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Unit 1
Energy Matters
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Reaction Rates
Enthalpy changes
Patterns in the Periodic Table
Bonding, Structure and Properties
The Mole
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Reaction Rates
Energy changes
• Exothermic changes cause heat to
be released to the surroundings.
• Endothermic changes cause
absorption of heat from the
surroundings.
• A potential energy diagram can be
used to show the energy pathway
for a reaction.
Energy Diagrams
• We can represent
what happens in a
chemical reaction
using an energy
diagram.
• The energy of the
reactants and
products is
shown.
Activated complex
E
n
e
r
g
y
Reactants
Products
• The enthalpy
(energy change)
for the reaction
can be calculated
 DH = HR – HP
• For an exothermic
reaction DH is
negative.
Activated complex
E
n
e
r
g
y
Reactants
DH
Products
• The enthalpy
(energy change)
for the reaction
can be calculated
 DH = HR – HP
• For an
endothermic
reaction DH is
positive.
Activated complex
E
n
e
r
g
y
Products
DH
Reactants
• An activated
complex is formed.
• This is an unstable
collection of atoms,
intermediate
between reactants
and products.
Activated complex
E
n
e
r
g
y
Reactants
Products
• The energy which
is needed to
produce the
activated complex
is called the
activation energy
(EA).
E
n
e
r
g
y
EA
• A similar energy
diagram can be
drawn for an
endothermic
reaction.
E
n
e
r
g
y
EA
DH
• A similar energy
diagram can be
drawn for an
endothermic
reaction.
E
n
e
r
g
y
EA
DH
Following the course of a
reaction
• Reactions can be followed by
measuring how some quantity we
can measure changes with time.
• Reactions can be followed by
measuring changes in
concentration, mass or volume of
either the reactants or products
• The average rate of a reaction, or
stage in a reaction, can be
calculated by dividing the
difference between the initial and
final quantities by the time
interval.
Rate = Dchange
Dtime
Volume
of gas
released
(ml)
Time (s)
Volume
of gas
released
(ml)
V
1
t1
Time (s)
V2
Volume
of gas
released
(ml)
V
1
t1
Time (s)
t2
V2
Volume
of gas
released
(ml)
V
Volume change (DV)
DV = V2 – V1
1
t1
Time (s)
t2
V2
Volume
of gas
released
(ml)
V
Time change (Dt)
Dt = t2 – t1
1
t1
Time (s)
t2
V2
Volume
of gas
released
(ml)
V
Average reaction rate
between t1 and t2.
1
Rate = DV/Dt
t1
Time (s)
t2
• The rate of a reaction, or stage in a
reaction, is proportional to the
reciprocal of the time taken.
•
Rate proportional to 1/t
Factors affecting rate
• The rates of reactions are affected
by changes in
• Concentration
• Particle size
• Temperature.
Collision Theory
• Reactions will only take place
when the reacting particles collide.
• Reactions will only take place
when the reacting particles collide.
• The particles need to collide at the
correct angle.
• The particles need to collide at the
correct angle.
• The particles need to collide at the
correct angle.
• Collision theory explains the effect
of concentration on reaction rates.
• The more particles there are in a
given volume, the greater the
chance of collision.
Concentration
Collision Theory
• Collision theory explains the effect
surface area on reaction rates.
• Collisions can only take place on
the surface.
• The larger the surface the more
collisions.
Surface Area
Activation energy
Two molecules approach each other
If they don’t have the required
activation energy nothing happens.
Two molecules approach each other
If they have the required activation
energy the molecules form the
Activated complex
If they have the required activation
energy the molecules form the
Activated complex
The activated complex splits apart
To form the products.
Temperature
• Each molecule has a kinetic
energy.
• Not all molecules in a material
have the same kinetic energy.
• Temperature is a measure of the
average kinetic energy of the
molecules.
Energy Distribution
• Energy distribution diagrams show
the numbers of molecules with
each kinetic energy.
number of
molecules
energy of molecules
• As the temperature increases the
kinetic energy distribution
changes.
number of
molecules
energy of molecules
• At higher temperatures the
average kinetic energy of the
molecules is greater.
