AP Chemistry 12

Energy Relationships in Chemistry

Thermochemistry

Thermodynamics – the study of energy and its transformations.

Thermochemical changes – energy changes associated with chemical reactions.

System that specify part of the universe of interest to us.

Surroundings the part of the universe not contained in the system.

3 types of Systems

open system exchanges mass and energy

closed system exchanges energy but no mass

isolated system no exchange of either mass or energy

Three Types of Systems

Open system Closed System

cork

Isolated System

insulation

Different Types of Energy

Energy – the ability to do work.Thermal energy – associated with the

random motions of atoms and molecules Heat energy – transfer of thermal energy

between two objects at different temperature.

Energy (cont’d)

Chemical energy – energy stored within the structural units of chemical substance.

Potential energy – the ability of an object to do work because of its position in a field of force.

Kinetic Energy – the work that can be performed by a moving object.

The unit of energy

1 Joule (J)

=1 kg m2/s2

An older unit of energy

1 calorie (cal)

= 4.184 J exactly

The Law of Conservation of Energy

The law of conservation of energy Energy is neither created nor destroyed in

ordinary chemical and physical processes.Converted from one type into another.

This is also stated in terms of the first law of thermodynamics.

if EEE E = internal energy change of the system

Ef and Ei the energy of the final and initial states, respectively

First Law of Thermodynamics

Chemical reactions either absorb or release energy.

Two termsExothermic reaction heat is released to the

surroundings.Endothermic reaction heat is supplied to the

system by the surroundings.

Exothermic

System

Heat

An Exothermic Process

surroundings

Endothermic

Heat

System

An Endothermic Process

surroundings

The First Law Restated

chemical systems – examine the conversion of heat energy into work.

wqEEE if

Signs for Heat and Work

Work done by system on surroundingsw ‘-’

Work done by surroundings on systemw ‘+’

q < 0, heat flows to surroundingsExothermic ‘-’

q > 0, heat flows to systemEndothermic ‘+’

State and Path Functions

E, H, V are examples of state functions. State functions – numerical value doesn’t

depend on how the process is carried out.

Work (w) and q (heat) are path functions The amount of work done or heat released

depends on how the system changes states.

Enthalpies of Formation – Standard Reaction Enthalpies

The enthalpy change for the reaction

Hrxn = H(products) - H(reactants)

We cannot measure the absolute values of the enthalpies!!

How do we ‘measure’ enthalpies (or heat contents) of chemical species?

The Formation Reaction

A "chemical thermodynamic reference point."

For CO and CO2

C (s) + O2 (g) CO2 (g)

C (s) + ½ O2 (g) CO (g)

The "formation" of CO and CO2 from its constituent elements in their standard states under standard conditions.

The Formation Reaction

The formation reaction For the formation of 1.00 mole of

Na2SO3(s)

2 Na(s) + S(s) + 3/2 O2 (g) Na2SO3 (s) The ‘formation enthalpy of Na2SO3(s)’,

symbolised Hf[Na2SO3 (s)]

Standard Conditions for Thermodynamic Reactions

The degree sign, either or , indicates standard conditionsP = 1.00 atm[aqueous species] = 1.00 mol/LT = temperature of interest (note 25C or

298.15 K is used in the tables in your text).

The Significance of the Formation EnthalpyHf° is a measurable quantity!

Compare CO (g) with CO2 (g)

C (s) + 1/2 O2 (g) CO (g)

Hf° [CO(g)] = -110.5 kJ/mole

C (s) + O2 (g) CO2 (g)

Hf° [CO2(g)] = - 393.5 kJ/mole

The formation enthalpy for CO2(g) is larger than the formation enthalpy of CO (g).

Reaction Enthalpies

Formation enthalpies – thermodynamic reference point,

Formation of the elements from themselves is a null reaction – Hf (elements) = 0 kJ / mole.

The Combustion of Propane

The General Equation

Calculate enthalpy changes from the formation enthalpies as follows.

Reverse a reaction, the sign of the enthalpy change for the reaction is reversed.Multiply a reaction by an integer, the enthalpy change is multiplied by the same integer.

reactants frfpr HnproductsHnH

The Measurement of Energy Changes – CalorimetryCalorimetry – the measurement of heat and

energy changes in chemical and physical processes.

Heat capacity (C) – the amount of heat (energy) needed to raise the temperature of a given mass of substance by 1°C.

Specific heat capacity (s) – the amount of heat energy (in Joules, J) required to raise 1 g of a substance by 1°C (units = J/g °C).

General expression for heat capacity

C = m s m is the mass of the substance (in grams).

Molar heat capacity

Cm = M s M – molar mass of the substance s – its specific heat capacity.

The Calorimeter

A calorimeter – a device which contains water and/or another substance with a known capacity for absorbing energy (heat).

Calorimeters are adiabatic systems. All energy changes take place within

the calorimeter.

Adiabatic System

Adiabatic system – thermally insulated from the rest of the universe

No heat exchange between system and surroundings!

For an adiabatic system,

qtotal = qrxn + qH2O + qcal = 0

-qrxn = qH2O + qcal

The Constant Volume (Bomb) Calorimeter

E = qv

The Constant Pressure Calorimeter

H = qp

Other important Enthalpy changes

Many other important processes have associated enthalpy changes.

The measurement of the heat changes for these process can give us some insight into the changes in intermolecular forces that occur during the transformation.

Heat of dilution and solution.

Hsol = the heat absorbed or given off when a quantity of solute is dissolved in a solvent.

Hsol = H(sol’n) - H(component)H(component) = H (solid) + H(solvent)

For the process,

HCl (aq, 6 M) HCl (aq, 1 M).A significant amount of heat is released

when the acid solution is diluted. This is the enthalpy of dilution of the acid.

Hdil = H(sol’n 2) – H(sol’n ,1)

Lattice Enthalpies

Look at the following process. NaCl (s) Na+ (g) + Cl- (g)

H = Hlat = 788 kJ/mole the lattice enthalpy

A very endothermic reaction! Due to the strength of the ionic bond!

Latent Heats

Latent heats are the enthalpy changes associated with phase transitions.

H2O (l) H2O (g)

Hr = Hvap the enthalpy of vapourization.

H2O (s) H2O (l)

Hr = Hfus the enthalpy of fusion.

H2O (s) H2O (g)

Hr = Hsub the enthalpy of sublimation.

Latent Heats

Foods and Fuels

Most of the chemical reactions that produce heat are combustion reactions.

Note – all combustion reactions are exothermic.

Fuel values are generally reported as positive quantities.

Obtaining fuel values – calorimetry.

Fossil Fuels

Coal, petroleum, and natural gas are known as fossil fuels. They are collectively the major source of energy for commercial and personal consumption.

Fossil fuels are mixtures of many different kinds of organic compounds.

The fuel values of fossil fuels is directly related to the amount of carbon and hydrogen in the fuel.

Hydrogen As a Fuel

Hydrogen has a huge fuel value (142 kJ/g).The combustion product is innocuous –

water. Obviously, there are problems!Two major difficulties with H2 as a fuel

source.Where do we get the hydrogen?How do we store the hydrogen?