1. Grade 11: Unit 6- Energetics IB Topic 5
IB Chemistry HL1
Grade 11: Unit 6- Energetics
IB Topic 5
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2. System and Surroundings
The system is the name we give the sample or reaction vessel of interest.
The surroundings are everything else in the universe.
When a chemical change happens in an open system matter and energy can be exchanged between the system and the surroundings.
In a closed system only energy can be exchanged with the surroundings.
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3. System and Surroundings
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4. Energy
Energy is defined as the ability to do work. Energy is often converted from one form to another during physical and chemical changes.
Thermochemistry is the study of energy changes associated with chemical reactions.
Chemical energy is the energy stored in chemical bonds. It is a type of potential energy.
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5. Thermochemistry
Most reactions absorb or evolve energy usually in the form of heat but chemical reactions can also produce light, electricity and mechanical energy – used to do work.
Energy is measured in joules, J.
1000 J = 1 kJ.
Physical changes like change of state/phase also have heat energy changes.
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6. Exothermic and Endothermic Reactions
5.1.1 Define the terms exothermic reaction, endothermic reaction and standard enthalpy change of reaction (ΔH°).
Standard enthalpy change is the heat energy transferred under standard conditions—pressure kPa, temperature 298 K. Only ΔH can be measured, not H for the initial or final state of a system.
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7. Exothermic and Endothermic Reactions
5.1.3 Apply the relationship between temperature change, enthalpy change and the classification of a reaction as endothermic or exothermic.
5.1.4 Deduce, from an enthalpy level diagram, the relative stabilities of reactant and products, and the sign of the enthalpy change for the reaction.
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8. Enthalpy
Enthalpy is the total energy of a system, some of which is stored as chemical potential energy in the chemical bonds.
Enthalpy is given the symbol H.
Enthalpy is also known as the heat content of a system.
We cannot measure the enthalpy content of a system but we can measure changes in it.
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9. Enthalpy Change
In chemical reactions, bonds are broken and made, but the energy absorbed breaking bonds is almost never exactly equal to that released in making new bonds.
All reactions are accompanied by a change in the potential energy of the bonds and hence an ENTHALPY CHANGE.
There is no “absolute zero” for enthalpy so absolute enthalpies cannot be measured only the change in enthalpy that occurs during a reaction.
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10. Enthalpy Change
The enthalpy change of a reaction is given the symbol ∆H.
∆H is the difference in the enthalpy between the products and the reactants.
∆H = H(products) - H(reactants) when at constant pressure.
Enthalpy level diagrams are used to show the change in enthalpy of a system during a change.
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11. Enthalpy Level Diagram
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12. Enthalpy Level Diagram
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13. Endothermic Reaction
An endothermic reaction is one where energy is transferred from the surroundings to the system.
If energy is absorbed during a reaction then the enthalpy of the products will be higher than that of the reactants.
This means the enthalpy change will have a positive sign.
This reaction will either get cooler or heat will need to be supplied - temperature decreases.
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14. Exothermic Reaction
An exothermic reaction is one where energy is transferred from the system to the surroundings.
If energy is released during a reaction then the enthalpy of the products will be lower than that of the reactants.
This means the enthalpy change will have a negative sign.
This reaction will “feel” hotter - temperature increases.
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15. Stability
Reactions in chemistry tend towards products that are more stable.
Stability increases as energy decreases so exothermic reactions increase the stability of the substances.
This is why most chemical reactions that occur in nature are exothermic.
Endothermic reactions usually need “help” in the form of energy to allow them to occur.
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16. Stability
Remember that lower energy is more stable. We can compare this to standing on top of a high building where you have more potential energy than someone on the ground.
If you fall to the ground you lose some of that potential energy you had on the roof but you are now more stable.
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17. Exothermic and Endothermic Reactions
5.1.2 State that combustion and neutralization are exothermic processes.
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18. Combustion Combustion is the scientific word for burning.
Most hydrocarbons burn easily in excess oxygen.
When they burn they produce carbon dioxide and water.
Ex. CH4(g) + 2O2(g)  CO2(g) + 2H2O(g)
Combustion produces lots of heat which is transferred to the surroundings so it is very EXOthermic.
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19. Combustion
When 1 mole of a substance is burned the energy released is called the Enthalpy Change of Combustion, ΔHc°.
The ° sign indicates this was measured under standard conditions in a controlled environment.
Standard conditions for thermochemistry experiments are T = 298K and P = 101.3kPa (= 1 atm).
