Exothermic: chemical reaction in which heat energy is given out.
Endothermic: chemical reaction in which heat energy is taken in.
(So, in an exothermic reaction the heat exits from the chemicals so temperature rises)
Calorimetry allows for the measurement of the amount of energy transferred in a chemical reaction to be calculated.
EXPERIMENT1: Displacement, dissolving and neutralisation reactions
Example: magnesium displacing copper from copper(II) sulfate
Method:
| Initial temp. of solution (oC) | Maximium temp. of solution (oC) | Temperature rise (oC) |
|---|---|---|
| 24.2 | 56.7 | 32.5 |
Note: mass of 50 cm3 of solution is 50 g
The cup used is polystyrene because:
polystyrene is an insulator which reduces heats loss
EXPERIMENT2: Combustion reactions
To measure the amount of energy produced when a fuel is burnt, the fuel is burnt and the flame is used to heat up some water in a copper container
Example: ethanol is burnt in a small spirit burner
Method:
| Mass of water (g) | Initial temp of water (oC) | Maximum temp of water (oC) | Temperature rise (oC) | Initial mass of spirit burner + ethanol (g) | Final mass of spirit burner + ethanol (g) | Mass of ethanol burnt (g) |
|---|---|---|---|---|---|---|
| 100 | 24.2 | 54.2 | 30.0 | 34.46 | 33.68 | 0.78 |
The amount of energy produced per gram of ethanol burnt can also be calculated:
Calorimetry allows for the measurement of the amount of energy transferred in a chemical reaction to be calculated.
EXPERIMENT1: Displacement, dissolving and neutralisation reactions
Example: magnesium displacing copper from copper(II) sulfate
Method:
| Initial temp. of solution (oC) | Maximium temp. of solution (oC) | Temperature rise (oC) |
|---|---|---|
| 24.2 | 56.7 | 32.5 |
Note: mass of 50 cm3 of solution is 50 g
EXPERIMENT2: Combustion reactions
To measure the amount of energy produced when a fuel is burnt, the fuel is burnt and the flame is used to heat up some water in a copper container
Example: ethanol is burnt in a small spirit burner
Method:
| Mass of water (g) | Initial temp of water (oC) | Maximum temp of water (oC) | Temperature rise (oC) | Initial mass of spirit burner + ethanol (g) | Final mass of spirit burner + ethanol (g) | Mass of ethanol burnt (g) |
|---|---|---|---|---|---|---|
| 100 | 24.2 | 54.2 | 30.0 | 34.46 | 33.68 | 0.78 |
The amount of energy produced per gram of ethanol burnt can also be calculated:
The symbol ΔH is used to represent the change in heat (or enthalpy change) of a reaction.
ΔH is measured in kJ/mol (kilojoules per mole).
The change in heat (enthalpy change) can be represented on an energy level diagram. ΔH must also labelled.
In an exothermic reaction, the reactants have more energy than the products.
Energy is given out in the form of heat which warms the surroundings.
ΔH is given a negative sign, because the reactants are losing energy as heat, e.g ΔH = -211 kJ/mol.
In an endothermic reaction, the reactants have less energy than the products.
Energy is taken in which cools the surroundings.
ΔH is given a positive sign, because the reactants are gaining energy, e.g ΔH = +211 kJ/mol.
During chemical reactions, the bonds in the reactants must be broken, and new ones formed to make the products.
Breaking bonds need energy and therefore is described as endothermic.
Energy is released when new bonds are made and therefore is described as exothermic.
If bonds are both broken and made during chemical reactions, why can a reaction overall be describe as either exothermic or endothermic?
Example: hydrogen reacts with oxygen producing water. Overall energy is released and therefore the reaction is exothermic.
The reaction is exothermic because the energy needed to break the bonds is less than the energy released in making new bonds.
If a reaction is endothermic then the energy needed to break the bonds is more than the energy released in making new bonds.
Each type of chemical bond has a particular bond energy. The bond energy can vary slightly depending what compound the bond is in, therefore average bond energies are used to calculate the change in heat (enthalpy change, ΔH) of a reaction.
