1. Cover Sheet Vocab Sheet
Energy & Equilibria
Cover Sheet
Vocab Sheet

2. EE1: How does temperature change during exothermic and endothermic reactions?
A reaction which releases energy to the surroundings. The energy given out in making the bonds is greater than the energy needed to break the bonds
Reactants products + (heat)
( the test tube gets hot)
ENDOTHERMIC
A reaction which takes in energy from the surroundings. The energy needed to break the bonds is greater than the energy given out in making the bonds.
Reactants + (heat) products
(the test tube gets cold)

3. Reactions which involve temperature changes
The temperature changes for the following reactions can be measured using a polystyrene cup and lid (foam calorimeter) the results can then be used to calculate heat energy:
Neutralisation of an acid with an alkali
Dissolving a solute in a solvent
Displacement reactions, a more reactive metal displaces a less reactive metal from its salt solution

4. How can heat energy be calculated? (Doubles)
The heat energy of an experiment can by given the symbol q.
q = mcΔT
m = The mass of the substance (water or a solution) being heated in g.
c = The heat capacity of the substance being heated, which is J °C-1 g-1 for water or solutions
ΔT = The change in temperature of the experiment
E.g. If 1g of Zinc causes 20 cm3 of copper sulphate to rise by 50C what is the heat energy given off?
q = 20 x 4.18 x 5
= J (negative because the reaction is exothermic)
Practical- Temperature Changes of selected Reactions

5. How can the enthalpy change for a reaction be calculated? (Triple only)
Enthalpy (DH) is the energy change per mole of the limiting reagent of a chemical reaction
E.g. If 1g of Zinc causes 20 cm3 of copper sulphate to rise by 50C what is the heat energy given off?
DH = q / n
n is the moles of the limiting agent
n of zinc = m / Mr
= 1 / 65
DH = q / n
= / (1 / 65)
= KJ mol-1
Practical- Temperature Changes of selected Reactions

6. EE2: How can you calculate the heat energy (enthalpy) of a combustion reaction?
Fuels can be used to heat water that is contained within a copper can (calorimeter).
The temperature change of the water can be measured and the results can be used to calculate the heat energy used to heat the water (and the copper can).
The enthalpy of the reaction can be calculated by measuring the mass of the fuel at the start and end of the experiment and calculating the moles of fuel used. Then use DH = q / n to calculate the enthalpy.
Practical- Designing a combustion experiment Key Assessment
Homework- Calorimetric Questions

7. EE3: How do you represent energy changes that occur during a chemical reaction?
ΔH is used to represent the molar enthalpy change for exothermic and endothermic reactions
In chemical reactions existing (old) bonds are broken (endothermic)
New bonds are formed (exothermic)

8. Breaking chemical bonds
Most chemicals will decompose (break up) if we heat them strongly enough.
This indicates that breaking chemical bonds requires energy – is an endothermic process.
Heat taken in
Energy in chemicals
Energy needed
Energy needed to overcome the bonding between the atoms

9. Energy given out as bonds form between atoms
Making chemical bonds
Bond making will be the opposite of bond breaking
Energy will be given out in an exothermic process when bonds are formed.
Heat given out
Energy in chemicals
Energy given out
Energy given out as bonds form between atoms

10. Summary – Bond Changes
Where the energy from bond forming exceeds that needed for bond breaking the reaction is exothermic.
Where the energy for bond breaking exceeds that from bond forming the reaction is endothermic.
This can be represented in a simple energy level diagram:
Worksheet- Exo and Endo Questions
Read t/b 7.2; ans qus 1-3
Potential Energy
reactants
products
-H
Bonds break
Bond forming
Potenial Energy
+H
Bonds form
Exo
Endo
Reaction Pathway
Reaction Pathway

11. Average Bond Energies (Triple only)
Average Bond Energies can be used to calculate the Enthalpy of a Chemical Reaction
E.g. Hydrogen peroxide decomposes as shown:
Calculate energy for bond breaking (remember the sign).
Calculate the energy for bond making (remember the sign)
What is the value of H for the reaction shown O H O H
Bond
Energy (kJ)
H-O
464
O-O
146
O=O
498 O O H O H

