Higher Chemistry

Energy MattersReaction RatesRates

Reaction rates are dependant on the concentration of the reactants particle size temperatureCatalysts speed up reactionsThe concentration of reactants in solution is measured in the units; mol/litre or mol litre-1 or mol l-1.

How to find the rate of a reaction?Usually found experimentally by measuring how fast a reactant is used up or how fast a product is formed.a) the rate of lass loss or increase in volume if a gas is a product. b) a change in pH if, for example, and acid is being neutralised. c) the rate at which a colour develops or disappears – this is gone using an instrument called colorimeter.

Definition of rate: defined as the change in concentration of reactants or products in unit time.

Experimentreaction between marble chips (calcium carbonate) and dilute hydrochloric acid.

CaCO3(s) + 2HCl(aq) CaCl2(aq) + H2O(l) + CO2(g)

1. by measuring the mass loss of the flask as the carbon dioxide is produced and escapes from the flask.

2. by measuring the fall in concentration of the aci as the reaction proceeds.

3. by collecting and measuring the volume of carbon dioxide produced.

Average Rate = loss of mass/time interval = ___ gs-1

Average Rate = decrease in acid concentration/time interval = ___ mol l-1 s-1

Effect of Concentration on Reaction Rate Persulphate Oxidation of Iodide Ions The Formation of Colloidal Sulphur from Thiosulphate ions.

For both of these reactions there is a straight line through the origin showing that rate is directly

proportional to concentration.

Effect of Temperature on Reaction Rate Oxidation of oxalic acid by potassium permanganate at different

temperatures. Reaction between thiosulphate ions and dilute hydrochloric acid

These reactions show a marked increase in rate as the temperature is increased.

* Roughly the rate of a chemical reaction is doubles for every 10oC rise in temperature *

Collision TheoryFor a reaction to occur the reactant molecules have to collide. This is the basis of the Collision Theory.

The collision provides energy to break the bonds in the reactant molecules and for new bonds to be formed to make the product molecules.

Everything that we know about the effect of concentration, particle size and temperature can be explained by the Collision Theory.

ConcentrationThe more concentrated the reactants the more collisions there are going to be between the reactant molecules and hence the faster the reaction. Particle SizeIn a reaction with solids the smaller the particle size the larger the surface area presented to the other reactant. This increases the chance of collision and so increases the reaction rate.TemperatureThe temperature of a substance the measure of the average kinetic energy of its molecules.At low temperature, molecules are moving slowly so have a smaller kinetic energy. At a high temperature, molecules are moving quickly so have a high kinetic energy.

If molecules are moving slowly they may collide but may not have enough energy to react. If they were moving quickly they are more likely to react. Meaning that there must be a minimum speed (kinetic energy) for the reaction to take place, this is called the Activation Energy.

If the temperature of a reaction is raised it means that more molecules will have successful collisions and will have the required Activation Energy. This is why rate increases so much with temperature.

Light as a supplier of the Activation EnergyWith some chemical reactions light can be used to increase the amount of particles with energy greater than the activation energy, they are called photochemical reactions. Photography

Light hitting a photograph film provides the activation energy for this reaction.

Ag+ + e- Ag(s)

Hydrogen and ChlorineWhen H and Cl are mixed together in the dark nothing happens. But when an ultraviolet light is shone on them there is a explosive reaction forming hydrogen chloride.

H2 + Cl2 2HClThe light provides the activation energy for the dissociation of some chlorine molecules.

Potential Energy DiagramsDuring an exothermic reaction some of its chemical potential energy in the reactant molecules is released as heat energy, meaning that the product contains less potential energy than the reactant molecules. During an endothermic reaction heat energy is absorbed from the surroundings, so the products have more potential energy than the reactant molecules.

Exothermic graphthe difference in the potential energy is called the enthalpy

change for the reaction and has the symbol ΔH. Exothermic reactions

have a negative ΔH.

Endothermic graphEndothermic reactions have a positive ΔH.

