1. Thermochemistry
Lesson # 1:
Heat & Energy

2. Background Definitions
VIDEO
Thermodynamics – the study of energy and energy transfer.
Force – directional push or pull of an object, measured in the Newton, N.
Work – process of energy transfer to an object when the object is moved by a force.
Energy – the ability to do work, measured in the joule, J. The joule is equal to the energy used when applying a force of one Newton over a distance on one metre
(1 J = 1 N.m)
Kinetic energy – energy of motion.
Potential energy – energy that is stored by an object as a result of its composition or position.

3. Definitions
Heat – energy transferred between two objects as a result of a difference in temperature between them.
Thermal energy – sum of the potential and kinetic energies of all the particles of an object or substance.
Temperature – measure of the average kinetic energy of the particles in a substance or an object. Celsius and Kelvin are two temperature scales commonly used.
Thermochemistry – the study of the energy changes involved in chemical and physical processes.
Law of conservation of energy – energy cannot be created or destroyed (it is only converted into another form).

4. Kinetic Molecular Theory
All matter is composed of particles. The state of a substance is determined by the energy of the particles, the distance between the particles, and the attractive forces between the particles.
Particles at higher temperature move faster, on average, than particles at a lower temperature.
Energy is involved when matter changes state. Melting, vaporization and sublimation from the solid state to the gaseous state involve the absorption of energy by particles.
Freezing, condensation and deposition from the gaseous state to the solid state involve the release of energy from particles.

5. Systems
Any sample under observation is referred to as a system. For example, the contents in a beaker is known as a system. Everything else in the entire universe is called the surroundings.
An open system can exchange both energy and matter with its surroundings (open beaker)
A closed system can exchange energy, but not matter, with its surroundings (closed beaker)
An isolated system cannot exchange energy or matter with its surroundings (closed beaker in an insulated container)

6. Variables in Heat & Energy
You cannot measure the thermal energy of a system, but you can measure the temperature. The temperature is directly related to the average kinetic energy of all the particles.
You can also measure the mass, volume, and pressure.
When you measure these variables during a reaction along with temperature, you get important information about whether heat is entering or leaving the system.
The specific heat capacity of a substance is the amount of energy needed to increase the temperature of one gram of a substance by one degree Celsius.

7. Heat Equation
The equation below enables us to calculate the amount of heat absorbed or released by a substance:
q = m c ΔT
q = heat (J)
m = mass (g)
c = specific heat capacity (J/g.°C)
ΔT = change in temperature (°C)
ΔT = Tfinal = Tinitial

8. Exothermic & Endothermic
During a chemical reaction, bonds in reactants are broken, and new bonds are formed.
The potential energy of stable compounds is lower than isolated atoms, therefore, breaking bonds requires energy and forming bonds releases energy.
If more energy is released from breaking bonds of reactants than needed to form new ones, energy is released from the system to the surroundings. This is an exothermic reaction (exo = “out of”).
Exothermic reactions would have a negative value of q (as it is a measure of absorption, and it is releasing energy).

9. Exothermic & Endothermic
If more energy is needed to break bonds than form new ones, energy is absorbed from the surroundings to increase potential energy of the system. This is called an endothermic reaction (endo = “inside”).
Endothermic reactions would have a positive value of q (as energy is taken absorbed).
Since energy cannot be created or destroyed,
qsystem + qsurroudings = 0, then qsystem = -qsurroundings.

10. Example
Example – when a 1.25kg sample of water was heated in a kettle, its temperature increased from 16.4°C to 98.9°C. How much heat did the water absorb?

11. Calorimetry
Calorimetry is the process of measuring energy changes during a physical or chemical reaction.
A calorimeter is a devised used to measure these energy changes. It consists of a well-insulated reaction chamber, a tight-fitting cover with insulated holes for a thermometer, and some mechanism to stir the calorimeter contents.
A simple calorimeter can be maid from two polystyrene cups and a lid to hold a stirrer and thermometer. The inner cup holds the system, usually a liquid, and the outer cup provides insulation.

12. Simple Calorimeter

13. Bomb Calorimetry
To study gases, a more complex calorimeter is used, called a bomb calorimeter. It is very rigid and tightly sealed to ensure no gas escapes while the volume remains constant.
The reaction produces a gas, the pressure in the calorimeter will increase. Bomb calorimeters are mostly used to study combustion reactions.

