Thermodynamics They study of energy and its transformations

Energy Defined as the ability to do work or produce heat Defined as the ability to do work or produce heat 2 types of energy 2 types of energy –Kinetic Energy – energy of motion (thermal energy or heat) –Potential Energy – energy of position (example – chemical bond energies) –Units – Joules = kg m 2 /sec 2

First Law of Thermodynamics – Energy is always conserved. Energy cannot be created or destroyed, it can only be converted into matter or another form of energy. E = mc 2 – Energy is always conserved. Energy cannot be created or destroyed, it can only be converted into matter or another form of energy. E = mc 2 We must keep track of the energy exchanged in a process. So we divide the universe into two parts. We must keep track of the energy exchanged in a process. So we divide the universe into two parts.

2 Parts of the Universe Are... The SYSTEM - part of the universe under study (generally a chemical process) The SYSTEM - part of the universe under study (generally a chemical process) The SURROUNDINGS – the rest of the universe (like the container where a chemical reaction is taking place. The SURROUNDINGS – the rest of the universe (like the container where a chemical reaction is taking place. In chemistry we always focus on the SYSTEM!! In chemistry we always focus on the SYSTEM!! In a closed system - energy can be exchanged with the surroundings but not matter. In a closed system - energy can be exchanged with the surroundings but not matter. In an open system – energy and matter can be exchanged with the surroundings. In an open system – energy and matter can be exchanged with the surroundings.

Internal Energy Internal energy is the sum of the kinetic and potential energy of a system. Internal energy is the sum of the kinetic and potential energy of a system. INTERNAL ENERGY = KE + PE INTERNAL ENERGY = KE + PE

Internal Energy We can only measure a change in energy (  E) We can only measure a change in energy (  E)  E = E final – E initial  E = E final – E initial It is a STATE FUNCTION – it only depends on current conditions. It does not matter how it got there. It is a STATE FUNCTION – it only depends on current conditions. It does not matter how it got there.

Heat and Work There are two ways for a system to exchange energy with its surroundings There are two ways for a system to exchange energy with its surroundings Heat (q) = the amount of energy transferred between two objects Heat (q) = the amount of energy transferred between two objects Work (w) = Energy that is a force acting over a distance w = F x  d Work (w) = Energy that is a force acting over a distance w = F x  d

Heat and Work Heat and work are not state functions and are therefore dependent on the pathway. Heat and work are not state functions and are therefore dependent on the pathway. The sum of heat and work = the change in Internal energy of a system The sum of heat and work = the change in Internal energy of a system  E = q + w

Sign Conventions for work and heat all thermodynamic variables include a number and a sign. all thermodynamic variables include a number and a sign. Heat (q) = (+) then heat is absorbed by the system (surroundings cool down Heat (q) = (+) then heat is absorbed by the system (surroundings cool down Heat (q ) = (-) then heat is released from the system (and heats up the surroundings) Heat (q ) = (-) then heat is released from the system (and heats up the surroundings) Work (w) = (+) then work is done on the system by the surroundings Work (w) = (+) then work is done on the system by the surroundings Work (w) = (-) then work is done by the system on the surroundings Work (w) = (-) then work is done by the system on the surroundings

Change in Energy  E = (+) then the system gains energy  E = (+) then the system gains energy  E = (-) then the system loses energy  E = (-) then the system loses energy Endothermic = energy is absorbed by the system by the surroundings as heat Endothermic = energy is absorbed by the system by the surroundings as heat Exothermic = energy is released by the system to the surroundings as heat. Exothermic = energy is released by the system to the surroundings as heat.

Potential energy of product bonds are less than Reactants. Products have stronger or more stable bonds Potential Energy of the Products are higher than reactants. Reactants have stronger more stable bonds.

Pressure Volume Work For chemical processes – work is done by a gas (through expansion) or work is done on a gas (by compression.) Using P = force/area and w= force x d Using P = force/area and w= force x d you can derive work = P  V you can derive work = P  V (P = external pressure) (P = external pressure) For an expanding gas w = - P  V For an expanding gas w = - P  V (work is being done on the surroundings) (work is being done on the surroundings)

Enthalpy Enthalpy (H) is the amount of heat flow under conditions when only Pressure volume work is done. Enthalpy (H) is the amount of heat flow under conditions when only Pressure volume work is done. PV work is only done under conditions of constant pressure with a volume expansion (+  V) or volume compression (-  V) PV work is only done under conditions of constant pressure with a volume expansion (+  V) or volume compression (-  V)  H =  E + P  V

 H =  E + P  V  H =  E + P  V (we also know  E = q p + w) (we also know  E = q p + w) Since  E = q p - P  V Since  E = q p - P  V (q p = heat at constant pressure) (q p = heat at constant pressure)  H = q p + P  V - P  V so.....  H = q p + P  V - P  V so.....  H = q p  H = q p Which means we can measure enthalpy by measuring the heat flow (or change in temperature) at constant pressure! Which means we can measure enthalpy by measuring the heat flow (or change in temperature) at constant pressure!

Heat at Constant Volume vs. Heat at Constant Pressure Constant Volume Constant Volume –Bomb Calorimeter q v = ∆E Volume must change for work to be done, so no work term! Constant Pressure Constant Pressure –Coffee Cup Calorimeter – volume of a gas can change. q p = ∆H q p = ∆H

Example -Hydrogen and oxygen gas are placed together in a container with a moveable piston, and the gases are ignited. As water is being produced, 1350 Joules of heat are lost to the surroundings. The gases expand, and the piston is moved, producing 350 Joules of work. What is the change in internal energy in this system? Example -Hydrogen and oxygen gas are placed together in a container with a moveable piston, and the gases are ignited. As water is being produced, 1350 Joules of heat are lost to the surroundings. The gases expand, and the piston is moved, producing 350 Joules of work. What is the change in internal energy in this system? Heat = q = J Negative because the heat is lost to the surroundings – Exothermic Reaction! Work = w = -350 J Negative because the work is done by the system The piston, part of the system, is expanding and working against the surrounding pressure!

∆E = q + w ∆E = q + w = (-1350J) + (-350J) = (-1350J) + (-350J) = J of energy transferred to the surroundings. = J of energy transferred to the surroundings. Exothermic Process! Gas expands to keep constant pressure! Exothermic Process! Gas expands to keep constant pressure! H2O2H2O2