Example: dehydration of ethanol
Note: bond energy tables will always be given in the exam, e.g:
| Bond | Average bond energy in kJ/mol |
|---|---|
| H-C | 412 |
| C-C | 348 |
| O-H | 463 |
| C-O | 360 |
| C=C | 612 |
So the enthalpy change in this example can be calculated as follows:
| Breaking bonds | Making bonds | ||
|---|---|---|---|
| Bonds | Energy (kJ/mol) | Bonds | Energy (kJ/mol) |
| H-C x 5 | (412 x 5) = 2060 | C-H x 4 | (412 x 4) = 1648 |
| C-C | 348 | C=C | 612 |
| C-O | 360 | O-H x 2 | (463 x 2) = 926 |
| O-H | 463 | ||
| Energy needed to break all the bonds | 3231 | Energy released to make all the new bonds | 3186 |
Enthalpy change, ΔH = Energy needed to break all the bonds - Energy released to make all the new bonds ΔH = 3231 – 3186 = +45 kJ/mol (ΔH is positive so the reaction is endothermic) |
|||
Calorimetry allows for the measurement of the amount of energy transferred in a chemical reaction to be calculated.
EXPERIMENT1: Displacement, dissolving and neutralisation reactions
Example: magnesium displacing copper from copper(II) sulfate
Method:
| Initial temp. of solution (oC) | Maximium temp. of solution (oC) | Temperature rise (oC) |
|---|---|---|
| 24.2 | 56.7 | 32.5 |
Note: mass of 50 cm3 of solution is 50 g
EXPERIMENT2: Combustion reactions
To measure the amount of energy produced when a fuel is burnt, the fuel is burnt and the flame is used to heat up some water in a copper container
Example: ethanol is burnt in a small spirit burner
Method:
| Mass of water (g) | Initial temp of water (oC) | Maximum temp of water (oC) | Temperature rise (oC) | Initial mass of spirit burner + ethanol (g) | Final mass of spirit burner + ethanol (g) | Mass of ethanol burnt (g) |
|---|---|---|---|---|---|---|
| 100 | 24.2 | 54.2 | 30.0 | 34.46 | 33.68 | 0.78 |
The amount of energy produced per gram of ethanol burnt can also be calculated:
The rate of a chemical reaction can be measured either by how quickly reactants are used up or how quickly the products are formed.
The rate of reaction can be calculated using the following equation:
The units for rate of reaction will usually be grams per min (g/min)
An investigation of the reaction between marble chips and hydrochloric acid:
Marble chips, calcium carbonate (CaCO3) react with hydrochloric acid (HCl) to produce carbon dioxide gas. Calcium chloride solution is also formed.
Using the apparatus shown the change in mass of carbon dioxide can be measure with time.
As the marble chips react with the acid, carbon dioxide is given off.
The purpose of the cotton wool is to allow carbon dioxide to escape, but to stop any acid from spraying out.
The mass of carbon dioxide lost is measured at intervals, and a graph is plotted:
Experiment to investigate the effects of changes in surface area of solid on the rate of a reaction:
The experiment is repeated using the same mass of chips, but this time the chips are larger, i.e. have a smaller surface area.
Since the surface area is smaller, the rate of reaction is less.
Both sets of results are plotted on the same graph.
If instead the chips were smashed into powder (and again same mass of chips used) the surface area would be much larger and so the rate of reaction higher (steeper line on graph).
Experiment to investigate the effects of changes in concentration of solutions on the rate of a reaction:
The experiment is again repeated using the exact same quantities of everything but this time with half the concentration of acid. The marble chips must however be in excess. The reaction with the half the concentration of acid happens slower and produces half the amount of carbon dioxide.
Experiment to investigate the effects of changes in temperature on the rate of a reaction:
The experiment is once again repeated using the exact same quantities of everything but this time at a higher temperature. The reaction with the higher temperature happens faster.
Experiment to investigate the effects of the use of a catalyst on the rate of a reaction:
Hydrogen peroxide naturally decomposes slowly producing water and oxygen gas.
Manganese (IV) oxide can be used as a catalyst to speed up the rate of reaction.
The rate of reaction can be measured by measuring the volume of oxygen produced at regular intervals using a gas syringe.
