Reaction Rates
• During the course of a chemical reaction, reactants are
being converted into products.
• Measurement of the rate of reaction involves measuring
the ‘change in the amount’ of a reactant or product in a
certain time.
• The rate of reaction changes as it progresses, being
relatively fast at the start and slowing towards the end.
• What is being measured is the average rate over the time
interval chosen.
• Reactions can be followed by measuring changes in
concentration, mass and volume.
Where property = mass/volume/concentration
The above is used when there is no change in
mass/volume/concentration measured, for example
during a colour change reaction.
Collision Theory
A chemical reaction can only occur if there is a
successful collision between reactant molecules.
From national 5 we know that we can speed up a
chemical reaction by;
1.
2.
3.
4.
Decreasing particle size (increasing surface area)
Increasing concentration (of reactant)
Increasing temperature
Adding a catalyst
Collision Theory – Particle Size
• The smaller the particle size, the higher the surface
area.
• The higher the surface area, the greater the number of
collisions that can occur at any one time.
• The greater the number of collisions, the faster the
reaction.
• Therefore the smaller the particle size, the faster the
reaction rate.
Collision Theory – Concentration
• The higher the concentration, the higher the number
of particles.
• The higher the number of particles, the greater the
chance of collisions.
• The greater the number of collisions, the faster the
reaction.
• Therefore the higher the concentration, the faster the
reaction rate.
Collision Theory – Temperature
• The higher the temperature, the higher the energy the particles
have.
• The higher the energy, the faster the particles move.
• The faster the particles move, the greater the chance that they
can collide with sufficient energy (activation energy)
• The greater the number of collisions, the faster the reaction.
• Therefore the higher the temperature, the faster the reaction
rate.
Potential Energy Diagrams
Labels include;
• Exothermic or Endothermic
• Activated complex
• Enthalpy change
• Reaction pathway
• Potential Energy (KJ)
• Activation Energy
Catalysts
• A catalyst is a substance which speeds up
a chemical reaction without getting used
up or changed itself.
• There are two main categories of catalyst;
a) Heterogeneous +
b) Homogenous.
Heterogeneous Catalysts
Heterogeneous catalysts have active sites on their surface.
Reactant molecules form weak bonds with the surface in a
process called adsorption.
At the same time bonds within the adsorbed reactant
molecules are weakened.
The reactant molecules are also held at a favourable angle
for a collision with another reactant molecule to occur.
The product molecules then leave the active site in a stage
called desorption. The active site is then available again.
surface
Active
site
Desorption – product molecules formed.
Unwanted substances can often be adsorbed onto
the active sites thus making them unavailable for
the normal reactants. (Example; lead in petrol.)
When this happens the catalyst is said to be
poisoned.
Sometimes it is not possible to regenerate a
poisoned catalyst and it must be replaced/renewed.
This adds to industry costs so every effort is
made to remove any impurities from reactants that
might poison a catalyst.
Homogeneous Catalysts
A catalyst that is in the same state as the reactants is
said to be a homogeneous catalyst.
The catalyst forms an intermediate compound with one
of the reactants. (this intermediate compound later
decomposes to reform the catalyst.)
For example (using Reactants A and B)
Reactant A + Catalyst
Intermediate
Intermediate + Reactant B
Product + Catalyst
Collision Theory – Catalyst
• A catalyst speeds up a chemical reaction by lowering
the activation energy. (i.e. catalyst provides another
‘easier’ route)
• The lower the activation energy, the greater the
chance of successful collisions.
• The more collisions in a period of time, the faster the
reaction rate.
Trends in the Periodic Table
Density (measured in g/cm3)
Across a period (starting from group 1) the
density increases towards the centre (group
4) and then decreases.
Density tends to increase down a group (as
atomic number increases.)
Atomic size
Atomic size (or covalent atomic radius) is half the distance
between the nuclei of two bonded atoms.
Single bond lengths between atoms of different elements can be
found by adding their individual covalent radii.
e.g. the covalent radii of hydrogen and chlorine are 37 and 99 pm
so the bond length in H-Cl is 37 + 99 pm = 136 pm
There are two clear trends in the periodic table;
1.Going across a period covalent radii decreases. This is because…
the nuclear charge increases but the number of electron shells stays the same – i.e.
the outer electrons are held more tightly, making the atom smaller.
2.Going down a group covalent radii increases. This is because…
the number of electron shells increases and therefore the inner (full) electron shells
shield the outer electrons from the nuclear charge. (known as ‘shielding effect’) i.e. the
outer electron are held less tightly, making the atom bigger.