• More molecules will be able to
provide the activation energy.
• The reaction will go faster.
Activation energy.
• In some chemical reactions light can be
used to increase the number of particles
with energy greater than the activation
energy (e.g. photography)
• In other reactions shock can increase
the number of particles with energy
greater than the activation energy.
Excess
• When we carry out a chemical
reaction we will usually use more
of one of the reagents than is
needed.
• That reagent is said to be in
excess.
• We can calculate the reagent in
excess using a mole equation.
• 1 g of carbon reacts with 10 g of
copper(II) oxide. Show by calculation
which reagent is in excess.
C
+ 2CuO  2Cu + CO2
1 mole 2 moles 2 moles 1 mole
12g
2x(63.5 + 16)
12g
159g
1g
159/12=13.3g
• 1 g of carbon reacts with 10 g of
copper(II) oxide. Show by calculation
which reagent is in excess.
C
+ 2CuO  2Cu + CO2
1g
13.3g
Since 13.3g of copper(II) oxide are
needed to react with 1g carbon, then
carbon is the reagent in excess in the
example given.
Catalysts
• Catalysts speed up reactions,
without being changed by the
reaction.
• Catalysts are used in many
industrial processes.
• They reduce the temperature
needed, so reducing energy costs.
Industrial Process
Catalyst used
Haber Process
Iron
Ostwald Process
Platinum
Contact Process
Vanadium(V) oxide
Catalytic Cracking
Aluminium oxide
Collision Theory
• Not all collisions are successful,
because they need to have the
appropriate activation energy.
• Activation energy is the energy
needed to start the reaction.
Catalysts
• Catalysts work by
providing an
alternative
pathway for the
reaction.
• This pathway has
a lower activation
energy
E
n
e
r
g
y
• Catalysts work by
providing an
alternative
pathway for the
reaction.
• This pathway has
a lower activation
energy
E
n
e
r
g
y
• Heterogeneous catalysts are in a
different state from the reactants
they catalyse.
• Homogeneous catalysts are in the
same state as the reactants they
catalyse.
• Heterogeneous catalysts work by
the adsorption of reactant
molecules.
• The adsorption of the molecules
loosens bonds and makes it easier
for the substance to react.
• The surface activity of a catalyst
can be reduced by poisoning, when
surface sites are taken over by
other substances, preventing
reactants being adsorbed.
• Impurities in the reactants result
in the industrial catalysts having to
be regenerated or renewed.
Catalytic converters
• Catalytic convertors are fitted to
cars to catalyse the conversion of
poisonous carbon monoxide and
oxides of nitrogen to carbon
dioxide and nitrogen.
• Cars with
catalytic
converters only
use ‘lead-free’
petrol to prevent
poisoning of the
catalyst.
Enzymes
• Enzymes catalyse the chemical
reactions which take place in the
living cells of plants and animals.
• Enzymes are used in many
industrial processes.
Reaction Rates
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Rates.
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Enthalpy changes
Enthalpy changes
• An enthalpy change is the energy
produced or released in a chemical
reaction.
• Enthalpy change is given the
symbol DH.
• Enthalpy change is measured in
kilojoules per mole (Kj/mol)
Enthalpy of combustion
• The enthalpy of combustion of a
substance is the enthalpy change
when one mole of the substance
burns completely in oxygen.
Enthalpy of solution
• The enthalpy of solution of a
substance is the enthalpy change
when one mole of the substance
dissolves in water.
Enthalpy of neutralisation
• The enthalpy of neutralisation of
an acid is the enthalpy change
when the acid is neutralised to
form one mole of water.
Calculating enthalpy
changes
• To calculate enthalpy changes we
use the equation:
DH = cmDT
DH is enthalpy change
c is (4.18 Kj/mol)
m is number of kilograms (litres)
of water
DT is the temperature change
Enthalpy changes
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Patterns in the
Periodic Table
Patterns in the Periodic
Table
• The modern
Periodic Table is
based on the
work of
Mendeleev.
