THERMOCHEMISTRY

THERMOCHEMISTRY Chapter 6

Energy = the capacity to do work or to produce heat Kinetic energy = the energy due to motion depends on mass & velocity Potential Energy = energy due to the position or composition.

Heat VS. Temperature • Temperature = reflects the random motions of particles in a substance. The more motion the higher the temperature. • Heat = Involves the transfer of energy between two objects due to a temperature difference.

System • The portions of universe that is identified • Two Types: Open system: the designated part is open to the atmosphere. Closed system: the designated part is closed to the atmosphere.

Energy transfer 2 ways 1. work = force x distance 2. heat

State Function • A property that is related only to the current conditions—There is no consideration as to how it got to the current situation. • Examples: pressure, volume, temperature, energy and enthalpy

Parts of the universe • System-reactants and products • Surrounding-everything else

First Law of Thermodynamics • Also known as the Law of Conservation of Energy • Energy cannot be created nor destroyed but may be conserved. • Concept describes the universe not a system.

Thus, energy can be lost or gained by a system. • Energy in the universe is constant. • Thermodynamics = the study of energy and its conversions

Internal energy (E) = sum of the kinetic and potential energies of all “particles” in the system. ΔE = q + w q = heat W = work

Thermodynamic Quantities • Consists of two parts: a. Number = gives the magnitude of the change b. sign 1. (+) = endothermic 2. (-) = exothermic

Energy exchanges • The energy is exchanged with the environment in terms of heat or work. • ΔE = q + w • q = (+) means that heat is added to the system • q = (-) means that heat is subtracted from the system

Work • Negative work = energy flows out of the system so the system does work on the surroundings --exothermic • Positive work = energy flows into the system so the surrounding do work on the system --endothermic • When the systems are under relatively standard conditions the effects of work is ignored.

Enthalpy of heat of reaction ΔH = ΔE + PΔV • ΔE = change in internal energy • P = pressure of the system • ΔV = change in volume of the system • ΔH = is equal to the energy flow as

ΔH = ΔHproducts – ΔHreactants • -ΔH = exothermic • +ΔH = endothermic

enthalpy • Loss or gain of heat by a system is enthalpy. (ΔH) • State Function • ΔH = Hf – Hi = qp • qp is heat associated with constant pressure

Positive value of ΔH means that the system has gained heat from the surrounding. (endo) • Negative value of ΔH means that the system has lost heat to the surroundings. (exo)

Heat of Reaction • Heat Capacity is the amount of heat required to raise the temperature of a substance 1°C. • Molar Heat Capacity is the heat capacity of one mole of the substance. • Specific Heat Capacity is the heat capacity of gram values of a substance. • The specific heat of a substance is the amount of heat required to raise 1 gram of the substance 1°C.

Specific heat • q = m x c x ΔT • q = heat • M = mass in grams • c = specific heat in J/g°C • ΔT is the difference between final and initial temperature in°C

Specific heat problem • A 2.50 kg piece of copper metal is heated from 25°C to 225°C. How much heat kJ, is absorbed by the copper. The specific heat is 0.384 J/g°C for copper. • q = 192 kJ

Hess’s law • The enthalpy of a reaction is equal to the sum of the enthalpies for each step. • Allows us to calculate the enthalpy of the reaction by using information about each reactant.

Standard enthalpy of reaction • ΔH° • Enthalpy for a reaction when all reactants and products are in their standard state. • Standard state is 25°C and 1 atm • ΔΔΔ

enthalpy of Formation • ΔHf • Represents the enthalpy change that occurs when a compound is formed from its constituent elements.

Standard enthalpy of formation • ΔH°f • References to one mole of a compound formed from its constituent elements in their standard state.

Thermodynamics Relationships of Energies of Reactants, Products and Reactions Chapter 16

Spontaneous Processes • Occurs without outside intervention • Can occur fast or slow • Ex. carbon to diamond

Entropy, S • = disorder • The driving force for a spontaneous process is an increase in entropy • Has to do with the probability everything is in order

Positional probability • Higher the positional probability the larger the entropy, +S • Increases going from a solid to a liquid, to a gas • Increases the larger the volume you have Sample 16.1 and 16.2

2nd law of thermodynamics • In any spontaneous process there is always an increase in the entropy of the universe • It occurs in one direction. ∆Suniv = ∆Ssys + ∆SSurr +∆Suniv = process is spontaneous in direction written -∆Suniv = spontaneous reverse direction

∆Ssurr and Temperature • The sign ∆Ssurr depends on the direction of the heat flow - ∆Ssurr = endothermic + ∆Ssurr = exothermic

∆Ssurr and Temperature • The magnitude of ∆S depends on temperature. -The impact of the transfer of energy as heat to and from the surroundings has greater impact at lower temperatures.

∆H review • ∆H = heat flow = change in enthalpy -∆H sys = endothermic +∆H sys = exothermic

Combining ∆H and ∆S ∆Ssurr = - ∆H / T *the minus sign changes the point of view from the system to the surroundings - For constant pressure and temperature Sample 16.4

Free energy, G ∆G = ∆H -T∆S H = enthalpy T = temperature in Kelvin (constant) S = entropy

Spontaneity and ∆G • A reaction is spontaneous if ∆G is negative and carried out under constant pressure and temperature. Sample 16.5

Entropy in Chemical Reactions • The change in positional entropy is dominated by the relative #’s of molecules of gaseous products and reactants

Does Entropy increase oRdecrease • N2 +3H2 2NH3 • H2  2H • 4NH3 + 5O2  4NO + 6H2O Sample 16.6

Third Law of Thermodynamics • The entropy of a perfect solids at 0K is zero. • An increase in motion is associated with higher entropy value.

THERMOCHEMISTRY