Chapter 6

= 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.

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.

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.

2 ways

1. work = force x distance

2. heat

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

System-reactants and products

Surrounding-everything else

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

= sum of the kinetic and potential energies of all “particles” in the system.

ΔE = q + w

q = heatW = work

Consists of two parts:a. Number = gives the magnitude of the change

b. sign1. (+) = endothermic2. (-) = exothermic

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

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.

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

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 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.

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

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

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.

ΔH°

Enthalpy for a reaction when all reactants and products are in their standard state.

Standard state is 25°C and 1 atm ΔΔΔ

ΔHf

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

ΔH°f

References to one mole of a compound formed from its constituent elements in their standard state.

Relationships of Energies of Reactants, Products and Reactions

Chapter 16

Occurs without outside intervention

Can occur fast or slow

Ex. carbon to diamond

= disorder The driving force for a spontaneous

process is an increase in entropy Has to do with the probability

everything is in order

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

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

1) The sign ∆Ssurr depends on the direction of the heat flow

- ∆Ssurr = endothermic

+ ∆Ssurr = exothermic

2) 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 = heat flow = change in enthalpy

-∆H sys = endothermic+∆H sys = exothermic

∆Ssurr = - ∆H / T

*the minus sign changes the point of view from the system to the surroundings

- For constant pressure and temperature

∆G = ∆H -T∆S

H = enthalpy T = temperature in Kelvin (constant)S = entropy

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

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

1) N2 + 3H2 2NH3

2) H2 2H

3) 4NH3 + 5O2 4NO + 6H2O

The entropy of a perfect solids at 0K is zero.

An increase in motion is associated with higher entropy value.

When a solid melts When a solid dissolves When a solid or liquid becomes a gas When a gaseous chemical reaction

produces more molecules When the temperature increases

Given in appendix 4

∆Soreaction = Σnp So – Σnr So

We cannot measure ∆G directly we must use other measured quantities.

The more negative the value of ∆G the further a reaction will go to the right to reach equilibrium

1) Using the formula∆Go= ∆Ho-T∆S

2) By taking advantage of the fact that ∆G is a state function and solving like Hess’s law

3) By using ∆Gof, standard free energy of

formation

∆Go = Σnp Gfo (products)– Σnr

Gfo(reactants)

Enthalpy is not affected by pressure

Entropy is dependent on pressure because entropy is dependent on volume.

Slarge volume> S small volume so

Slow pressure > Shigh pressure

∆G = ∆Go + RT lnQ ∆G = free energy change for rxn for

specified pressure ∆Go= free energy change at standard

pressure R = 8.31 J/K * mol T = temp in Kelvin Q = reaction quotient

The equilibrium point occurs at the lowest value of free energy available to the reaction system

∆GO= 0 the system is at equilibrium, K = 1, pressure = 1 atm

∆Go < O not at equilibrium, shift to the right, K >1, pressures of products > 1 atm, pressure of reactants < 1

∆G0 > 0 not at equilibrium, K <1, shift left because reactants have less energy

Free energy can tell us how much work can be done with a given process.

- ∆G means the amount of free energy available \ to do use work

+ ∆G is the minimium amount of work that must be expended to make the reaction occur.