Both sets of results are plotted on the same graph.
Experiment to investigate the reaction between varying concentrations of sodium thiosulfate and hydrochloric acid
Sodium thiosulfate (Na2S2O3) and hydrochloric acid (HCl) are both colourless solutions. They react to form a yellow precipitate of sulfur.
sodium thiosulfate + hydrochloric acid → sodium chloride + sulfur dioxide + sulfur + water
Na2S2O3(aq) + 2HCl(aq) → 2NaCl(aq) + SO2(g) + S(s) + H2O(l)
To investigate the effects of changes in concentration of sodium thiosulfate on the rate of a reaction, the conical flask is placed above a cross. The reaction mixture is observed from directly above and the time for a cross to disappear is measured. The cross disappears because a precipitate of sulfur is formed.
In order to change the concentration of sodium thiosulfate, the volumes of sodium thiosulfate and water are varied (see results table). However the total volume of solution must always be kept the same as to ensure that the depth of the solution remains constant.
In this reaction, sulfur dioxide gas (SO2), which is poisonous is produced therefore the experiment must be carried out in a well ventilated room.
The results are recorded in the table below and then plotted onto a graph.
| Volume of Na2S2O3(aq) (cm3) | Volume of water (cm3) | Concentration of Na2S2O3(aq) (mol/dm3) | Time taken for cross to disappear (s) | Rate of reaction (s-1) (1/time) |
|---|---|---|---|---|
| 50 | 0 | 0.10 | 45 | 0.0222 |
| 40 | 10 | 0.08 | 60 | 0.0167 |
| 30 | 20 | 0.06 | 80 | 0.0125 |
| 20 | 30 | 0.04 | 13 | 0.0769 |
| 10 | 40 | 0.02 | 255 | 0.0039 |
The graph shows that the rate of reaction is directly proportional to the concentration.
The experiment can also be repeated to show how temperature affects the rate of reaction.
In this experiment the concentration of sodium thiosulfate is kept constant but heated to range of different temperatures.
As a rough approximation, the rate of reaction doubles for every 10oC temperature rise.
Increasing the surface area of a solid increases the rate of a reaction.
Increasing the concentration of a solution increases the rate of a reaction.
Increasing the pressure of a gas increases the rate of a reaction.
Increasing the temperature increases the rate of a reaction.
Using a catalyst increases the rate of a reaction.
Increasing the surface area of a solid:
Increasing the concentration of a solution or pressure of a gas:
Increasing the temperature:
A catalyst is a substance that increases the rate of a reaction, but is chemically unchanged at the end of the reaction.
Catalyst: A substance that speeds up a chemical reaction while remaining chemically unchanged at the end of
the reaction.
A catalyst is not used up in a reaction.
A catalyst speeds up a reaction by providing an alternative pathway with lower activation energy.
Below is a diagram showing the reaction profile for the reaction of hydrogen with oxygen, which is EXOTHERMIC:
The activation energy is the minimum amount of energy required to start the reaction.
For an exothermic reaction, the products have less energy than the reactants. The difference between these energy levels is ΔH.
For an exothermic reaction, more energy is released when bonds are formed than taken in when bonds are broken.
Below is a diagram showing the reaction profile for the thermal decomposition of calcium carbonate, which is ENDOTHERMIC:
The activation energy is the minimum amount of energy required to start the reaction.
For an endothermic reaction, the products have more energy than the reactants. The difference between these energy levels is ΔH.
For an endothermic reaction, more energy is taken in to break bonds than is released when new bonds are formed.
The rate of a chemical reaction can be measured either by how quickly reactants are used up or how quickly the products are formed.
The rate of reaction can be calculated using the following equation:
The units for rate of reaction will usually be grams per min (g/min)
An investigation of the reaction between marble chips and hydrochloric acid:
Marble chips, calcium carbonate (CaCO3) react with hydrochloric acid (HCl) to produce carbon dioxide gas. Calcium chloride solution is also formed.
Using the apparatus shown the change in mass of carbon dioxide can be measure with time.
As the marble chips react with the acid, carbon dioxide is given off.
