chemical reactions - PowerPoint

Document Sample
chemical reactions - PowerPoint Powered By Docstoc
The study of heat released or required by chemical reactions
Fuel is burnt to produce energy - combustion (e.g. when fossil fuels are burnt) CH4(g) + 2O2(g) CO2(g) + 2H2O(l) + energy

What is Energy? Energy Kinetic energy (EK) Potential energy (EP)

Energy due to motion

Energy due to position (stored energy)

Total Energy = E

Kinetic Energy = EK +

+ EP

Potential Energy

Kinetic energy & potential energy are interchangeable

Ball thrown upwards slows & loses kinetic energy but gains potential energy

The reverse happens as it falls back to the ground

Law of Conservation of Energy: the total energy of the universe is constant and can neither be created nor destroyed; it can only be transformed.
The internal energy, U, of a sample is the sum of all the kinetic and potential energies of all the atoms and molecules in a sample i.e. it is the total energy of all the atoms and molecules in a sample

Systems & Surroundings
In thermodynamics, the world is divided into a system and its surroundings

A system is the part of the world we want to study (e.g. a reaction mixture in a flask)
The surroundings consist of everything else outside the system SYSTEM



OPEN SYSTEM: can exchange both matter and energy with the surroundings (e.g. open reaction flask, rocket engine) CLOSED SYSTEM: can exchange only energy with the surroundings (matter remains fixed) e.g. a sealed reaction flask

ISOLATED SYSTEM: can exchange neither energy nor matter with its surroundings (e.g. a thermos flask)

HEAT is the energy that transfers from one object to another when the two things are at different temperatures and in some kind of contact e.g. kettle heats on a gas flame cup of tea cools down (loses energy as heat) Thermal motion (random molecular motion) is increased by heat energy

i.e. heat stimulates thermal motion

Work is the transfer of energy that takes place when an object is moved against an opposing force i.e. a system does work when it expands against an external pressure Car engine: petrol burns & produces gases which push out pistons in the engine and transfer energy to the wheels of car •Work stimulates uniform motion • Heat and work can be considered as energy in transit

S.I. unit of energy is the joule (J)

Heat and work ( energy in transit) also measured in joules
1 kJ (kilojoule) = 103 J Calorie (cal): 1 cal is the energy needed to raise the temperature of 1g of water by 1oC 1 cal = 4.184 J

Internal energy changes when energy enters or leaves a system

U = Ufinal - Uinitial U change in the internal energy

Heat and work are 2 equivalent ways of changing the internal energy of a system

Change in internal energy


Energy supplied to system as heat


Energy supplied to system as work

U = q (heat) + w (work)






U like reserves of a bank: bank accepts deposits or withdrawals in two currencies (q & w) but stores them as common fund, U.

First Law of Thermodynamics:
the internal energy of an isolated system is constant

Signs (+/-) will tell you if energy is entering or leaving a system + indicates energy enters a system - indicates energy leaves a system

WORK •An important form of work is EXPANSION WORK i.e. the work done when a system changes size and pushes against an external force e.g. the work done by hot gases in an engine as they push back the pistons

In a system that can’t expand, no work is done (w = 0) U = q + w when w = 0, U = q (at constant volume)

•A change in internal energy can be identified with the heat supplied at constant volume

ENTHALPY (H) (comes from Greek for “heat inside”) • the change in internal energy is not equal to the heat supplied when the system is free to change its volume • some of the energy can return to the surroundings as expansion work

 U < q

The heat supplied is equal to the change in another thermodynamic property called enthalpy (H)

i.e. H = q
• this relation is only valid at constant pressure

As most reactions in chemistry take place at constant pressure we can say that: A change in enthalpy = heat supplied

EXOTHERMIC & ENDOTHERMIC REACTIONS Exothermic process: a change (e.g. a chemical reaction) that releases heat. A release of heat corresponds to a decrease in enthalpy Exothermic process: H < 0 (at constant pressure) Burning fossil fuels is an exothermic reaction