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20. Neutralization Neutralization reactions involved acids and bases.
If an acid and a base react completely the resulting solution will be pH neutral.
The products are a salt (ionic compound) + water.
Ex. NaOH(aq) + HCl(aq)  NaCl(aq) + H2O(l)
When weak acids and bases are involved such as NaHCO3 then CO2(g) may be another product.
These reactions release energy to the surroundings so they are EXOthermic.
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21. Negative ΔH
As combustion and neutralization reactions are always exothermic the enthalpy changes will always have negative values.
You can find many ΔH values in scientific literature ex. your textbook, the books in the classroom and online.
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22. Spontaneous Reactions
A spontaneous reaction is one that occurs when the reactants are mixed without the need to be heated or have some other outside influence.
Most spontaneous reactions are EXOthermic but there are some spontaneous endothermic reactions ex. Dissolving NH4Cl in water.
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23. Summary of Enthalpy Changes
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24. Enthalpy Changes
In an exothermic reaction the products are more stable than the reactants so the bonds made are stronger than the bonds broken.
In an endothermic reaction the products are less stable than the reactants so the bonds made are weaker than the bonds broken.
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25. Enthalpy Changes
Enthalpy changes are usually written alongside the chemical equation for the process with a positive or negative sign.
State symbols are VERY IMPORTANT as changes of state have their own enthalpy change values.
Enthalpy changes are usually reported per mole so the units are kJ mol-1.
If this is not the case then just kJ is used.
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26. Enthalpy Changes For example: 2NaHCO3(s)  Na2CO3(s) + H2O(l) + CO2(g)
H = kJ mol-1
The + sign indicates it’s an endothermic reaction.
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27. Standard Conditions
To compare enthalpy changes conditions must be the same.
The thermochemical standard conditions are
Temperature = 25°C = 298K (this is room temp)
Pressure = 1 atm = 101.3kPa
Solutions have a concentration of 1 mol dm-3
Standard conditions are sometimes indicated by the symbol  or °: H orΔH°
Sometimes the temperature is included too: H 298.
The values in bold are the SI units.
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28. Learning Check: Do NOW
Write the equation for the formation of chlorine oxide, Cl2O from its elements
What bonds are broken and what bonds are made in this process?
Do the processes in 2. absorb or release energy?
What is an enthalpy change?
In this reaction the bonds made are less strong than those broken. Will the enthalpy change be positive or negative?
Will this be an exothermic or endothermic reaction?
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29. Calculation of Enthalpy Changes
5.2.1 Calculate the heat energy change when the temperature of a pure substance is changed.
Students should be able to calculate the heat energy change for a substance given the mass, specific heat capacity and temperature change using q = mcΔT.
5.2.2 Design suitable experimental procedures for measuring the heat energy changes of reactions. Students should consider reactions in aqueous solution and combustion reactions.
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30. Calculation of enthalpy changes
5.2.3 Calculate the enthalpy change for a reaction using experimental data on temperature changes, quantities of reactants and mass of water.
5.2.4 Evaluate the results of experiments to determine enthalpy changes. Students should be aware of the assumptions made and errors due to heat loss.
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31. Temperature
Temperature is a measure of the average kinetic energy of the particles in a system.
Units are K or °C.
Note: there is no ° sign used with K!
Heat is a measure of the total energy in a substance.
T°C = TK
TK – 273 = T°C
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32. Specific Heat Capcity
Some substances will conduct heat and therefore change temperature more easily than others.
Ex. A metal pan on a stove will become very hot before the water inside it does.
A measure of how easily something changes temperature is called SPECIFIC HEAT CAPACITY.
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33. Specific Heat Capacity
Specific heat capacity is the amount of heat energy required to increase 1 g of a substance by 1K or 1°C.
It is used in the following equation:
q = m x c x T where q is heat energy, m is mass, c is specific heat capacity and T is change in temperature.
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34. Specific Heat Capacity
Units of specific heat capacity (c) are J g-1K-1
Units of heat energy (q) are J or kJ
Units of mass (m) are g or kg
Units of temperature (T) are K or °C
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35. Calorimetry
Calorimetry is a method used to measure the enthalpy associated with a particular change.
The temperature change of a liquid is measured inside a well insulated container called a calorimeter.
Often a styrofoam cup is used as it has a very low heat capacity and is a good insulator.
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36. Measuring Energy Changes
To measure the enthalpy change of a reaction that occurs in solution you can carry it out in a styrofoam (polystyrene) cup and monitor the temperature during the reaction.