First Ionisation Energy
Definition;
• First ionisation energy is the energy required to remove one
mole of electrons from an element in gaseous state.
(example at top of page 11 in data book – no excuses!)
Trends;
• Across a period… increases due to an increased nuclear
charge holding the outer electrons more closely (smaller
atoms.) This means you need MORE energy to remove a mole
of electrons.
• Down a group… decreases due to the shielding effect
(bigger atoms.) This means the outer electrons are further
away from the nucleus and therefore this attraction is less,
thus it is easier (less energy) to remove one mole of
electrons.
Other types of ionisation questions;
1. Calculation (whiteboard examples)
2. The first ionisation energy of lithium = 520kJmol-1
but its second ionisation energy value = 7298kJmol-1.
Why is there such a big difference between the two
values?
Lithium achieves a stable outer electron shell (octet)
when it loses one electron. Therefore first ionisation
energy is small. The second electron would therefore
be removed from an stable octet which is
unfavourable and requires a lot of energy.
3. The first compounds of the noble gases were formed
using Xenon (Xe).
Tetrafluoroammonium octafluoroxenate (NF4)2XeF8) is a
strong oxidiser and can be used in propelling rockets.
Suggest why xenon compounds were the first to be made.
(Hint it’s related to atomic size…)
Xenon has the largest atoms of any noble gas. This means
that the electron removed in ionisation to create Xe+ is
the furthest from the nucleus. Hence, the first
ionisation energy of Xenon is the lowest of any of the
noble gases.
Electronegativity
Electronegativity is a measure of the tendency of an
atom to attract electrons. (think pulling power)
Electronegativity is measured on the Pauling scale.
Trends;
Electronegativity increases across a period.
Electronegativity decreases down a group.
(Hint – Fluorine is the highest value)
Melting and Boiling Points
Generally, the stronger the bond between atoms, the higher the energy
required to break that bond. Melting/boiling points are varied and don't
generally form a trend across a period however;
• Metals generally possess a high melting point.
• Most non-metals possess low melting points.
• Group 4 have the highest values.
Metal group example;
In group 1 (alkali metals), the melting/boiling pts decrease as the atomic
number increases. This is because there is an decrease in the attraction
between the particles. (refer to bonding)
Non metal group example;
In group 7 (halogens) the melting/boiling pts increase as the atomic
number increases. This is because there is an increase in the attraction
between the particles. (refer to bonding)
Summary of Bonding types in first 20 elements.
• Inter – means in between.
In other words an INTERmolecular bonds means bonds in
between the molecules.
• Intra – means within.
In other words an INTRAmolecular bond means bonds
within the molecule.
Types of Bonding in elements
There are 3 types;
1.
Metallic Bonding – (intramolecular)
2. Covalent Bonding – (intramolecular)
3. Van der Waal’s
a) London Forces – (intermolecular)
Metallic Bonding
• Metallic bonding only appears in metal elements.
• Metallic bonding occurs between (positively
charged) metal ions and delocalised outer shell
electrons.
• ‘delocalised’ means the electrons are common to all
of the ions (i.e. they move from one to another)
• The movement of delocalised electrons allow metal elements
to conduct electricity
Covalent Molecular Bonding
• Covalent bonding (usually) occurs between two
non metal atoms.
• Covalent bonds are held together through the
attraction between the positively charged
nucleus of one atom and the negatively
charged outer electrons of the other atom.
• Outer electrons are shared in covalent
bonding.
Van der Waals’ Forces
• Van der Waals forces are weak bonding which occurs BETWEEN
molecules.
• London forces are one of three types of Van der Waals’ forces.
• London forces are temporary dipole to temporary dipole
attractions.
• Temporary dipoles occur when electrons lie slightly closer to one
atom than the other. This means for a short time one of the
atoms is slightly negative and the other is slightly positive (i.e.
electrons not shared equally)
• Van der Waals forces are useful when explaining patterns in the
periodic table e.g. melting/boiling points.
Bonding in Specific Groups
Groups 1, 2 and 3
All elements in groups 1, 2 and 3 have strong metallic
bonds holding them together in a giant lattice structure.
Metallic bonds allows metals to be shaped (i.e. malleable
and ductile)
Metals have high melting/boiling pts due to strong
metallic bonds
Boron is the only exception as it has very complex
bonding. B12 is almost as hard as diamond. This suggests a
covalent network structure.
The 3 Structures of Carbon (group 4)
• Each atom covalently bonds
to 4 other atoms.
• This means covalent bonds must
be broken to melt/ boil = very
high m.pt/b.pt values.
• No free electrons = no conduction.
• Tunnels between atoms allow
light through = transparent
structure.
Each atom forms 3 covalent bonds and its
last valence electron becomes delocalised.
As the delocalised electrons are only held
weakly they can flow i.e. graphite
conducts electricity.
The delocalised orbitals sit between the
layers - as a result there are 3 strong
covalent bonds WITHIN the layers but
only weak interaction BETWEEN the
layers.
Due to these weak interactions, graphite
is flaky as the layers can be easily
separated.
Graphite layers are offset (i.e. not above
each other) - light can’t travel through it
meaning it is not transparent.
3. Buckminsterfullerene (aka ‘Bucky Ball’)
• The fullerenes, despite being
large molecules, are discrete
covalent molecules.
• The smallest of fullerenes is a
molecule known as
Buckminsterfullerene (C60).
• This is a spherical molecule
containing 5 and 6 membered
carbon rings.
• The properties are still being
researched so the full
applications are still unknown.
Group 5 (nitrogen and phosphorus)
• Nitrogen atoms form diatomic molecules with
a triple covalent bond.
• This means that nitrogen only has London
forces between the molecules.
• London forces are easily broken and as a
result nitrogen has a low boiling pt. This is
why nitrogen is a gas at room temperature.
• Phosphorus forms tetrahedral P4 molecules
which are larger than N2 molecules. As a
result it has stronger London forces between
its molecules. This stronger attraction means
phosphorus has a higher boiling pt and is a
solid at room temp.
Group 6 (oxygen and sulphur)
• Oxygen atoms form diatomic molecules
with a double covalent bond.
• This means that oxygen only has London
forces
interaction
between
the
molecules.
• London forces are easily broken and as
a result oxygen has a low boiling pt.
Hence oxygen is gas @ room temp.
• Sulphur forms 8 membered rings. The
London forces between the molecules are
strong enough in sulphur to make it solid
at room temperature.
Group 7 (the halogens)
The halogens form diatomic molecules (i.e. they bond with
themselves)
As with oxygen and nitrogen this results in the halogens only
having London forces between each other molecule.
Fluorine and chlorine are very volatile (and therefore
reactive) gases due to these weak intermolecular forces.
Group 8 (the noble gases)
As the noble gases have a stable outer electron shell they do
not form bonds. As a result they remain monoatomic.
Explaining the Melting and Boiling pts Trend
In small discrete covalent molecules the melting and boiling
points are low.
This is because only weak intermolecular London forces
have to be overcome when boiling or melting. The strong
covalent bonds are left unaffected.
In the covalent network solids (carbon, silicon and boron)
strong covalent bonds MUST be broken when melting or
boiling. Breaking these bonds requires a lot more energy and
therefore we get very high values.
In the metal groups (1, 2, 3) strong metallic bonds MUST be
overcome thus they have high melting/boiling points.
Summary of Bonding types in first 20 elements.
Except Buckminsterfullerenes !
Other small
molecules
Metallic
bonding
Covalent
network
Diatomic
molecules
Discrete covalent
molecules
Monatomic
elements
Types of Bonding in compounds
There are 4 main types;
1.
Ionic Bonding – intramolecular
2. Covalent Network Bonding – intramolecular
3. Polar/Non Polar Covalent Bonding - intramolecular
4. Intermolecular (Van der Waals)
Strongest to
Weakest
1. Hydrogen bonding (present in H2O, NH3 and HF)
2. Permanent dipole – Permanent dipole interactions
3. London forces – Temporary dipole interactions
Ionic Bonding
• Ionic bonding is an electrostatic attraction between the
positive ions and negative ions.
• Ionic bonding is related to the electronegativities of elements.
The greater the difference in ‘e.n’ the less likely the elements
are to share outer electrons. (electronegativity definition and
trends are found in previous section of jotter.)
• Instead the element with the higher ‘e.n’ value will gain the
electrons to form a negative ion and the element with lower
‘e.n’ value will lose the electrons to form a positive ion.
• Due to the trends of electronegativity, the elements that are
far apart from one another in the periodic table form ionic
bonds. (normally metal and non metal.) Caesium fluoride is the
compound with the greatest ionic character.
Structure
• Ionic compounds do not form molecules. Instead the positive and
negative ions come together to form lattice structures.
• When the lattice forms, energy is released. This is known as
lattice energy or enthalpy.
• The overall charge of the lattice must be zero and therefore this
affects the number of each ions we have present.
• In sodium chloride (NaCl) there is an equal number of Na+ and Clions.
• In calcium fluoride (CaF2) there are twice as many F- ions than Ca2+
Covalent Compounds
There are 3 types of covalent bonding in
compounds (all involving combinations of non
metals) ;
1. Covalent network structures.
2. Polar covalent molecules
3. Non Polar covalent molecules
Covalent Network
• These covalent network compounds have the
same properties as covalent network elements.
• Both SiC and SiO2 have very high melting pts.
as melting requires breaking strong covalent
bonds.
• Silicon carbide – known as ‘carborundum’ - is
structurally similar to diamond and has many
uses due to it’s strength, durability and low
cost.
Polar Covalent Bonding
Most covalent compounds are made from atoms with slightly different
electronegativity (e.n.)
This difference is not significant enough for one of the atoms to fully
remove an electron from the other. (Approx difference of between 0.5
and 1.6)
As a result the atom (element) with the higher e.n. holds the electron
slightly closer to itself and therefore becomes slightly negative (∂-)
The atom with the lesser e.n. is therefore slightly positive (∂+) as the
electron is sitting further away from it.
Covalent bonds with unequal electron sharing are called polar covalent
bonds.
Molecules are attracted via permanent to permanent dipole interactions
Non Polar Covalent Bonding
• Non polar (or pure) covalent bonding normally occurs when;
1.
electrons are equally shared between the two different
atoms. i.e. equal electronegativity.
E.g. Phosphorus Hydride
2.
the compound structure is symmetrical and therefore
charges are overall balanced.
E.g. CH4 (methane) and CO2 (carbon dioxide)
• Molecules are therefore attracted via weak London forces.
(i.e. temporary to temporary dipole attraction)
Both polar and non polar?
It’s possible for non polar covalent molecules
to have individual polar bonds.
For example; Carbon Tetrachloride (CCl4)
Summary of electronegativity values + bonding
In general;
1.
If the electronegativity difference (usually
called ΔEN) is less than 0.5, then the bond is
non polar (pure) covalent.
2. If the ΔEN is between 0.5 and 1.6, the bond is
considered polar covalent
3. If the ΔEN is greater than 2.0, then the bond
is ionic.
Properties of Polar/ Non Polar Covalent Bonds
Boiling Points
Polar covalent molecules have higher boiling points than
non polar covalent molecules with a similar mass.
This is because the intermolecular forces are stronger
(changing from London forces to permanent – permanent
dipole interactions.)
Permanent dipole to dipole interaction is caused via the
constant attraction between the ∂+ atoms and ∂- atoms
of neighbouring molecules.
Solubility
‘like substances dissolve in like
substances’
This means that polar molecules will dissolve in polar solutions
but not in non polar solutions and vice versa.
This is due to the attraction between ∂+ and ∂- atoms of the
water and the polar substance.
Ionic compounds dissolve in polar solutions in a similar way due
to the interaction between the ions and the ∂+ and ∂- atoms.
Ions surrounded by a layer of water molecules – held by
electrostatic attraction – are said to be hydrated.
Behaviour in electric field
Copy figure 4.8 on page 49
Viscosity (thickness/ability to pour)
Summarise textbook notes on page 54
(diagram + 2/3 sentences max)
Miscibility (ability to mix)
Summarise textbook notes on page 57
(definition + 2/3 sentences max)
Boiling pts of some hydrides (elements bonded to hydrogen)
gfm = 18
gfm = 20
gfm = 17
gfm = 123
The above bonds are very polar due to the large
difference in the electronegativity values.
This interaction is called hydrogen bonding.
Hydrogen bonding occurs in any molecule that contains
any of the above bonds but mainly in hydrogen fluoride
(HF), water (H2O) and ammonia (NH3). These
intermolecular forces affect the properties of these
compounds.
Hydrogen bonding is stronger than both London forces and
permanent dipole to dipole interactions but weaker than
covalent bonding.
Density of Water
Water is unusual as it’s solid form (ice) is less dense than it’s
liquid form (water). This means that ice floats on water whereas
most other solids sink in their own liquid forms. This
phenomenon is due to the structure and bonding which takes
place between the water molecules.
As water molecules cool, they contract. However, at 4oC they
begin to expand. This is because of hydrogen bonds between the
water molecules. This decreases the density of ice (greater
volume/same mass) compared to that of the liquid water.
Ice floating is vital to ‘real life’ – i.e. fish/marine life surviving
under frozen lakes etc (https://www.youtube.com/watch?v=T4GCShGvw-M)