The purpose of the cotton wool is to allow carbon dioxide to escape, but to stop any acid from spraying out.
The mass of carbon dioxide lost is measured at intervals, and a graph is plotted:
Experiment to investigate the effects of changes in surface area of solid on the rate of a reaction:
The experiment is repeated using the exact same quantities of everything but using larger chips. For a given quantity, if the chips are larger then the surface area is lesson. So reaction with the larger chips happens more slowly.
Both sets of results are plotted on the same graph.
Experiment to investigate the effects of changes in concentration of solutions on the rate of a reaction:
The experiment is again repeated using the exact same quantities of everything but this time with half the concentration of acid. The marble chips must however be in excess. The reaction with the half the concentration of acid happens slower and produces half the amount of carbon dioxide.
Oxygen (O2) is made in the lab from hydrogen peroxide (H2O2) using manganese(IV) oxide (MnO2) as a catalyst.
Different catalysts could be used to investigate which is the most effective in decomposing hydrogen peroxide. Examples of other substances which could be tested are:
Only some of these are effective catalysts when used with hydrogen peroxide. If a substance is not a catalyst, there will be no bubbles of oxygen produced. For other substances, such as liver which is a very effective catalyst in the decomposition of hydrogen peroxide, bubbles of oxygen will be produced quickly.
Some reactions are reversible. This is indicated by the symbol ⇌
For example:
Dehydration of copper(II) sulfate
Heating ammonium chloride
On heating, white solid ammonium chloride decomposes forming ammonia and hydrogen chloride gas. On cooling, ammonia and hydrogen chloride react to form a white solid of ammonium chloride:
A dynamic equilibrium can only be achieved in a sealed container so substances cannot escape.
Features of a reaction mixture that is in dynamic equilibrium:
A catalyst is a substance which increases the rate of reaction without being chemically changed at the end of the reaction.
A reversible reaction is one where the forward reaction and the backward reaction happen simultaneously. For example:
3H₂ + N₂ ⇋ 2NH₃
In such a reaction a catalyst speeds up both the forward and the backward reactions. Hence, although the system will reach dynamic equilibrium more quickly, the addition of a catalyst will not affect the position of equilibrium.
In a reversible reaction the position of the equilibrium (the relative amounts of reactants and products) is dependent on the temperature and pressure of the reactants.
If the conditions of an equilibrium reaction are changed, the reaction moves to counteract that change.
Therefore by altering the temperature or pressure the position of the equilibrium will change to give more or less products.
Adding a catalyst does not affect the position of the equilibrium.
If a change in conditions moves equilibrium to the right, the yield of the substances on the right is increased.
Changing the temperature:
All reactions are exothermic in one direction and endothermic in the other way.
For this reaction the enthalpy change, ΔH is negative therefore the forward reaction is exothermic:
CO(g) + 2H2(g) ⇌ CH3OH(g) ΔH = –91 kJ/mol
If temperature is decreased the position of the equilibrium will shift to the right because it is an exothermic reaction.
For this reaction the enthalpy change, ΔH is positive therefore the forward reaction is endothermic:
CH4(g) + H2O(g) ⇌ CO(g) + 3H2(g) ΔH = +210 kJ mol–1
If temperature is increased the position of the equilibrium will shift to the right because it is an endothermic reaction.
Key point: an increase (or decrease) in temperature shifts the position of equilibrium in the direction of the endothermic (or exothermic) reaction
Changing the pressure:
Reactions may have more molecules of gas on one side than on the other.
For this reaction there are 2 molecules on the left and 4 molecules on the right:
CH4(g) + H2O(g) ⇌ CO(g) + 3H2(g) ΔH = +210 kJ mol–1
If the pressure is increased the position of the equilibrium will shift to the left because there are fewer molecules on the left-hand side.
For this reaction there are 3 molecules on the left and 1 molecule on the right
CO(g) + 2H2(g) ⇌ CH3OH(g) ΔH = –91 kJ/mol
If the pressure is decreased the position of the equilibrium will shift to the left because there are more molecules on the left-hand side.
Key point: an increase (or decrease) in pressure shifts the position of equilibrium in the direction that produces fewer (or more) moles of gas