• Mendeleev arranged the known
elements in order of increasing
atomic masses.
• He combined this putting elements
with similar chemical properties in
the same vertical group.
• He left gaps for elements that had
not been discovered at that time.
• Certain physical properties show
trends, which repeat from one
period of the periodic table to the
next.
Density and hardness
• Elements at the left side of the
Periodic Table indicate a gradual
increase in these properties
• From Group 1 elements to Group 4
elements. Then there is a gradual
decrease to low values at Group 7
and Group 0 (mostly gases).
Melting and boiling points
• The melting point and boiling point
of an element gives an indication
of the size of the forces that hold
together the atoms or molecules.
• The higher the melting and boiling
point the stronger the forces.
Boiling points
Melting points
Atomic size
• Covalent radius is
half the distance
between the
centres of two
covalently bonded
atoms.
• We use it as a
measure of
atomic size.
Covalent radius
• We can see that covalent radius
decreases as we move from left to
right across a period.
• This is because the nuclear charge
is increasing.
• Attraction to outer electrons
increases.
• Thus the atom is smaller
• We can see that covalent radius
decreases as we move from top to
bottom down a period.
• This is because atoms have more
shells of electrons.
• Nuclear charge is shielded by the
inner electron shells.
• Thus the atom is bigger
Ionisation enthalpy
• First ionisation enthalpy (energy)
is the energy to remove one mole
of electrons from one mole of free,
gaseous atoms.
X(g)  X+(g) + e
Ionisation enthalpies of
the first 20 elements
Ionisation enthalpy
• If we plot ionisation enthalpy
against atomic number we can see
two things.
• Ionisation enthalpy increases as
we move from left to right across a
period.
• Ionisation enthalpy decreases as
you move down a group.
• As we move across a period from
left to right
Atomic size decreases
Charge on the nucleus increases
The force of attraction on the outer
electrons is greater
More energy is needed to remove
an electron
• As we move down a group
Atomic size increases
Extra layers of electrons help shield the
outer electrons from nuclear attraction.
The force of attraction on the outer
electrons is less
Less energy is needed to remove an
electron
• Second ionisation enthalpy
(energy) is the energy to remove a
second mole of electrons.
X+(g)  X2+(g) + e
• If we look at successive ionisation
enthalpies for the same element,
we see that there is a point where
there is a large increase.
• This comes when the next electron
has to be taken from an inside
shell, nearer the nucleus where the
electron is more strongly held.
Element
1st I.E.
(kJmol-1)
2nd I.E.
(kJmol-1)
3rd I.E.
(kJmol-1)
4th I.E.
(kJmol-1)
Na
502
4560
6920
9540
Mg
744
1460
7750
10500
Al
584
1830
2760
11600
Electronegativity
• Atoms of different
elements have
different
attractions for
bonding
electrons.
• Pauling worked
out values for
electronegativity.
• Electronegativity is a measure of
the attraction an atom involved in
a bond has for the electrons of the
bond.
• Electronegativity values increase
across a period and decrease down
a group.
Patterns in the Periodic
Table
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Periodic Table.
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Bonding, structure
and properties
Bonding
• Bonds are forces which hold
particles together.
• There are two main kinds of forces
• Intramolecular forces
• Intermolecular forces
Intramolecular forces
• Intramolecular forces are the forces
which hold atoms together.
• Covalent bonds are typical
intramolecular forces.
Intermolecular forces
• Intermolecular forces are the forces
which exist between molecules.
• Van der Waal’s forces, dipole-dipole
attractions and hydrogen bonds are
typical intermolecular forces.
Metallic bonding
• Metals can lose their outer shell
electrons to gain a stable electron
arrangement.
• Metal atoms are arranged in a lattice
and can also delocalise their outer
shell electrons, allowing them to
move freely between the atoms in the
lattice.
• This brings the metals closer to
obtaining a stable electron
arrangement.
• Since the metal atoms become
positively charged ions they attract
the free moving electrons in the
lattice.