If the activation energy is low, many molecules have enough energy to over comes the barrier and the reaction will be fast.

Energy Distribution and the Activated Complex

This shows that at 20oC only a few molecules have energy greater than Ea (Activation Energy), whereas at 30oC many more molecules have greater energy than Ea.

Activation Energy and the Activated ComplexWhen two molecules collide the first thing that happens is they join to form a species called the activated complex. The activated complex is very unstable so only exists for a short period of time. The activation energy required energy for its formation because there may be bonds in the reactants to be weakened or broken. This energy required comes from the Ek of the colliding molecules and is stored as Ep in the bonds of the activated complex. The energy needed to form the activated complex is the activation energy.

*If the energy available from the collision is less than the Ea, the activated complex cannot be formed so no reaction can occur *

The Effect of Catalysts on Reaction Rate

Definition of Catalyst: A catalyst is a substance that alters the rate of a chemical reaction without itself undergoing any permanent chemical change.

Because the catalyst is not used up the amount of catalyst is usually small compared to the other reactants. Catalysts are normally used to speed up a reaction, but can also be used to slow down a reaction.

When it is used to slow down a reaction it is called a inhibitor. Inhibitors are used in rubber to increase its stability, in antifreeze to slow down rusting, to stabilise monomers for polymerisation.

A catalyst provides an alternative pathway between the reactants and products. To speed up a reaction the alternative route has a lower activation energy than the route without the catalyst so the reaction is speeded up.

example with an exothermic reaction.

* The ΔH for the reaction is the same whether the catalyst is used

or not *

Types of CatalystsDefinition of Homogeneous Catalysts: Catalysts that are in the same physical state as the reactants.Definition of Heterogeneous Catalysts: Catalysts that are in a different physical state from the reactants. Examples of Heterogeneous Catalysts;

How do heterogeneous catalysts work?

Catalyst PoisoningDefinition of a Catalyst Poison: A substance which reacts with a catalyst and prevents the catalyst from doing their job. This is done by preventing the reactant molecules being absorbed and making the catalyst ineffective.

Even with precautions in place, impurities remain in the reactants mean that catalysts must be regenerated or renewed from time to time.

EnzymesDefinition of an Enzyme: An enzyme is a catalyst that working in biological systems. They catalyse the reactions that happen in the cells of plants and animals. Examples

1. yeast contains enzymes that convert carbohydrates into alcohol. 2. amylase is an enzyme that is used to hydrolyse starch. 3. catalase is an enzyme that helps with the decomposition of hydrogen

peroxide.

Enzymes are proteins which are complicated large molecules. Its 3D shape is very important to its functions as an enzyme. If the enzyme changes in temperature or pH the shape can be denatured, so this means that they work best as certain temperatures and pHs.

AdsorptionMolecules of the reactants create bonds with the catalyst. This

weakens the bonds within the

molecules.

ReactionThe molecules

react on the surface of the

catalyst.

DesorptionThe products leave

the catalyst and the empty active

site can be used by another molecule.

One enzyme can only catalyse one reaction.

Enzymes in our bodies can be poisoned by arsenic, cyanide, and carbon monoxide. This prevents the enzyme from doing its job which can lead to death.

Chemical ReactionsChemical EquilibriumReversible ReactionsReactions where the products can form the reactants are called reversible reactions.

A + B C + DIf we start with A and B and allow them to reaction, the rate of the forward reactions (rf) will be high as the concentrations of A and B are high. The rate of the reverse reactions is zero for a start as the concentrations of C and D is zero. As the reactions proceed the concentrations of A and B decrease while the concentrations of C and D increase. This means rf falls and rr increases. This continues until the two rates become equal. At this point the concentration of A, B, C and D do not change (unless the conditions are altered) and is said to have reached equilibrium.

Definition of Dynamic equilibrium: At the molecular level the forward and backward reactions are continuing but, because their rates are equal the concentrations of the four substances remain constant.

Equilibrium is only reached in a closed system (no substances are added or removed).

