Thermal Chemistry
Day 1:
Thermal Chemistry: The study of the heat changes that occur during a chemical reaction and phase changes.
Energy: The ability to do work or produce heat
v The unit for energy is the Joule (J)
Two Forms:
v Potential Energy: The energy is the energy due to position. Ex: A ball sitting on a hill has potential energy.
v Kinetic Energy: The energy of motion. Ex. The ball moving down the hill
Chemical Potential Energy: The energy stored in a substance because of its composition (what it is made of).
v The stronger the bonds between atoms the more chemical potential energy
States of Matter:
*How does kinetic energy and potential energy change as particles are heated?
*How does pressure and temperature determine state of matter?
*How does temperature and concentration of particles affect vapor pressure?
Heat: a form of energy that always flows from warmer objects or surroundings to cooler surrounding or objects.
v Symbol for heat = q
v To measure heat we use the calorie as our unit (cal)
v Calorie: The amount of heat required to raise 1g of pure water 1 degree Celsius.
v 1000 cal= 1 kcal (kilocalorie)
v 1Joule= 0.2390 cal
v 1cal= 4.184 Joules
Temperature- a measure of the average kinetic energy ( Joule, Celsius and Kelvin)
Endothermic Reactions: reactions that require heat to occur.
Exothermic Reactions: reactions that produce heat.
Calorimeter: an insulated device used to measure the release or absorption of heat in a chemical or physical change.
Closed System: a system in which matter is conserved inside a container but energy is transferred between different parts of the system. No energy is lost or gained.
Specific Heat: the heat that is needed to raise the temperature of 1gram of a substance by one degree.
Q=c x m x T
Q= heat transferred
C= the specific heat of substance
M= mass of substance
T= Final temperature- Initial Temperature
Example: What is the heat absorbed by 12g of water if the temperature changes from 13 oC to 45oC?
What is the final temperature of copper if 35J is added to 3g of copper at 47 oC?
Day 2:
Phase Diagrams: graph that shows how temperature and pressure affect phase change.
Phase Diagram of Water: Phase Diagram for CO2
*Know the definitions of the following vocabulary: melting, freezing, vaporization, condensation, sublimation and deposition. Be able to label a phase diagram using these vocabulary words.
Law of conservation of Energy: in any chemical reaction or physical process (phase change), energy cannot be created or destroyed but is changed from one form to another.
If a substance gets hotter something else must get colder.
Heat lost = heat gained
mc Tlost = - mc Tgained
For example: a piece of iron with a mass of 21.5 g at a temperature of 100.0 oC is dropped into an insulated container of water. The mass of the water is 132 g and its temperature before adding the iron is 20.0 oC. What is the final temperature of the system?
Day 3:
Boiling point: is the temperature at which vaporization/condensation occurs.
Heat of Vaporization:heat needed to vaporize or condense a substance; amount of heat need to vaporize or condense 1 g or kg of a substance.
Q = mHV
Q= energy
m = mass
HV = heat of vaporization
Example: How much heat would it take to change 3 kg of water at 100 degrees C to steam? (Heat of vaporization is 2.26 x 10^6 J/kg)
Freezing point of a substance is the temperature at which freezing/ melting occurs. Heat of Fusion: amount of heat that must be removed for a liquid to freeze or added for a solid to melt; amount of heat necessary to freeze/melt 1 g or kg of a substance.
Example: The heat of fusion of water is 3.34 x 10^5 J/kg
Q = mHF
Q = heat energy
m = mass
HF = heat of fusion
Example: How much heat would it take to melt 0.5 kg of ice at 0 oC?
Heating and Cooling Curve: represents the changes in the state of matter as energy is added.
Phase of matter and change? Specific Heat or Heat of…?
1. 1.
2. 2.
3. 3.
4. 4.
5. 5.
Example: What is the total amount of heat needed to raise the temperature of H2O from
-5oC to 100 oC?
Day 4:
Collision Theory: in order for reactants to react (change into products) they must collide or come into contact with each other. In order for a collision to be effective (cause the change) it must have two things :
a. Enough energy to break the bonds of the reactant particles andto form the bonds of the product particles. This is called activation energy.
b. Proper orientation - reactant particles must collide in the correct positions to break and reform bonds
Potential Energy Diagrams: represents the change in the energy during a reaction.
Day 5:
Reaction rate: number of reactant particles that react to form product particles per unit of time. Four factors which affect reaction rate are :
a. Temperature - the higher the temperature, the faster the particles move and the more likely the reactant particles will have enough energy to overcome the energy barrier to form products.
b. Concentration - the higher the concentration the higher the reaction rate. The more particles present, the more likely they are going to collide with each other and react.
c. Particle size - the smaller the particle size, the greater the surface area. The increase in surface area makes it more likely that reactant particles will collide to react to form product particles.
d. Catalysts - catalysts reduce the activation energy needed for a reaction to occur and so increase the number of effective collisions. Our body has many enzymes (organic catalysts) which speed up reaction rates allowing the body to control its metabolism.
Reversible reactions: reactants change into the products and the products change into the reactants simultaneously. Reversible reactions are indicated by a double yields sign between reactants and products as follows:
2SO2(g) + O2(g)D 2SO3(g)
Chemical equilibrium: forward and backward rates of a reversible reaction occur at the same rate causing no net change in the amounts of reactants and products present.
The equilibrium constant (Keq) for a reaction at equilibrium is the ratio of the product concentrations to reactant concentrations, each raised to the power of the number of moles there are in the balanced equation. For example:
2SO2(g) + O2(g) D 2SO3(g) 4 H2(g) + CS2(g) D CH4(g) + 2 H2S(g)
Keq >1: More products than reactants at equilibrium; reaction is more complete
Keq < 1: More reactants that products at equilibrium; reaction is less complete
Keq = 1: High concentrations of both reactants and products are present.
Day 6:
Le Châtelier's principle states that if a stress is applied to a system at equilibrium, the system will shift its equilibrium position to relieve the stress.
Factors which can create stress on a system at equilibrium are :
Concentration : changing the concentration of reactants or products by the addition or removal of reactants or products will shift the equilibrium position. For example if more SO2 is added to the following reaction
2SO2(g) + O2(g) D 2SO3(g)
the equilibrium position would shift to the right (products) causing more SO3 to form and removing some of the added SO2.
Temperature : changing the temperature of a reaction at equilibrium will shift the equilibrium position to undo the stress of the temperature change. You can treat temperature just like a reactant or product. In the following exothermic reaction the heat energy can be treated like a product.
2SO2(g) + O2(g) D 2SO3(g) + heat
If heat is added (temperature raised) the equilibrium position will shift to the left to remove the added heat and favor the formation of more of the reactants.
Pressure : changing the pressure on a system at equilibrium will change the equilibrium position of reversible reactions with unequal numbers of moles of gaseous reactants and products. It will not affect nongaseous reactants or products. For the following reaction a change in pressure would not affect the equilibrium position. The reason is that there are equal numbers of moles of gaseous reactants and products.
2 HI(g) D H2(g) + I2(g)
Example: If pressure were added to the following equation it would shift to the right because there are fewer moles of gas.
4 H2(g) + CS2(g) D CH4(g) + 2 H2S(g)