Styrofoam is a good insulator so the amount of heat lost to the surroundings will be reduced.
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37. Measuring Energy Changes
Burning substances in a “bomb calorimeter” measures the temperature change in the water surrounding the burning item.
This system is also well insulated to try and reduce heat loss.
There may be some losses from incomplete combustion.
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38. Calorimetry
If calorimeters made of other materials are used then the heat absorbed by the calorimeter must be added to that absorbed by the liquid:
Heat absorbed = (mcT)liquid + (mcT)calorimeter
Calorimetry assumes no heat is transferred to or from the surroundings so they must be well insulated.
However this is hard to achieve and is a major source of error in high school labs.
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39. Calorimetry
Once you have calculated the energy released from a process then you can calculate how many kJ of energy were released per mole of the reactant. This is usually referred to as ΔH or molar enthalpy in a test question.
ΔH = q/n
Where n represents the number of moles of the reactant that is reacting/burning.
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40. Sample Problems Specific heat, c of liquid water = 4.18 kJ dm-3K-1
= 4.18J g-1K-1
How much heat energy is required to increase the temperature of 10 g of nickel (c = 440 J g-1K-1) from 50°C to 70°C?
The enthalpy of combustion of ethanol (C2H5OH) is 1370 kJ mol-1. How much heat is released when mol undergo complete combustion?
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41. Sample Problems
3. H2(g) + 1/2 O2(g)  H2O(l) ∆H for the reaction above is -286 kJ mol-1. What mass of oxygen must be consumed to produce 1144 kJ of energy? 4. Calculate the molar enthalpy change when excess zinc is added to 50 cm3 of a 1 mol dm-3 solution of CuSO4. The temperature increases from 20°C to 70°C when the zinc is added. Assume the solution has the same density as water = 1.00 g cm-3
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42. Bond Enthalpies 5.4.1 Define the term average bond enthalpy.
5.4.2 Explain, in terms of average bond enthalpies, why some reactions are exothermic and others are endothermic.
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43. Bond Enthalpies
All chemical reactions involve the making and breaking of bonds.
The bonds in the reactants are broken which absorbs energy so this is an endothermic process.
The bonds in the products form which releases energy so this is an exothermic process.
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44. Bond Enthalpy
Bond enthalpy is defined as the energy needed to break one mole of bonds in gaseous molecules under standard conditions.
Ex. ½ H2(g) + ½ Cl2(g)  H(g) + Cl(g)
Breaking a bond is endothermic so these values are always positive.
Bond enthalpies depend on the rest of the molecule so values are usually averages.
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45. Bond Enthalpy
The energy released when a bond is made is the same value as the bond enthalpy but with a negative sign.
A higher bond enthalpy indicates a stronger bond.
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46. Bond Enthalpy
If the bonds being broken (bonds in the reactants) are weaker than those being made (bonds in products) then the reaction will be exothermic.
If the bonds being broken are stronger than those being made then the reaction will be endothermic.
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47. Bond Enthalpies
As bond enthalpies are averages and they are only for gases they are not the most accurate way to calculate an enthalpy change but they are usually within about 10% of more accurate values and are a useful tool.
∆Hºreaction= ∑BEbonds broken- ∑BEbonds made
Where BE stands for bond enthalpy
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48. Bond Enthalpy
Bond enthalpy values are given in the data booklet. A sample is shown here. When using these values be careful to check if bonds are single or double etc.
Bond
E / kJ mol-1
H-H
436
C-H
413
C-C
347
C=C
612
O=O
498
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49. Bond Enthalpy Calculations
When water is formed from its elements, what bonds are broken and formed? What is the enthalpy change predicted by bond enthalpies?
2H2(g) + O2(g)  2H2O(l)
Bonds broken: 2(H-H), 1 (O=O)
Bonds formed: 4 (O-H)
ΔH = Σbonds broken – Σbonds formed
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50. Bond Enthalpy Calculations
= (2(436) ) – (4(464))
= ( ) – (1856)
=1370 – 1856 = -486 kJ
What does this sign tell you about the reaction?
Does this make sense?
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51. Practice Problem
Find the enthalpy change when CO2(g) is formed from its elements using bond enthalpy values in the data booklet.
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52. Hess’s Law
5.3.1 Determine the enthalpy change of a reaction that is the sum of two or three reactions with known enthalpy changes.
Students should be able to use simple enthalpy cycles and enthalpy level diagrams and to manipulate equations. Students will not be required to state Hess’s Law.
