Thermochemistry

Energy Review Energy – the capacity to do work or transfer heat. Work – the energy used to move an object against a force. Heat – energy used to cause the temperature of an object to rise. Kinetic energy – energy of motion. Potential energy - stored energy based on the position of the particles. Chemistry is all about electrostatic energy.

Electrostatic Interactions

Other Definitions and Relationships Units of energy are Joules and calories J = 1 cal; 1000cal = 1Cal The calorie is the amount if energy required to raise one gram of water one degree Celsius. Law of Conservation of Energy. (1 st Law of Thermo) The system is the part that we are studying. The surroundings is everything outside of the system. Closed systems exchange energy but not matter. Isolated systems don’t exchange anything.

Work vs. Heat Work is the energy required to put a force on an object to move it. Force is any push or pull on an object. Heat is energy transferred from a hot object to a cold one. (Endo and exothermic) Internal energy of the system is the total of all potential and kinetic energies of the system. Rather than focus on total energy, we focus on changes in energy. ∆E = q + w

State Functions State functions depend only on the present state of the system. The path used to get to that state is irrelevant. Energy is a state function, so any changes in energy are independent of path as well. Energy changes are state functions, but whether that change was through heat and work is not. Enthalpy changes are state functions as well which is why Hess’ Law works. Also explains why enthalpy changes can be found through Enthalpy of Formation calculations. The reaction mechanism has no effect.

Hess’ Law N 2(g) + 2O 2(g)  2NO 2(g)  H = 68 kJ/mol Can be determined by: N 2(g) + O 2(g)  2NO (g)  H = 180 kJ/mol 2NO (g) + O 2(g)  2NO 2(g)  H =- 112 kJ/mol N 2(g) + 2O 2(g)  2NO 2(g)  H = 68 kJ/mol

Enthalpy of Formation Amount of energy required to make one mole of a substance from elements. ∆H o f = standard enthalpy of formations. All ∆H o f for elements = 0 kJ/mol if they are under standard conditions. ∆H = ∆H o f (products) - ∆H o f (reactants) These values will be provided when needed. Since it is kJ/mole, then the ∆H o f values must be multiplied by the coefficients in the balanced equation to get energy.

Enthalpy Enthalpy is the total energy of the system plus the product of pressure and volume. H = E + PV PV is pressure-volume work When the pressure is constant, w = -P∆V ∆H = ∆E + P∆V and ∆E = q + w Since w = -P∆V Then for constant pressure systems, ∆H = q, or simply ∆H = change in heat.

Calorimetry Calorimetry is a way of measuring heat flow. We measure the temperature change of a substance to determine the heat flow. We use water since it is easy to measure to determine the energy change of something we can’t measure. Heat capacity is the amount of energy an object needs to raise its temperature 1 o C Specific heat capacity is the amount of energy per gram.

Calorimetry ∆H = m C p ∆T This equation can be used if the substance is not reacting or changing phase. 2 types of situations – A substance at a different temperature than water is put in the water and left. – A reaction is taking place in water In both cases, ∆H (H2O) = - ∆H (system) Reactions in solutions will require solution to be treated as water.

Types of Calorimetry Constant pressure – done in a container that will change volume, or done in an open air container. Constant volume – Bomb Calorimeter. – Seal reaction and water in a container. – Since V doesn’t change, no work is done, so ∆E = q – Calculate the heat capacity of the container. – ∆E = ∆H essentially.