12. Bond breaking. (endothermic) 4(O-H) + 2(O-O) =1856+292 = +2148kJ
Bond forming: (exothermic)
4(O-H) + 1(O=O)
= = -2354kJ
H = – 2354 = -206kJ
(Exothermic)
Worksheet- Average Bond Energies
Bond
Energy (kJ)
H-O
464
O-O
146
O=O
498 O H O H O O H O H

13. Activation Energy.
Most chemical reactions, including exothermic reactions, seem to need an input of energy to get the reaction started.
This fits completely with what we have already explained:
Before new bonds can be formed we need to break at least some existing chemical bonds.
This requires an energy input –known as the activation energy (Ea or Eact)
Once an exothermic reaction is underway it can provide its own activation energy (from the bond forming stage) and so sustains the reaction.

14. Activation Energy and Exothermic Reactions
reactants
H= -
Energy / kJ)
products
Progress of reaction

15. Activation Energy and Endothermic Reactions
products
H=+
Energy / kJ)
reactants
Progress of reaction

16. Copy the energy diagram and use it to help you explain why petrol stations can store petrol safely but always have notices about not smoking near the petrol pumps.
ActivationEnergy
Petrol + oxygen
Energy / kJ)
Carbon dioxide + water
Progress of reaction

17. The reaction is exothermic but requires the Activation energy to be provided before the reaction can get underway. This is necessary to break some of the bonds in the oxygen or petrol before new bonds can start forming.
ActivationEnergy
Petrol + oxygen
Energy / kJ)
Carbon dioxide + water
Progress of reaction

18. EE4: Reversible 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 cannot un-react back to magnesium
Wood reacting with oxygen
When wood burns it cannot be un-burned back into wood again!

19. However, some reactions are reversible. Heating copper sulphate.
The change from blue hydrated copper sulphate to white anhydrous copper sulphate is one of the most commonly known reversible reactions.
Heat
Water
Hydrated copper sulphate
anhydrous copper sulphate
steam
CuSO4.5H20  CuSO H2O

20. Heating ammonium chloride.
Ammonium salts are made by reacting ammonia with an acid but some of these salts will decompose back into reactants when heated.
Practical- Reversible Reactions
Read t/b 9.1; ans qus 1-3
Heat
Heat makes the solid disappear as it changes into gases.
Solid reappears as it changes back again in the cool part of the tube
Ammonium chloride
Ammonia
Hydrogen chloride
NH4Cl(s) NH3(g) + HCl(g)

21. EE5: Dynamic Equilibrium
Both the forward and backward reactions take place at the same time, leading to a state of balance with both reactants and products present in unchanging amounts (concentrations are kept constant). This is called a dynamic equilibrium. A B
These combine
These decompose

22. Equilibrium – because of the unchanging amounts.
Dynamic – because the forwards and backwards reactions are still occurring.
It is rather like the situation where a man is walking the wrong way along a moving pavement or escalator. Neither have stopped but the man could remain in the same place for ever!
Moving pavement 
Walking man
Overall – the man stays in the same place!

23. Effecting Equilibrium
Equilibrium means balance, but this balance does not have to be at the half-way point.
There are 3 factors that can influence the position of an equilibrium:
Temperature
Pressure (gas reactions)
Concentration (solution reactions)
This leads into Le Chatelier’s principle
When a reversible reaction is in equilibrium and you make a change, it will do what it can to oppose that change
Finding the conditions that gives the most product is really important in industrial chemical reactions.

24. Equilibrium and Temperature
All reactions are exothermic (give out heat) in one direction and endothermic (take in heat) in the other. E.g. nitrogen dioxide (NO2) a yellow gas joins to form dinitrogen tetroxide (N2O4) a colourless gas, exothermically
Gets hot going forward (exothermic)
Gets cold going backward (endothermic)
2NO  N2O4
The rule is: the hotter a reaction is, the more likely it is to go in the endothermic direction.
Heating will give more NO2 in the equilibrium mixture (yellow)
Cooling would give more N2O4 in the equilibrium mixture (colourless)

25. Equilibrium and Pressure
This applies to gas reactions.
Here the rule depends upon the number of gas molecules on each side of the equation
Get less gas molecules in forward direction
Get more gas molecules in backward direction
2NO2(g)  N2O4 (g)
The higher the pressure the more the reaction moves in the direction with less gas molecules.
Increasing the pressure will give more N2O4
Decreasing pressure gives more NO2 at equilibrium..