Endothermic process: a change (e.g. a chemical reaction) that requires (or absorbs) heat.
An input of heat corresponds to an increase in enthalpy Endothermic process: H > 0 (at constant pressure) Forming Na+ and Cl- ions from NaCl is an endothermic process

Photosynthesis is an endothermic reaction (requires energy input from sun)

Measuring Heat
Exothermic reaction, heat given off & temperature of water rises



Endothermic reaction, heat taken in & temperature of water drops

How do we relate change in temp. to the energy transferred? Heat capacity (J/oC) = heat supplied (J)

temperature (oC)
Heat Capacity = heat required to raise temp. of an object by 1oC • more heat is required to raise the temp. of a large sample of a substance by 1oC than is needed for a smaller sample

Specific heat capacity is the quantity of energy required to change the temperature of a 1g sample of something by 1oC

Specific Heat Capacity (Cs)

Heat capacity


J / oC / g

J / oC


Vaporisation Energy has to be supplied to a liquid to enable it to overcome forces that hold molecules together • endothermic process (H positive)

Energy is supplied to a solid to enable it to vibrate more vigorously until molecules can move past each other and flow as a liquid

• endothermic process (H positive) Freezing
Liquid releases energy and allows molecules to settle into a lower energy state and form a solid • exothermic process (H negative) (we remove heat from water when making ice in freezer)

Reaction Enthalpies All chemical reactions either release or absorb heat

Exothermic reactions:
Reactants products + energy as heat (H -ve) e.g. burning fossil fuels

Endothermic reactions: Reactants + energy as heat e.g. photosynthesis products (H +ve)

Bond Strengths

Bond strengths measured by bond enthalpy HB (+ve values)
• bond breaking requires energy (+ve H) • bond making releases energy (-ve H) Lattice Enthalpy

A measure of the attraction between ions (the enthalpy change when a solid is broken up into a gas of its ions)
• all lattice enthalpies are positive • I.e. energy is required o break up solids

Enthalpy of hydration Hhyd • the enthalpy change accompanying the hydration of gasphase ions •Na+ (g) + Cl- (g) Na+ (aq) + Cl- (aq) • -ve H values (favourable interaction)

• If dissolves and solution heats up : exothermic •If dissolves and solution cools down: endothermic

Breaking solid into ions


Ions associating with water



Lattice Enthalpy


Enthalpy of Hydration


Enthalpy of Solution

Substances dissolve because energy and matter tend to disperse (spread out in disorder) 2nd law of Thermodynamics

Second Law of Thermodynamics:
the disorder (or entropy) of a system tends to increase
ENTROPY (S) •Entropy is a measure of disorder

• Low entropy (S) = low disorder
•High entropy (S) = greater disorder • hot metal block tends to cool • gas spreads out as much as possible

Total entropy change


entropy change + of system

entropy change of surroundings

disorder of solution

disorder of surroundings

• must be an overall increase in disorder for dissolving to occur

1. If we freeze water, disorder of the water molecules decreases , entropy decreases ( -ve S , -ve H)

2. If we boil water, disorder of the water molecules increases , entropy increases (vapour is highly disordered state)
( +ve S , +ve H)

A spontaneous change is a change that has a tendency to occur without been driven by an external influence e.g. the cooling of a hot metal block to the temperature of its surroundings

A non-spontaneous change is a change that occurs only when driven e.g. forcing electric current through a metal block to heat it

•A chemical reaction is spontaneous if it is accompanied by an increase in the total entropy of the system and the surroundings

• Spontaneous exothermic reactions are common (e.g. hot metal block spontaneously cooling) because they release heat that increases the entropy of the surroundings. •Endothermic reactions are spontaneous only when the entropy of the system increases enough to overcome the decrease in entropy of the surroundings

System in Dynamic Equilibrium
A + B C + D

Dynamic (coming and going), equilibrium (no net change) • no overall change in disorder  S  0 (zero entropy change)

Shared By: