1. Chapter 17 Thermochemistry
1st Law of Thermodynamics: Energy cannot be created or destroyed!!!!
Chapter 17 Thermochemistry

2. Virtually every chemical reaction is accompanied by a change in energy
Virtually every chemical reaction is accompanied by a change in energy. Chemical reactions usually absorb or release energy as heat. Thermochemistry is the study of the transfers of energy as heat that accompany chemical and physical changes.

3. Heat
Heat can be thought of as the energy transferred between the samples of matter because of differences in their temperatures. Energy transferred as heat always moves spontaneously from matter at a higher temperature to matter at a lower temperature.

4. Temperature
Temperature if a measure of the average kinetic energy of the particles in a sample of matter.
The greater the kinetic energy, the higher the temperature.
The ability to measure temperature is based on heat transfer.
Temperature is measured in Joules, as are most other forms of energy.

5. Measuring Heat
The energy absorbed or released in a chemical or physical change is measured in a calorimeter.
In one kind of calorimeter, known quantities of reactants are sealed in a reaction chamber, which is immersed in a know quantity of water in an insulated vessel. Therefore, the energy given off (or absorbed) during the reaction is equal to the energy absorbed (or given off) by the known quantity of water. The amount of energy is determined from the temperature change of the known mass of surrounding water.

6. Heat Capacity and Specific Heat
The quantity of energy transferred as heat during a temperature change depends on the nature of the material changing temperature, the mass of the material and the size of the temperature change.
Specific heat is the amount of energy required to raise the temperature of one gram of a substance by one Celsius degree.
cp = __q___
M x ∆T
cp= specific heat in J/(g  K)
Q = energy lost or gained in Joules
M= mass of sample in grams
∆T= change in temperature in Kelvin

7. Practice
A 40.0 gram sample of glass was heated from 274 K to 314 K, a temperature increase of 40 K, and was found to have absorbed 32 J of energy as heat.
What is the specific heat of this type of glass?
How much energy will the same glass sample gain when it is heated from 314 K to 344 K?

8. Heat of Reaction 2H2(g) + O2(g) → 2H2O
The heat of reaction is the quantity of energy released or absorbed as heat during a chemical reaction.
You can think of heat of reaction as the difference between stored energy of the reactants and the products.
If a mixture of hydrogen and oxygen is ignited, water will form and energy will be released explosively. The energy that is released comes from the reactants as they form products. Because energy is released, the reaction is exothermic, and the energy of the product, water, must be less that the energy of the reactants before ignition.
2H2(g) + O2(g) → 2H2O
The equation does not tell you that energy is evolved as heat during the reaction. Experiments have show that kJ of energy are evolved when 2 mol of gaseous water are formed at K from its elements.

9. Thermochemical equations
An equation that includes the quantity of energy released or absorbed as heat during the reaction as written is called a thermochemical equation. Modifying the chemical equation from the previous slide to show the amount of energy produced gives the following expression: 2H2(g) + O2(g) → 2H2O kJ The quantity of heat released in this or any reaction depends on the amounts of reactants or products. In this example, the quantity of energy released during the formation of water is proportional to the quantity of water formed. Producing twice as much water vapor would require twice as many moles of reactants and would release 2 X kJ of energy as heat: 4H2(g) + 2O2(g) → 4H2O kJ

10. Producing twice as much water vapor would require twice as many moles of reactants and would release 2 X kJ of energy as heat: 4H2(g) + 2O2(g) → 4H2O kJ Producing one-half as much water would require one-half as many moles of reactants and would release only on-half as much energy H2(g) + ½ O2(g) → H2O kJ **fractional coefficients are sometimes used in thermochemical equations.

11. Enthalpy Change
Enthalpy change is the amount of energy absorbed or lost by a system as heat during a process at constant pressure. Enthalpy change is always the difference between the enthalpies of the products and the reactants ∆H = Hproducts – Hreactants

12. Exothermic vs. endothermic
For exothermic reactions, ∆H is always given a minus sign because the system loses energy.
2H2(g) + O2(g) → 2H2O ∆H = kJ
Ouch, that’s hot!!
For endothermic reactions, ∆H is always given a positive value because the system gains energy.
2H2(g) + O2(g) → 2H2O ∆H = kJ
Wow, that’s cold!

14. Energy diagrams
Reactants to top of curve, Energy (heat) is being put in to break bonds in the reactants.
At the top of the curve, the bonds in the reactants have been broken. The amount of energy put in to break these bonds is called the activation energy.
Going from the top of the curve to the products, you are going down the energy scale. Energy (heat) is given out as bonds form in the products.
The reactants are higher up the energy scale than are the products. The amount of energy (heat) you need to put in  is less than the amount of energy (heat) you get out. This is a typical exothermic reaction.
The difference in energy levels between the reactants and the products is given the symbol ∆H This is the amount of heat given out (or taken in) during the reaction. For an exothermic reaction,  ∆H is negative. For an endothermic reaction,  ∆H is positive.

15. Heat of Formation
The energy released or absorbed as heat when one mole of a compound is formed by combination of its elements. To make comparisons meaningful, heats of formation are given for the standard states of reactants and products- these are the states found at atmospheric pressure and usually room temperature. Thus, the standard state of water is liquid and the standard state of iron is solid.

16. Calculating Heats of Reaction
Hess’s Law state that the overall enthalpy change in a reaction is equal to the sum of the enthalpy changes for the individual steps in the process.

17. Hess’s Law calculations simplified
Balance the equation.
Look up the heat of formation for each substance in the equation. Keep in mind that elements do not have heats of formation.
Multiply the heats of coefficients by the coefficients from the balanced equation.
∆H = Hproducts – Hreactants

18. ∆H = Hproducts – Hreactants 33.2kJ/mol – 90.29 kJ/mol = -57.1 kJ/mol
Calculate the heat of reaction of nitrogen monoxide gas, NO, to form nitrogen dioxide gas, NO2, as given in the following thermo chemical equation.
NO (g) + ½ O2(g) → NO2(g)
kJ/mol kJ/mol kJ/mol
∆H = Hproducts – Hreactants
33.2kJ/mol – kJ/mol = kJ/mol
This reaction is exothermic.
**note that zero is the assigned value for the heats of formation of elements in their standard states

19. Reaction Rates- Rate influencing factors
Nature of Reactants
Surface area
Temperature
Concentration
Prescence of Catalysts