14. Bomb Calorimeter

15. Laws of Thermodynamics
The First Law: The total energy of the universe of constant. Therefore, ΔEuniverse = 0
The Second Law: When two objects are in thermal contact, heat is always transferred from the object at a higher temperature to an object at a lower temperature until the two objects are at the same temperature (thermal equilibrium).
We have already learned that q represents the change in thermal kinetic energy of the surroundings.
Enthalpy is the change in potential energy of the system (ΔH). It is also known simply as enthalpy change, heat of reaction, or change in heat content. It is measured in kJ.
ΔHsystem = ± │qsurroundings│
VIDEO – Calorimetry & q = mcΔT & ΔH
3:05 – 9:05 only

16. Dissolution & Enthalpy
Physical Change.
Forces between molecules or ions of a solute must be broken to make room for solvent molecules
Forces between solvent molecules must be broken to make room fro the solute molecules
Forces must form between the solvent molecules at the solute molecules or ions.
The sum of the enthalpy changes of all these above processes is called the enthalpy of solution, or ΔHsol.
ΔH can be exothermic or endothermic, depending on the nature of the solute and solvent.

17. Solvent & Solute Reaction Potential Diagram

18. Phase Change & Enthalpy
A significant amount of heat must be added or removed for a substance to change its state (its phase).
ΔH°melt = enthalpy of melting (energy required to change solid to liquid)
ΔH°vap = enthalpy of vaporization (energy required to change liquid to gas)
ΔH°cond = enthalpy of condensation (energy released to change gas to liquid)
ΔH°fre = enthalpy of freezing (energy released to change liquid to solid)
The ° (“nought”) indicates “at standard atmospheric pressure”.
ΔH°melt = - ΔH°fre – sometimes referred to collectively as ΔH°fus (fusion)
ΔH°vap = - ΔH°cond – sometimes referred to collectively as ΔH°vap

19. Phase Change Reaction Potential Diagram

20. The Heating Curve of Water

21. The Heating Curve of Water
Between A and B – water is solid and absorbing enough heat to increase its temperature from -20°C to 0°C. You can calculate the heat needed by q = mcΔT.
Between B and C, 6.01 kJ of heat are added to water with no change in temperature. Here all the energy is being used to break the intermolecular forces between ice molecules, converting it to liquid water. The heat added is ΔH°melt.
Between C and D, the heat gained by the water increases the temperature from 0°C to 100°C.
Between D and E, 40.7 kJ of heat are added to water with no change in temperature. Here all the energy is being used to break the intermolecular forces between water molecules, converting it to water vapour. The heat added is ΔH°vap.
Between E and F, the heat that is gained by the water vapour increases the temperature above 100°C.

22. Reactions & Enthalpy
We will go through specific enthalpies of chemical reactions in the next lesson, but note that the enthalpy changes in chemical reactions are much larger than those of physical reactions, as it involves breaking bonds, versus only intermolecular forces.

23. Nuclear Fusion & Enthalpy
All nuclear reactions are exothermic, and release much more energy than any other type of chemical reaction.
Nuclear fusion is a reaction in which nuclei of small atomic mass combine to form larger, heavier nuclei.
These reactions release vast amounts of energy in stars, including the Sun, as it contains mostly hydrogen and helium.
Hydrogen isotopes of deuterium and tritium under pressure undergo fusion to form helium, and releases huge amounts of energy.
Scientists have tried to develop fusion reactors to create energy on earth, but the pressure and heat needed for it is too high and difficult to achieve economically.

24. Nuclear Fusion

25. Nuclear Fission & Enthalpy
Nuclear fission is a reaction in which large nuclei with high atomic mass are split into smaller, lighter nuclei by collision with a neutron.
All elements above 83 are unstable and can undergo fission, though it does not normally occur in nature as it requires vast quantities of energy.
Nuclear power plants use fission of uranium-235 to produce electricity.
When uranium is bombarded with neutrons, it is broken into smaller nuclei and release more neutrons, which thus in turn react with more uranium.
This is called a fission chain reaction.

26. Nuclear Fission

27. Videos – Crash Course Energy & Heat – review of lesson 1 terms
Enthalpy – prep for lesson 2
Calorimetry
(3:05 – 9:05 on calorimeters and q=mcΔT and ΔH from lesson 1, the rest as prep for lesson 2)