in Pressure A (g) + B (g) ⇄ C (g)
2 gas molecules 1 gas molecule increasing P will move equilibrium to RHS A (g) + B (g) ⇄ 2C (g) 2 gas molecules gas molecules increasing P will have no effect A (g) + B (g) ⇄ 2C (g) + D (g) 2 gas molecules 3 gas molecules increasing P will move equilibrium to LHS © DGMcC

Increased pressure gives more N2O4.
Which of these is true about the effect of pressure on the reaction below? Increased pressure gives more N2O4. Increased pressure does not affect the equilibrium. Increased pressure makes N2O4 decompose. Increased pressure slows down the reaction. 2NO2(g) N2O4 (g) © DGMcC

Changes in Temperature
During a chemical reaction heat is either given out by the reaction or heat is required by the reaction. During an exothermic reaction heat is given out During an endothermic reaction heat is taken in. © DGMcC

Changes in Temperature
Exothermic Heat can be considered a product. A + B ⇄ C + D heat If temperature is increased equilibrium moves to LHS ∴ K gets smaller Endothermic Heat can be considered a reactant. heat + A + B ⇄ C + D If temperature is increased equilibrium moves to RHS ∴ K gets larger © DGMcC

2NO2(g) N2O4 (g) H=-58kJ/mol
Which of these is true about the effect of increased temperature on the reaction? gives more N2O4. does not affect the equilibrium. slows down the reactions. Achieves equilibrium more quickly. 2NO2(g) N2O4 (g) H=-58kJ/mol © DGMcC

should understand: RATE OF REACTION - qualitative effects of concentration, temperature and pressure changes (including those used on an industrial scale) and the effect of particle size (although relatively unusual on an industrial scale) - catalysts and their effect - reaction profiles, collision theory, activation energy - the effect of a change of temperature on energy distribution (Maxwell–Boltzmann distribution curves) - use of energy distribution curves to explain the effect of a change in temperature and the use of a catalyst - rate equations - order of reaction (zero, first, second) and the rate constant - determining order of reaction and rate equation from ……...data on initial rates; © DGMcC

What is the “rate" of a reaction?
The rate of a reaction is the speed of the reaction. It is not “how much” of a product is made, but instead “how quickly” a reaction takes place. © DGMcC

Reaction Rates Reaction rates vary widely, from instantaneous to very slow, from a bomb exploding to concrete setting. The factors that may influence a rate are; Concentration (solutions only) 2. Temperature (all) 3. Catalyst (all) Particle size (solids only) 5. Pressure (gases only) 6. Light intensity (photochemical reactions only) © DGMcC

THE EFFECT OF CONCENTRATION ON REACTION RATES
For many reactions involving liquids or gases, increasing the concentration of the reactants increases the rate of reaction. In a few cases, increasing the concentration of one of the reactants may have little noticeable effect of the rate. Don't assume that if you double the concentration of one of the reactants that you will double the rate of the reaction. It may happen like that, but the relationship may well be more complicated. The mathematical relationship between concentration and rate of reaction is related with the orders of reaction. © DGMcC

Changing Concentration
By increasing [ ] the number of particles in the given volume will increase ∴ the total number of collisions/sec will increase ∴ the number of effective collisions/sec will increase ∴ the reaction rate will increase. © DGMcC

THE EFFECT OF SURFACE AREA ON REACTION RATES
The more finely divided the solid is, the faster the reaction happens. A powdered solid will normally produce a faster reaction than if the same mass is present as a single lump. The powdered solid has a greater surface area than the single lump. © DGMcC

The effect of particle size
Solids with a smaller particle size (e.g. powders or small chips) react more quickly than solids with a larger particle size (e.g large chips). © DGMcC

THE EFFECT OF TEMPERATURE ON REACTION RATES
As you increase the temperature the rate of reaction increases. As a rough approximation, for many reactions happening at around room temperature, the rate of reaction doubles for every 10°C rise in temperature. © DGMcC