• This attraction
forms a metallic
bond which is very
strong.
Covalent bonding
• As with all bond formations, the
atoms must first collide with one
another.
• When some atoms collide with each
other, the electrons in the outer shell
can be shared between the atoms.
• The electrons of the two atoms are both
negatively charged and repel each other.
• When a collision takes place with
sufficient energy to form a compound, the
outer energy levels overlap and the atoms
share the electrons.
• The overlap area has an increase in
negative charge, which is strongly
attracted by the positive nuclei of both
atoms.
• This draws the two atoms closely
together.
• The electrostatic force of attraction
between the nuclei and the shared
electrons forms a strong covalent bond.
+
ee-
+
Polar covalent bonding
• Sometimes on atom has a greater
force of attraction than the other.
• This leads to polar covalent
bonding, where there are slight
charges (shown by d+ and d-) on
d+
the atoms. d+
ee-
+
What kind of bonding?
• The type of bond
depends on the
difference in
electronegativity
between the
bonded atoms.
Covalent
0
Borderline
1.5
Ionic
3.0
Ionic bonding
• An ionic bond usually occurs
between a metal and a non-metal and
involves ions, which are charged
atoms (or groups of atoms).
• In ionic bonding, electrons are
transferred from one atom to another
allowing both atoms to achieve a
stable electron arrangement.
• For example, sodium and chlorine
atoms would form an ionic bond
making the compound sodium
chloride as shown below :
van der Waal’s forces
• van der Waals’ forces are forces of
attraction which can operate
between all atoms and molecules.
• van der Waals’ forces are much
weaker than all other types of
bonding.
• Random
movement of
electrons in an
atom can cause a
temporary
imbalance in
charge.
• The force of
attraction
between two such
atoms is called
van der Waal’s
force.
• Thus van der Waals’ forces are a
result of electrostatic attraction
between temporary dipoles,
caused by movement of electrons
in atoms and molecules.
• The strength of van der Waals’
forces is related to the size of the
atoms or molecules.
Polar molecules
• A molecule is
described as polar
if it has a
permanent dipole
e.g. hydrogen
chloride
d+
H
Cl
d-
Cl
• A molecule with
polar bonds will
not be polar if it is
arranged
symmetrically e.g.
Cl
C
Cl
Cl
tetrachloromethane
non-polar
Cl
Cl
C
Cl
H
trichloromethane
polar
• The dipoles of
adjacent
molecules will
attract each
other, as shown
opposite.
• Permanent dipole-permanent
dipole interactions are additional
electrostatic forces of attraction
between polar molecules.
• Permanent dipole-permanent
dipole interactions are stronger
than van der Waals’ forces for
molecules of equivalent size.
• The melting and boiling points of
polar substances are higher than
the melting and boiling points of
non-polar substances with similar
molecular sizes.
• This is due to the extra forces
between molecules.
Solvents
• Ionic compounds and polar molecular
compounds tend to be soluble in polar
solvents such as water and insoluble in
non-polar solvents.
• Non-polar molecular substances tend to
be soluble in non-polar solvents and
insoluble in polar solvents.
Hydrogen bonds
• Hydrogen bonds
are electrostatic
forces of
attraction
between
molecules which
contain these
highly polar
bonds.
• A hydrogen bond is stronger than other
forms of permanent dipole-permanent
dipole interaction but weaker than a
covalent bond.
• Bonds consisting of a hydrogen atom
bonded to an atom of a strongly
electronegative element such as
fluorine, oxygen or nitrogen are highly
polar.
• The anomalous boiling points of
ammonia, water and hydrogen
fluoride are a result of hydrogen
bonding.
• Boiling points, melting points,
viscosity and miscibility in water
are properties of substances which
are affected by hydrogen bonding.
• Hydrogen bonding between
molecules in ice results in an
expanded structure which causes
the density of ice to be less than
that of water at low temperatures.
Structure
• Different materials have different
structures
• Metallic
• Ionic
• Covalent molecular
• Covalent network
• Monatomic
Metallic Structure
• A metallic
structure consists
of a giant lattice
of positively
charged ions and
delocalised outer
electrons.
