Thermodynamics X Unit 9

Energy: Basic Principles  Thermodynamics – the study of energy changes  Energy – the ability to do work or produce heat Kinetic Energy – energy of motion Potential Energy – energy due to position or composition

Law of Conservation of Energy  A.k.a. First Law of Thermodynamics  Energy can be converted from one form to another but can’t be created or destroyed This means the total energy of the universe is CONSTANT!

Heat vs. Temperature  Temperature – measure of the random motion of a substance Temperature is proportional to kinetic energy (it is a measure of the average kinetic energy in a substance)  Heat (q) – flow of energy due to a temperature difference

System vs. Surroundings  A system is the part of the universe we are studying.  The surroundings are everything else outside of the system.

Direction of Heat Flow  Heat transfer occurs when two objects are at two different temperatures.  Eventually the two objects reach the same temperature At this point, we say that the system has reached equilibrium.

Thermal Equilibrium  Heat transfer always occurs with heat flowing from the HOT object to the COLD object.

Exothermic vs. Endothermic  Exothermic process  heat is transferred from the system to the surroundings Heat is lost from the system (temperature in system decreases)  Endothermic process  heat transferred from the surroundings to the system Heat is added to the system (temperature in system increases)

Units of Energy  Joule (J) is the SI unit of energy & heat One kilojoule (kJ) = 1000 joules (J)  calorie (cal) = heat required to raise the temperature of 1.00 g of water by 1 °C 1 calorie = J

Units of Energy  Food is measured in Calories (also known as kilocalories) instead of calories 1 Cal = 1 kcal = 1000 calories

Units of Energy 3800 cal = __________ Cal = _________ J

Units of Energy The label on a cereal box indicates that 1 serving provides 250 Cal. What is the energy in kJ?

Heat Transfer Direction and sign of heat flow – MEMORIZE! ENDOTHERMIC: heat is added to the system & the temperature increases ( +q ) EXOTHERMIC: heat is lost from the system (added to the surroundings) & the temperature in the system decreases ( -q )

Specific Heat (Specific Heat Capacity)  Specific Heat (C) - The quantity of heat required to raise the temperature of one gram of a substance by 1 °C  Units: J/(g · °C) or J/(g · K) cal/(g · °C) or cal/(g · K)

Examples of Specific Heat At the beach, which gets hotter, the sand or the water? Higher specific heat means the substance takes longer to heat up & cool down!

Examples of Specific Heat  Specific heat (C)= the heat required to raise the temperature of 1 gram of a substance by 1 °C C water = J/(g  °C) C sand = J/(g  °C)

Calculating Changes in Thermal E q = mCΔT q = heat (J) m = mass (g) C = specific heat capacity, J/(g  °C) ΔT = change in temperature, T final – T initial (°C or K) ***All units must match up!!!***

Example How much heat in J is given off by a 75.0 g sample of pure aluminum when it cools from 84.0°C to 46.7°C? The specific heat of aluminum is J/(g°C). q = mCΔt

Example What is the specific heat of benzene if 3450 J of heat are added to a g sample of benzene and its temperature increases from 22.5 °C to 35.8 °C? q = mCΔt

Enthalpy  Enthalpy (H) The heat content of a reaction (chemical energy)  ΔH = change in enthalpy The amount of energy absorbed or lost by a system during a chemical process  ΔH = ΔH final - ΔH initial

Representation of Enthalpy as a Graph

Two Ways to Designate Thermochemical Equations Endothermic: a) H 2 (g) + I 2 (s)  2 HI (g) ΔH = 53.0 kJ b) H 2 (g) + I 2 (s) kJ  2 HI (g)

Two Ways to Designate Thermochemical Equations Exothermic: a) ½ CH 4 (g) + O 2 (g)  ½ CO 2 (g) + H 2 O (l) ΔH = kJ b) ½ CH 4 (g) + O 2 (g)  ½ CO 2 (g) + H 2 O(l) kJ

Two Ways to Designate Thermochemical Equations Note the meaning of the sign in ΔH in the equations above!! Endothermic: ΔH = + Exothermic: ΔH = -