26. Equilibrium and Concentration
This applies to reactions in solution.
Worksheet- Equilibrium Questions
Read t/b 9.2; ans qus 1-3
Increasing the concentration of a substance tips the equilibrium in the direction that uses up (decreases) the concentration of the substance added.
Eg. Bismuth chloride reacts with water to give a white precipitate of bismuth oxychloride.
BiCl3(aq) + H2O (l)  BiOCl(s) + 2HCl(aq)
Adding water will produce more BiOCl solid (to use up the H2O).
Adding acid (HCl) will result in less BiOCl solid to use up the HCl.

27. Equilibrium Questions (Answers)
The reaction of nitrogen and hydrogen to form ammonia (NH3) is exothermic. How will temperature affect the composition of the equilibrium mixture?
Gets hot going forward (exothermic)
Gets cold going backward (endothermic)
3H N2  2NH3
Which direction is endothermic.
Which direction do reactions move when heated?
Will heating give more or less NH3 in the equilbrium mixture?
backward
backward
less

28. Look at the reaction of nitrogen and hydrogen to form ammonia
Get less gas molecules in forward direction
Get more gas molecules in backward direction
3H2(g) N2 (g)  2NH3 (g)
Which direction produces less gas molecules.
Which direction do reactions move when compressed?
Will high pressure give more or less NH3 in the equilbrium mixture?
forward
The side that has less gas molecules
more

29. Chlorine gas reacts with iodine chloride ( a brown liquid) converting it to iodine trichloride (a yellow solid).
ICl(l) Cl2(g)  ICl3(s)
Brown pale green yellow
What effect will adding more chlorine have upon the colour of the mixture in the U-tube?
If the U-tube is turned on its side heavy chlorine gas pours out of the tube. Which way will this tip the equilibrium?
Produce more ICl3 and so more yellow solid
Produce less ICl and so more brown liquid

30. 3H2(g) + N2 (g)  2NH3 (g) H=-92kJ/mol
Is the forward reaction exothermic or endothermic?
Will heating the mixture give an equilibrium mixture with more or less ammonia?
Are there more gas molecules of reactant or product?
Will raising the pressure give an equilibrium mixture with more or less ammonia?
exothermic
less
reactant
more

31. EE6: Production of Ammonia (The Haber Process)
Nitrogen from air, and hydrogen from natural gas or the cracking of hydrocarbons, are used in the manufacture of ammonia.
Manufacture of ammonia by the Haber process, includes the essential conditions:
i a temperature of about 450°C
ii a pressure of about 200 atmospheres
iii an iron catalyst
The cooling of the reaction mixture liquefies the ammonia produced and allows the unused hydrogen and nitrogen to be recirculated.
Ammonia that is produced by the haber process is used to make nitric acid and fertilisers.

32. Watch the video and answer the questions on the worksheet
/use this animation if you wish to show the process step by step.

33. What does the graph show about the effect of temperature on the Haber process?
Suggest why a temperature of 400oC is chosen when a lower temperature gives an equilibrium mixture with greater % conversion to ammonia.
Reduces %conversion
Hint: reaction rates?
3H2(g) + N2 (g)  2NH3 (g) H=-92kJ/mol

34. What does the graph show about the effect of pressure on the Haber process?
Suggest why a pressure of 200 atm is chosen when a higher pressure gives an equilibrium mixture with greater % conversion to ammonia.
Increases %conversion
Hint: costs?
3H2(g) + N2 (g)  2NH3 (g) H=-92kJ/mol

35. 3H2(g) + N2 (g)  2NH3 (g) H=-92kJ/mol
The Haber Compromise
3H2(g) + N2 (g)  2NH3 (g) H=-92kJ/mol
The aim of the chemical industry is not to make chemicals. It is to make money!
If low temperatures are used it takes too long to reach equilibrium. It’s better to get a 40% yield in 2 minutes than an 80% yield in 2 hours!
If very high pressure is used, the cost of the equipment used increases drastically and there are also safety issues. Better 90% conversion at 200atm than 95% conversion at 600 atm.
Unchanged reactants can always be recycled.
Iron is used as a catalyst in the process to further speed up the rate at which equilibrium is reached
Worksheet- The Haber Process
Read t/b 16.2; ans qus 1-3

36. Revision for Test