Ionic Structure
• An ionic structure
consists of a giant
lattice of
oppositely
charged ions.
• This structure
only is found in
compounds.
Covalent Molecular
Structure
• A covalent molecular
structure consists of
discrete molecules
held together by
weak intermolecular
forces.
• This can be found in
elements (Cl2) or
compounds (CH4)
Covalent Network
Structure
• A covalent network
structure consists of
a giant lattice of
covalently bonded
atoms.
• This structure is
found in elements
(C(diamond)) or
compounds (SiO2 or
SiC)
Monatomic Structure
• A monatomic
structure consists
of discrete atoms
held together by
van der Waals’
forces.
• This is found in
the Noble Gases.
Structure
• The first 20 elements in the
Periodic Table can be categorised
according to bonding and
structure.
Li, Be, Na, Mg, Al, K, Ca have metallic structure
H2, N2, O2, F2, Cl2, P4, S8 and C (fullerenes)
have covalent molecular structure
B, C (diamond, graphite), Si)
have covalent network structure
He, Ne and Ar have monatomic structure
Bonding, structure and
properties
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structure and properties
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The mole
The Mole
• One mole of any substance
contains the formula mass in
grams.
mass( g )
moles 
gram formula mass
• A 1 mole per litre (mol/l) solution
contains one mole in each litre of
solution.
number of moles
concentration 
volume(l )
• Each material is made up of
different types of particles: atoms,
molecules, ions.
• The basic blocks of each substance
are called formula units.
Formula Units
Substance
Example
Atomic
element
Molecular
element
Covalent
compound
Ionic
compound
copper
Formula
Units
Cu atoms
chlorine
Cl2 molecules
methane
CH4
molecules
1xCa2+ ions
2xCl- ions
calcium
chloride
The Avogadro Constant
• One mole of any substance
contains 6.02×1023 formula units.
• This number is called the Avogadro
Constant
• Equimolar amounts of substances
contain equal numbers of formula
units.
•
•
•
•
How many atoms in 2g of calcium?
Calcium is Ca, formula mass 40.
2g is 2/40 = 0.05 moles
Number of atoms = 0.05 x 6 x 1023
= 3 x 1022
• How many ions in 10 of calcium
carbonate?
• Calcium carbonate is CaCO3, formula
mass 100.
• 10g is 10/100 = 0.1 moles
• Number of formula units
= 0.1x6x1023 = 6x1022
• Each formula unit has 2 ions.
Number of ions = 2 x 6x1022 = 1.2x1023
Molar Volume
• The molar volume (in units of l
mol-1) is the same for all gases at
the same temperature and
pressure.
• This is because the sizes of
molecules is insignificant
compared to the distances
between them.
• The volume of a gas can be
calculated from the number of
moles and vice versa.
• How many atoms are in 2.5 litres
of hydrogen gas? (molar volume =
25l)
• 2.5 litres =2.5/25 = 0.1 moles
• Number of H2 molecules =
0.1x6x1023 =6x1022
• Number of atoms =2x6x1022
=1.2x1023
Reacting Volumes
• Since equal volumes of any gas
contain equal numbers of
molecules then we can relate mole
numbers to molecules.
CH4 + 2O2  CO2 + 2H20
1 mole 2 moles 1mole 2 moles
1 vol
2 vols
1 vol
2 vols
Reacting Volumes
• What volume of oxygen will be
needed to burn 20 ml of propane?
C3H8 + 5O2  3CO2 + 4H20
1 mole 5 moles 3 moles 4 moles
1 vol
5 vols
3 vol
4 vols
20 ml 100 ml
The Mole
• The mole is the key to all our
calculations.
• The next slide shows how the mole
is related to all the other quantities
we need.
• It also includes a quantity we will
not meet until later.
1 litre of
1 mol/l solution
1 gram
formula
mass
1 mole
6 x 1023
formula units
1 molar
volume
n x 96500 coulombs
The Mole
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