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53. Hess’s Law
Hess’s Law is a special case of the law of conservation of energy.
It states that the enthalpy change for a reaction will have the same value no matter how many steps were taken to go from reactants to products.
Another way we can use it is to say the enthalpy change for the reaction is equal to the enthalpy of the products - enthalpy of reactants.
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54. Hess’s Law
Hess’s Law enables us to calculate an enthalpy change for a reaction without carrying out the actual reaction.
We do this by measuring the enthalpy change for other related reactions.
Hess’s Law is very useful for reactions that are difficult to carry out in a lab or that do not occur.
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55. Hess’s Law
For example sodium hydrogen carbonate can be reacted with hydrochloric acid as follows
NaHCO3(s) + HCl(aq)  NaCl(aq) + CO2(g) + H2O(l) ∆H1
To give the same products we could carry out 2 other reactions:
2NaHCO3(s)  Na2CO3(s) + CO2(g) + H2O(l) ∆H2
Na2CO3(s) + 2HCl (aq)  2NaCl(aq) + CO2(g) + H2O(l) ∆H3
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56. Hess’s Law
The sodium hydrogencarbonate could be heated to form sodium carbonate which is then reacted with hydrochloric acid.
2NaHCO3(s)  Na2CO3(s) + CO2(g) + H2O(l) ∆H2
Na2CO3(s) + 2HCl (aq)  2NaCl(aq) + CO2(g) + H2O(l) ∆H3
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57. Hess’s Law
If these two equations are added together the Na2CO3 cancel out and the result is twice the overall equation.
2NaHCO3(s) + 2HCl(aq) --> 2NaCl(aq) + 2CO2(g) + 2H2O(l)
Hess’s Law says that the enthalpy change for the two stage reaction must be equal to the single stage process:
2∆H1 = ∆H2 + ∆H3
The number 2 is present in front of ΔH1 as the equations must all be balanced.
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58. Hess’s Law
This can also be shown using an enthalpy cycle:
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59. Hess’s Law
To read an energy cycle you must identify the reactants and products that you are being asked about and make sure all 3 sides are balanced.
Then you must identify the alternative route you will take to get from reactants to products.
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60. Hess’s Law
If you are following an arrow then you add all those ΔH values. You then subtract any ΔH values when you go in the opposite direction of the arrow.
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61. Hess’s Law
You can also use a series of equations and compare them to find the unknown enthalpy change.
For example:
S(s) + 1½O2(g)  SO3(g) ΔH°=-395 kJ ΔH1
SO2(g) + ½O2(g)  SO3(g) ΔH°=-98 kJ ΔH2
Calculate the standard enthalpy change, ΔH° for the reaction:
S(s) + O2  SO2(g)
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62. Hess’s Law ΔH1 has the same reactants as the reaction in question.
ΔH2 relates SO3(g) to SO2(g) but the SO2 is on the opposite side of the equation from the reaction in question. This means we reverse the chemical change and the sign of ΔH2.
The equations can now be combined and the values added to find the ΔH°:
ΔH°= = -297 kJ
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63. Hess’s Law Activity Worksheet
1. Calculate ΔH for the reaction
C2H4 (g) + H2 (g)  C2H6 (g),
from the following data.
C2H4 (g) + 3 O2 (g)  2 CO2 (g) + 2 H2O (l) ΔH = kJ/mole
C2H6 (g) + 7/2 O2 (g) 2 CO2 (g) + 3 H2O (l) ΔH = kJ/mole
H2 (g) + 1/2 O2 (g) H2O (l) ΔH = kJ/mole
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64. Solution to #1 problem
Reactions that were reversed or multiplied by a constant are shown in italics.
1. ΔH = kJ
C2H4 (g) + 3 O2 (g)  2 CO2 (g) + 2 H2O (l) ΔH = kJ
2 CO2 (g) + 3 H2O (l)  C2H6 (g) + 7/2 O2 (g) ΔH = kJ
H2 (g) + 1/2 O2 (g)  H2O (l) ΔH = kJ
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65. Hess’s Law Worksheet 2. Calculate ΔH for the reaction
4 NH3 (g) + 5 O2 (g) 4 NO (g) + 6 H2O (g),
from the following data.
N2 (g) + O2 (g) 2 NO (g) ΔH = kJ
N2 (g) + 3 H2 (g) 2 NH3 (g) ΔH = kJ
2 H2 (g) + O2 (g)  2 H2O (g) ΔH = kJ
Answer:
ΔH = kJ
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