Chapter 10 - Thermochemistry

10.1 Energy as a Reactant or Product

Kinetic Energy: energy of an object in motion

Potential Energy: energy of an object because of position or composition (chemical energy)

Internal Energy: sum of all kinetic and potential energies of a system

Thermodynamics: study of energy and its transformations

First Law of Thermodynamics: energy cannot be created or destroyed but it can change from one form to another

Universe = system + surroundings

Work: force over a distance (F*d)

Internal energy of a system decreases when heat flows out of a system or if work is done by the system on its surroundings. The internal energy of a system increases when heat flows into a system or if work is done on a system. ΔE = q+w, where q is heat and w is work.

State function: internal energy depends on the potential and kinetic energy of a system, not the means of which these energies changed; ΔE = Efinal - Einitial

Heat: energy in the process of being transferred from a higher temperature object to a lower temperature one

Thermal Energy: total internal energy of a system is related to its heat

Translational Kinetic Energy: motion of particles as they collide with each other; average kinetic energy is directly proportional to its absolute kelvin temperature

Non-translational Kinetic Energy: rotational kinetic energy and vibrational kinetic energy

10.2 Transferring Energy and Doing Work

Open System: matter and energy can be transferred between a system and its surroundings

Closed System: only energy can flow between system and surroundings

Isolated System: neither matter nor energy can flow between system and surroundings

Exothermic: transfer of energy from system to surroundings, usually detected by an increase in temperature ΔH < 0

Endothermic: transfer of energy from surroundings to system ΔH > 0

Pressure-Volume Work: work associated with the expansion or compression of a gas ΔE=−PΔV

10.3 Enthalpy and Enthalpy Changes

Enthalpy: measure of the total energy of a system H=E+PV

Enthalpy Change (ΔH): quantity of heat transferred in or out from a system during a chemical reaction or phase change ΔH = ΔE + Δ(PV)=qp

Enthalpy of Fusion: energy required to convert one mole of a substance at its melting point into the liquid state

Enthalpy of Vaporization: energy required to convert one mole of a liquid substance at its boiling point into the gaseous state

10.4 Heating Curves and Heat Capacity

specific heat capacity: quantity of thermal energy required to riase the temperature of one gram of a substance by 1 degree Celsius at a constant pressure

Q=mcΔT, where m is mass, c is specific heat capacity and ΔT is the change in temperature

molar heat capacity: quantity of energy required to raise the temperature of one mole of a substance by one degree Celsius at a constant pressure, q=ncΔT where c is the molar heat capacity

The relative lengths of line segments in a heat curve represent phase changes; longer lengths correspond with greater enthalpies of phase changes

10.5 Enthalpies of Reaction and Calorimetry

Calorimetry: changes in temperature with a known heat capacity, a calorimeter, are used to determine the energy released or absorbed by a process occurring inside the calorimeter

Enthalpy of Reaction: magnitude of energy absorbed or released during a chemical reaction; proportional to quantity of reactants consumed; difference in enthalpy between products and reactants

Thermochemical Equation: adding enthalpy change that accompanies mole ratio

Coffee Cup Calorimeter: q*(rxn)=-q(calorimeter)=C(calorimeter)*ΔT=(4.18)mΔT

Bomb Calorimeter: used to measure energy released during a combustion reaction; sealed vessel capable of withstanding high pressures that is submerged in water in a heavily insulated container, oxygen is introduced and mixture is ignited. the thermal energy generated by the reaction flows into the walls of the bomb and into the water of the calorimeter

• The heat capacity of the calorimeter (calorimeter constant) also needs to be known because it absorbs some of the energy from the reaction q=ΔE

10.6 Hess’s Law and Standard Enthalpies of Reaction

Hess’s Law: change in enthalpy that accompanies a process that occurs in more than one step is the sum of the enthalpy changes that occur in each of those steps

• we can combine chemical equations by describing reactions in a way that gives us an equation for the final reaction by multiplying coefficients and reversing reactions, which directly impact the enthalpy value

Standard Enthalpy of Reaction: enthalpy change that accompanies a reaction under standard conditions: 1 bar of pressure and solutions with a concentration of 1 M. No standard temperature, but 25 C is common.

Standard State: most stable physical state of a substance under standard conditions

10.7 Enthalpies of Reaction from Enthalpies of Formation and Bond Energies

Standard Enthalpy of Formation: relative enthalpy change for a formation reaction; pressure of 1 bar when one mole of substance is formed from constituent elements in their standard states - will be a value of 0.

Breaking bonds is endothermic and forming bonds is exothermic; exothermic reactions occur when more energy is released forming bonds than energy needed to break bonds; endothermic reactions occur when more energy is consumed in breaking bonds than released in forming bonds

Bond energy: enthalpy change that occurs when one mole of bonds in the gas phase is broken

10.9 More Applications of Thermochemistry

Fuel Values: energy produced during the combustion of 1 gram of a substance

• as the number of carbon atoms per molecule increases, the hydrogen to carbon ratio decreases, making fuels less effective

Fuel Density: quantity of energy released per unit volume of a liquid fuel