THERMODYNAMICSPart I (Yep, there’ll be a Part II)

Energy

The capacity to do work or transfer heat

Measured in Joules Two Types

Kinetic (motion) Potential (based on position)

1st Law of Thermodynamics

1st Law of Thermodynamics

aka Law of Conservation of Energy

Energy can be converted from one form to another but cannot be created or destroyed

Heat vs Temperature

Heat Temperature

A measure of energy content

What is transferred during a temperature change

Units Joules

Reflects random motion of particles in a substance

Indicates the direction in which heat energy will flow

Units °C or K

Energy, Heat, and Work

∆E = q + w Change in energy equals heat plus

work

Energy and heat are state functions. Work is not A state function is independent of the

pathway taken to get to that state. (only the beginning and end matter)

Sign Conventions for q, w, and ∆E

q + means system gains

heat (endothermic

)

- means system loses

heat (exothermic)

w + means work done on

system

-means work done by system

∆E + mean net gain of

energy by system

-means net loss of

energy by system

Work, Pressure, and Volume

w = -P∆V Work done

By a gas(through

exapansion)

To a gas(by

compression)

∆V + -

w - +

Chemical Energy

Endothermic Exothermic

Rxns in which energy is absorbed from the surroundings

Energy flows into the system to increase the potential energy of the system

Rxns that give off energy as they progress

Some of the potential energy stored in the chemical bonds is converted to thermal energy (random KE) through heat

Products are more stable (stronger bonds) than reactants

Enthalpy

H = E + PV H – enthalpy E – internal energy P- pressure V – volume

In systems at constant pressure, where the only work is PV, the change in enthalpy is due only to energy flow as heat ∆H = heat of rxn

Sample Problem

Indicate the sign of the enthalpy change, ∆H, in each of the following processes carried out under atmospheric pressure, and indicate whether the process is endothermic or exothermic An ice cube melts 1g of butane is combusted

Enthalpies of Reaction

∆H = Hproducts- Hreactants

1. Enthalpy is an extensive property2.  The enthalpy change for a reaction

is equal in magnitude, but opposite in sign, to ∆H for the reverse reaction.

3.  The enthalpy change for a reaction depends on the state of the reactants and products

Sample Problem

How much heat is released when 4.50g of methane gas is burned in a constant pressure system? (∆H = -890 kJ)

Calorimetry

Science of measuring heat Heat Capacity (C)

C= heat absorbed/Temp increase Specific Heat Capacity

Energy required to raise the temp of 1 gram of a substance 1 °C

Molar Heat Capacity Energy required to raise the temp of 1

mole of a substance 1 °C

Calorimetry

∆H = mc∆T or q = mc∆T m – mass (g) c – specific heat (J/g·K) ∆T – change in temperature (K)

Sample Problem

How much heat is needed to warm 250 g of water from 22°C to near its boiling point, 98°C? (The specific heat of water is 4.18 J/g·K)

What is the molar heat capacity?

Constant Pressure Calorimetry

Coffee Cup Calorimeter

Sample Problem When a student mixes 50 mL of 1.0 M

HCl and 50 mL of 1.0 M NaOH in a coffee-cup calorimeter, the temperature of the resultant solution increases from 21.0°C to 27.5°C. Calculate the enthalpy change for the reaction in kJ/mol HCl, assuming that the calorimeter loses only a negligible quantity of heat, that the total volume of the solution is 100 mL, that its density is 1.0 g/mL, and that its specific heat is 4.18 J/g·K.

Constant Volume Calorimetry

Bomb Calorimeter

Hess’s Law

In going from a particular set of reactants to a particular set of products, the change in enthalpy (∆H) is the same whether the reaction takes place in one step or a series of steps

Hess’s Law

Using Hess’s Law

1. Work backward from the final reaction

2. Reverse reactions as needed, being sure to also reverse the sign of ∆H

3. Remember that identical substances found on both sides of the summed equation cancel each other out.

Sample Problem

The enthalpy of reaction for the combustion of C to CO2 is -393.5 kJ/mol·C, and the enthalpy for the combustion of CO to CO2 is -283.0 kJ/mol·C. Using this data, calculate the enthalpy for the combustion of C to CO.

Standard State

For a compound Gaseous state

Pressure of 1 atm Pure liquid or solid

Standard state IS the pure liquid or solid Substance in a soln

Concentration of 1 M

For an element The form in which the element exists at

1 atm and 25°C

Standard Enthalpies of Formation (∆Hf°) The change in enthalpy that

accompanies the formation of one mole of a compound from its elements with all elements in their standard state

Calculating Enthalpy Change1. When a rxn is reversed, the magnitude of

∆H remain the same, but its sign changes.2. When the balanced eqn for a rxn is

multiplied by an integer, the value of ∆H must be multiplied by the same integer

3. The change in enthalpy for a rxn can be calculated from the enthalpies of formation of the reactants and products

∆H°rxn=Σ∆Hf°products - Σ∆Hf°reactants

4. Elements in their standard states are not included

For elements in their standard state, ∆Hf° = 0

Sample Problem

For which of the following reactions at 25°C would the enthalpy change represent a standard enthalpy of formation? For those where it does not, what changes would need to be made in the reaction conditions?

a) 2Na(s) + ½ O2(g) Na2O(s)

b) 2K(l) + Cl2(g) 2KCl(s)

c) C6H12O6(s) 6C (diamond) + 6H2(g) + 3O2(g)

Sample Problem

Calculate the standard enthalpy change for the combustion of 1 mol of benzene, C6H6.

Sample Problem

Compare the quantity of heat produced by combustion of 1.00g propane, C3H8, to that produced by 1.00 g benzene.