Position of EquilibriumEquilibrium does not imply 50% reactants and 50% products. Sometimes equilibrium is established when the forward reaction is nearly complete – so we would say that equilibrium lies to the right. In other cases equilibrium is reached when the forward reaction is barely started – equilibrium lies to the left.

Factors Which Affect the Position of EquilibriumEquilibrium is reached when the rates of two opposing reactions become equal. The rate is affected by;a) catalystsb) concentrationc) pressure (of gases)d) temperature

CatalystsA catalyst has the effect of lowering the energy barrier between reactants and products by providing an alternative reaction path. The net effect is that a catalyst does not alter the position of equilibrium. However, a catalyst speeds up both the forward and reverse reaction, so the same equilibrium is reached more quickly.

PositionA + B (reversible reaction) C + D

Increasing the concentration of A + B will speed up the forward reactions so producing more C and D until a new equilibrium position further to the right is established. Decreasing the concentration of C or D will slow down the reverse reaction which converts C and D into A and B. This means the concentration will increase again moving the equilibrium to the right.

* The intensity of the colour indicates the position of equilibrium *

Pressure

A change of pressure can only affect equilibria in which gasses are involved. The pressure exerted by a gas I caused by the freely moving gas molecules colliding with the walls of the containing vessel. An increase in the number of molecules in the vessel will cause an increase in pressure; so long the container size is kept constant. Increasing the pressure favours whichever reaction brings about a reduction in the total number of gas molecules. Decreasing the pressure favours the reaction that increases the total number of gas molecules.

If an equilibrium system has the same number of gas molecules on both sides of the arrow, a change in pressure will have no effect on the position of equilibrium. However an increase in pressure will increase both forward and reverse reactions and so reduce the time for equilibrium to be established.

TemperatureIn a system at equilibrium, if the forward reaction is exothermic the reverse reaction must be endothermic, and vice versa.

If the temperature is raised, then the rate of both reactions increases but not equally. A rise in temperature favours the reaction that needs to have heat supplied (the endothermic reaction). A decrease in temperature has the opposite effect and favours the exothermic reaction.

Le Chatelier’s PrincipleThe effect of changes in concentration, pressure and temperature on an equilibrium can be predicted using Le Chatelier’s principle.

“If a system at equilibrium is subjected to a change, the system will adjust to oppose the effect of the change”

The Haber Process for Ammonia

N2(g) + 3H 2NH3(g) ΔH = -92kJ mol-1

CatalystIn the absence of a catalyst the nitrogen and the hydrogen hardly combine. The high temperature needed to make thee nitrogen and hydrogen combine

would forc the equilibrium to the left so little ammonia would be formed. An iron catalyst is used in the Haber Process; this allows a fast reaction rate at a lower temperature and gives a reasonable yield of ammonia.

PressureThe formation of ammonia gives a decrease in the number of molecules of gas, so a high pressure favours the ammonia production. However plants that operate at high pressure and costly to build and require expensive compressors.

TemperatureA low temperate would give a high equlibirium yield of ammonia. However a low temperature means a slow rate and a long time to come to equilibrium. A higher temperature increases the rate but gives a reduced yield of ammonia.

Raw materials for the Haber processNitrogen comes from the air. The Hydrogen comes from the syngas manufactured from natural gas and steam.

Marketability of AmmoniaThis is clearly important for any industrial product. There is a large market for ammonia because it is further converted into fertilisers, nitric acid and nylon.

Equilibrium in aqueous solutionsPure water conducts electricity slightly, due to the slight dissociation of water molecules..

H2O(l) H+(aq) + OH-

(aq)

Equilibrium lies to the left; only 1 in every 555 million water molecules dissociates.

The pH ScaleDilution of 1 mol l-1 HClWith a pH meter we can see that 1 mol l-1 HCl hada pH of 0. We took 10ml of this acid solution and made the volume up to 100ml with distilled water – this is a ten fold dilution and left us with 0.1 mol l-1 HCl with a pH of 1. This dilution was repeated several times.

