THERMOCHEMISTRY

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THERMOCHEMISTRY
The study of heat released or required by
chemical reactions
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What is Energy?
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Total Energy Kinetic Energy Potential
Energy E EK EP
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
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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
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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
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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)
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HEAT and WORK
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
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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

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UNITS OF ENERGY
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
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INTERNAL ENERGY (U)
Internal energy changes when energy enters or
leaves a system
Heat and work are 2 equivalent ways of changing
the internal energy of a system
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?U q (heat) w (work)
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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
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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

HEAT
In a system that cant expand, no work is done (w
0) ?U q w when w 0, ?U q (at
constant volume)
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• 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 lt q

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• 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
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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 lt 0 (at constant pressure)
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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 gt 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)
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Measuring Heat
Exothermic reaction, heat given off temperature
of water rises
Endothermic reaction, heat taken in temperature
of water drops
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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

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Specific heat capacity is the quantity of energy
required to change the temperature of a 1g sample
of something by 1oC
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• Vaporisation
• Energy has to be supplied to a liquid to enable
  it to overcome forces that hold molecules
  together
• endothermic process (?H positive)

• Melting
• 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)

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Reaction Enthalpies
All chemical reactions either release or absorb
heat
Exothermic reactions Reactants
products energy as heat (?H -ve)
e.g. burning fossil fuels
e.g. photosynthesis
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• 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

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Enthalpy of hydration ?Hhyd

• the enthalpy change accompanying the hydration
  of gas-phase ions
• Na (g) Cl- (g) Na (aq)
  Cl- (aq)
• -ve ?H values (favourable interaction)

WHY DO THINGS DISSOLVE?

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

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Ions associating with water
Breaking solid into ions


Dissolving
Substances dissolve because energy and matter
tend to disperse (spread out in disorder)
2nd law of Thermodynamics
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Second Law of Thermodynamics the disorder (or
entropy) of a system tends to increase

• hot metal block tends to cool
• gas spreads out as much as possible

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Total entropy change
entropy change of system
entropy change of surroundings

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

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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)
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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
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• 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

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System in Dynamic Equilibrium

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