Changing Temperature By increasing the temperature the total kinetic energy of the mixture is increased ∴ The individual particles, on average, will have a higher kinetic energy ∴ More collisions/sec will be effective, as more will exceed the activation energy Less importantly, there are more collisions as the particles are moving faster. ∴ The rate increases. © DGMcC

THE EFFECT OF PRESSURE ON REACTION RATES
Changing the concentration of a gas is achieved by changing its pressure. Increasing the pressure on a reaction involving reacting gases increases the rate of reaction. Changing the pressure on a reaction which involves only solids or liquids has no effect on the rate. © DGMcC

COLLISION THEORY Collision theory states that...
particles must COLLIDE before a reaction can take place not all collisions lead to a reaction reactants must possess at least a minimum amount of energy - ACTIVATION ENERGY plus particles must approach each other in a certain relative way - the STERIC EFFECT © DGMcC

COLLISION THEORY According to the collision theory, to increase the rate of reaction you need... more frequent collisions increase particle speed or have more particles present more successful collisions give particles more energy or lower the activation energy © DGMcC

Energy Level Diagrams EA ΔH -ive This is an EXOTHERMIC reaction Energy
Reactants Products Reaction Coordinate © DGMcC

Energy Level Diagrams EA ΔH +ive This is an ENDOTHERMIC reaction
Products ΔH +ive Reactants Reaction Coordinate © DGMcC

Activation Energy You can think if the particles collide with less energy than the activation energy, nothing important happens. They bounce apart. Only those collisions which have energies equal to or greater than the activation energy result in a reaction. © DGMcC

should understand: RATE OF REACTION - qualitative effects of concentration, temperature and pressure changes (including those used on an industrial scale) and the effect of particle size (although relatively unusual on an industrial scale) - catalysts and their effect reaction profiles, collision theory, activation energy - the effect of a change of temperature on energy distribution (Maxwell–Boltzmann distribution curves) - use of energy distribution curves to explain the effect of a change in temperature and the use of a catalyst - rate equations - order of reaction (zero, first, second) and the rate constant - determining order of reaction and rate equation from ……...data on initial rates; © DGMcC

Catalysts A catalyst provides an alternative pathway of lower activation energy. EA cat Energy EA uncat ΔH Reactants Products © DGMcC

THE EFFECT OF CATALYSTS ON REACTION RATES
A catalyst is a substance which speeds up a reaction, but is chemically unchanged at the end of the reaction. © DGMcC

Catalysts and activation energy with a simple analogy.
Suppose you have a mountain between two valleys so that the only way for people to get from one valley to the other is over the mountain. Only the most active people will manage to get from one valley to the other. Now suppose a tunnel is cut through the mountain. Many more people will now manage to get from one valley to the other by this easier route. You could say that the tunnel route has a lower activation energy than going over the mountain. But you haven't lowered the mountain! The tunnel has provided an alternative route but hasn't lowered the original one. The original mountain is still there, and some people will still choose to climb it. © DGMcC

The Maxwell-Boltzmann Distribution
Because of the key role of activation energy in deciding whether a collision will result in a reaction, it would obviously be useful to know what sort of proportion of the particles present have high enough energies to react when they collide. © DGMcC

Maxwell-Boltzmann Distribution NUMBER OF MOLECUES WITH
A PARTICULAR ENERGY MOLECULAR ENERGY Because of the many collisions taking place between molecules, there is a spread of molecular energies and velocities. This has been demonstrated by experiment. It indicated that ... no particles have zero energy/velocity some have very low and some have very high energies/velocities most have intermediate velocities. © DGMcC

The Maxwell-Boltzmann Distribution
The area under the curve is a measure of the total number of particles present. The graph only applies to gases, but the conclusions that we can draw from it can also be applied to reactions involving liquids. © DGMcC

INCREASING TEMPERATURE NUMBER OF MOLECUES WITH
A PARTICULAR ENERGY T2 TEMPERATURE T2 > T1 MOLECULAR ENERGY Increasing the temperature alters the distribution get a shift to higher energies/velocities curve gets broader and flatter due to the greater spread of values area under the curve stays constant - it corresponds to the total number of particles © DGMcC

DECREASING TEMPERATURE NUMBER OF MOLECUES WITH
A PARTICULAR ENERGY TEMPERATURE T1 > T3 MOLECULAR ENERGY Decreasing the temperature alters the distribution get a shift to lower energies/velocities curve gets narrower and more pointed due to the smaller spread of values area under the curve stays constant - it corresponds to the total number of particles © DGMcC

Ea ACTIVATION ENERGY - Ea
MAXWELL-BOLTZMANN DISTRIBUTION OF MOLECULAR ENERGY NUMBER OF MOLECUES WITH A PARTICULAR ENERGY NUMBER OF MOLECULES WITH SUFFICIENT ENERGY TO OVERCOME THE ENERGY BARRIER Ea MOLECULAR ENERGY ACTIVATION ENERGY - Ea The Activation Energy is the minimum energy required for a reaction to take place The area under the curve beyond Ea corresponds to the number of molecules with sufficient energy to overcome the energy barrier and react. © DGMcC

The Maxwell-Boltzmann Distribution and activation energy
Remember that for a reaction to happen, particles must collide with energies equal to or greater than the activation energy for the reaction: © DGMcC

INCREASING TEMPERATURE
T2 > T1 MAXWELL-BOLTZMANN DISTRIBUTION OF MOLECULAR ENERGY T1 NUMBER OF MOLECUES WITH A PARTICULAR ENERGY T2 EXTRA MOLECULES WITH SUFFICIENT ENERGY TO OVERCOME THE ENERGY BARRIER Ea MOLECULAR ENERGY Explanation increasing the temperature gives more particles an energy greater than Ea more reactants are able to overcome the energy barrier and form products a small rise in temperature can lead to a large increase in rate © DGMcC

ADDING A CATALYST MAXWELL-BOLTZMANN DISTRIBUTION OF MOLECULAR ENERGY NUMBER OF MOLECUES WITH A PARTICULAR ENERGY NUMBER OF MOLECULES WITH SUFFICIENT ENERGY TO OVERCOME THE ENERGY BARRIER MOLECULAR ENERGY Ea The area under the curve beyond Ea corresponds to the number of molecules with sufficient energy to overcome the energy barrier and react. If a catalyst is added, the Activation Energy is lowered - Ea will move to the left. © DGMcC

ADDING A CATALYST MAXWELL-BOLTZMANN DISTRIBUTION OF MOLECULAR ENERGY NUMBER OF MOLECUES WITH A PARTICULAR ENERGY EXTRA MOLECULES WITH SUFFICIENT ENERGY TO OVERCOME THE ENERGY BARRIER MOLECULAR ENERGY Ea The area under the curve beyond Ea corresponds to the number of molecules with sufficient energy to overcome the energy barrier and react. Lowering the Activation Energy, Ea, results in a greater area under the curve after Ea showing that more molecules have energies in excess of the Activation Energy © DGMcC

Notice that the large majority of the particles don't have enough energy to react when they collide. To enable them to react we either have to change the shape of the curve, or move the activation energy further to the left. You can change the shape of the curve by changing the temperature of the reaction. You can change the position of the activation energy by adding a catalyst to the reaction. © DGMcC

should understand: RATE OF REACTION - qualitative effects of concentration, temperature and pressure changes (including those used on an industrial scale) and the effect of particle size (although relatively unusual on an industrial scale) reaction profiles, collision theory, activation energy catalysts and their effect - the effect of a change of temperature on energy distribution (Maxwell–Boltzmann distribution curves) - use of energy distribution curves to explain the effect of a change in temperature and the use of a catalyst - rate equations - order of reaction (zero, first, second) and the rate constant - determining order of reaction and rate equation from ……...data on initial rates; © DGMcC

Orders of reaction Orders of reaction are always found by doing experiments. You can't deduce anything about the order of a reaction just by looking at the equation for the reaction. So let's suppose that you have done some experiments to find out what happens to the rate of a reaction as the concentration of one of the reactants, A, changes. Some of the simple things that you might find are: © DGMcC

Orders of reaction One possibility: The rate of reaction is proportional to the concentration of A That means that if you double the concentration of A, the rate doubles as well. If you increase the concentration of A by a factor of 4, the rate goes up 4 times as well. You can express this using symbols as: Writing a formula in square brackets is a standard way of showing a concentration measured in moles per cubic decimetre (litre). © DGMcC

Orders of reaction You can also write this by getting rid of the proportionality sign and introducing a constant, k. © DGMcC

Orders of reaction Another possibility: The rate of reaction is proportional to the square of the concentration of A This means that if you doubled the concentration of A, the rate would go up 4 times (22). If you tripled the concentration of A, the rate would increase 9 times (32). In symbol terms: © DGMcC

Orders of reaction Generalising this
By doing experiments involving a reaction between A and B, you would find that the rate of the reaction was related to the concentrations of A and B in this way: This is called the rate equation for the reaction. The concentrations of A and B have to be raised to some power to show how they affect the rate of the reaction. These powers are called the orders of reaction with respect to A and B. © DGMcC

Orders of reaction If the order of reaction with respect to A is 0 (zero), this means that the concentration of A doesn't affect the rate of reaction. Mathematically, any number raised to the power of zero (x0) is equal to 1. That means that that particular term disappears from the rate equation. © DGMcC

Overall order of the reaction
The overall order of the reaction is found by adding up the individual orders. For example, if the reaction is first order with respect to both A and B (a = 1 and b = 1), R= k[A] [B] the overall order is 2. We call this an overall second order reaction. © DGMcC

Some examples Each of these examples involves a reaction between A and B, and each rate equation comes from doing some experiments to find out how the concentrations of A and B affect the rate of reaction. © DGMcC

Example 1 Rate = k[A] [B] In this case, the order of reaction with respect to both A and B is 1. The overall order of reaction is 2 - found by adding up the individual orders. Note: Where the order is 1 with respect to one of the reactants, the "1" isn't written into the equation. [A] means [A]1. © DGMcC

Example 2 Rate = k [B]2 This reaction is zero order with respect to A because the concentration of A doesn't affect the rate of the reaction. The order with respect to B is 2 - it's a second order reaction with respect to B. The reaction is also second order overall (because = 2). © DGMcC

Example 3 Rate = k[A] This reaction is first order with respect to A and zero order with respect to B, because the concentration of B doesn't affect the rate of the reaction. The reaction is first order overall (because = 1). © DGMcC

Orders of reaction What if you have some other number of reactants?
It doesn't matter how many reactants there are. The concentration of each reactant will occur in the rate equation, raised to some power. Those powers are the individual orders of reaction. The overall order of the reaction is found by adding them all up. © DGMcC

Order of Reaction For the reaction A + 2B ⇄ C + D We know rate α [A]x
We know rate α [A]x and rate α [B]y x is called the order of reaction with respect to A y is called the order of reaction with respect to B they are small whole numbers (1,2,3,) or 0. x+y is called the order of reaction. 0 = zero order, 1 = first order, 2 = second order etc. rate  [A]x[B]y or rate = k[A]x[B]y where k = rate constant © DGMcC

  N.B. ! For the equilibrium constant you use the stoichiometry of
the equation but for rate equations you can’t. Orders can only be determined experimentally. for A + 2B ⇄ 3C + 4D Kc = [C]3x [D]4 [A] x [B]2  rate = k[A][B]2  © DGMcC

should understand: RATE OF REACTION - qualitative effects of concentration, temperature and pressure changes (including those used on an industrial scale) and the effect of particle size (although relatively unusual on an industrial scale) reaction profiles, collision theory, activation energy catalysts and their effect - the effect of a change of temperature on energy distribution (Maxwell–Boltzmann distribution curves) - use of energy distribution curves to explain the effect of a change in temperature and the use of a catalyst - rate equations - order of reaction (zero, first, second) and the rate constant © DGMcC

Chemical Equilibrium & Rate of Reaction Belfast Lecture

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