* For dilutions such as this, C1V1 = C2V2*

Square brackets are used in chemistry to denote the concentration. [H+] means “the concentration of H+ ions” and is normally measured in mol l-1.

0.1 mol l-1 HCl [H+] = 10-1 mol l-1

0.01 mol l-1 HCl [H+] = 10-2 mol l-1

0.001 mol l-1 HCl [H+] = 10-3 mol l-1

Results

Dilution of 1 mol l-1 NaOHDoing a similar experiment we got these results;

Ionic Product for WaterH2O(l) H+

(aq) + OH-(aq)

When pure water dissociates one H+ ion is produced for every OH- ion, so [H+] = [OH-]

For water to have a pH 7[H+] = [OH-] = 10-7mol l-1

The ionic product of water, Kw = [H+] [OH-] = 10-7 x 10-7 mol2 l-2

= 10-14 mol2 l-2

* This is true at all pH values. *

[H+] [OH-] = 10-14 mol2 l-2 must be true at all times in aqueous solutions.

Calculating the pH of solutionsExample 1;What is the pH of a 0.01 mol l-1 solution of hydrochloric acid?

HCl(g) H+(aq) + Cl-(aq)

1 mole 1 mole .[H+] = 0.01mol l-1

= 10-2 mol l-1

So the pH = 2Example 2;What is the pH of a 0.001 mol l-1 solution of sodium hydroxide?

NaOH Na+(aq) + OH-

(aq)

1 mole 1 mole[OH-] = 0.001 mol l-1

= 10-3 mol l-1

In any aqueous solution:[H+][OH-] = 10-14 mol2 l-2

So [H+] = 10-14/[OH-]= 10-14/10-3

= 10-11 mol l-1

So the pH = 11

Strong and Weak acidsThe pH of a 0.1 mol l-1 solution of hydrochloric acid is 1. Hydrochloric acid is a strong acid and is fully dissociated into ions in aqueous solution.

HCl(g) H+(aq) + Cl-(aq)

1 mole 1 mole .So if the HCl concentration is 0.1mol l-1 then [H+] is also 0.1 mol l-1

Ie [H+] 0.1 = 10-1 mol l-1

So the pH = 1When the pH of 0.1 mol l-1 ethanoic acid is measured it is found to be 3. this indicates a lower hydrogen ion concentration. Ethanoic acid is a weak acid because it is not fully dissociated into ions in aqueous solution.

CH3COOH(aq) CH3CO|O-(aq) + H+

(aq)

Comparison of Strong and Weak AcidsEquimolar (0.1 mol l-1) solution of hydrochloric and ethanoic acids were compared in a number of experiments. The results were:

The higher concentration of [H+](aq) ions in hydrochloric acid accounts for the lower pH, the high conductivity and the faster reaction rates.

Those experiments can be used to distinguish strong and weak acids

Amount of Alkali Neutralised by Strong and Weak AcidsNeutralisation is the joining of H+ and OH- ions to form water.

H+(aq) + OH-

(aq) H2O(l)

It would be expected that weak acids with the lower concentration of H+ ions would neutralise a small amount of alkali than a strong acid, however 0.1 mol l-1 ethanoic acid neutralises exactly the same volume of sodium hydroxide as 0.1 mol l-1 hydrochloric acid.

CH3COOH(aq) CH3COO-(aq) + H+

(aq)

removed by OH-(aq) to form water.

The amound of alkali neutralised cannot be used to distinguish strong and weak acids.

Examples of Strong AcidsHydrochloric Acid – HCl H+

(aq) + Cl-(aq)

Nitric Acid – HNO3 H+(aq) + NO3

-(aq)

Sulphuric Acid – H2SO4 2H+(aq) + SO4

2-(aq)

Examples of Weak AcidsCarboxylic Acids

Carbonic Acid CO2 + H2O(l) H2SO3(aq) 2H+

(aq) + CO32-

(aq)