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Duncanrig Secondary School cfe Higher Chemistry Chemical Changes and Structures
Page 1 of 53
Duncanrig Secondary School
CfE Higher Chemistry
Unit 1
Chemical Changes &
Structure
Part 1 Controlling the Rate
Part 2 Trends in the Periodic Table
Part 3 Structure and Bonding
Duncanrig Secondary School cfe Higher Chemistry Chemical Changes and Structures
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Circle a face to show how much understanding you have of each statement: if you fully
understand enough to do what the outcome says, if you have some understanding of the
statement, and if you do not yet understand enough to do what the statement says. Once
you have completed this, you will be able to tell which parts of the topic that you need to
revise, by either looking at your notes again or by asking for an explanation from your teacher
or classmates.
Learning Outcomes – Controlling the Rate
By the end of this topic I will be able to:
1 Explain the effect of temperature, concentration and particle
size in terms of the energy and number of collisions (Collision
Theory).
2 State which reactions are slowest or fastest at different
points using the slope of rates graphs.
3 State that activation energy is the minimum energy required for
particles to react.
4 Draw a graph showing the effect of temperature on the kinetic
energy of particles .
5 Use activation energy on this graph to explain why higher
temperatures speed up reactions.
6 State that catalysts speed up reactions by providing an
alternative reaction pathway with lower activation energy.
7 Describe the difference between a homogeneous catalyst and a
heterogeneous catalyst.
8 Explain the adsorption, reaction and desorption stages in the
action of a heterogeneous catalyst.
9 State that catalyst poisons occupy the active site in a catalyst
and prevent it working.
10 State that enzymes are biological catalysts and give examples
of some enzymes.
11 Explain why enzymes operate at optimum temperatures and pH
values.
12 Draw potential energy diagrams for exothermic and
endothermic reactions.
13 State that enthalpy change represents the difference: ∆H =
H(products) – H(reactants).
14 State that the activated complex is an unstable arrangement of
atoms formed at the maximum of the potential energy barrier,
during a reaction.
15 Use potential energy diagrams to illustrate the effect catalysts
have on the activation energy and reaction pathway.
Duncanrig Secondary School cfe Higher Chemistry Chemical Changes and Structures
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Learning Outcomes – Trends in the Periodic Table
By the end of this topic I will be able to:
1 Define the density of an element as its mass per unit volume,
usually in gcm-3.
2 Define the covalent radius as a measure of the size of an atom
(specifically that it is half the distance between the nuclei of
two bonded atoms of an element).
3 State that the atomic size decreases across a period and
increases down a group.
4 Explain why there are changes in atomic size across a period and
down a group.
5 Define the first ionisation energy as the energy required to
remove one mole of electrons from one mole of gaseous atoms
6 Understand that the second and subsequent ionisation energies
refer to the energies required to remove further moles of
electrons.
7 Explain the trends in first ionisation energy across periods and
down groups in terms of atomic size, nuclear charge and the
screening effect due to inner shell electrons
8 Understand that atoms of different elements have different
attractions for bonding electrons.
9 Define electronegativity as a measure of the attraction an atom
involved in a bond has for the electrons of the bond.
10 State that electronegativity values increase across a period and
decrease down a group.
11 Explain the trends in electronegativity across periods and down
groups in terms of nuclear charge, covalent radius and the
presence of “screening” inner shell electrons.
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Learning Outcomes – Bonding and Structure
By the end of this topic I will be able to:
1
The bonding types of the first twenty elements; metallic (Li, Be,
Na, Mg, Al, K and Ca); covalent molecular (H2, N2, O2, F2, Cl2, P4,
S8 and C60 [fullerenes]); covalent network (B, C (diamond,
graphite), Si) and monatomic (noble gases)
2 Describe the bonding continuum moving from pure non-polar
covalent to ionic.
3 Explain how polar covalent bonds arise
4 Explain how van der Waals forces arise between molecules.
5 Describe what causes dispersion forces to exist between
gaseous atoms and molecules.
6 Explain how the polarity of molecules affects the strength of
dispersion forces.
7 Explain why certain molecules have a stronger type of van der
Waal force called a hydrogen bond
8 Explain how the properties of substances are affected by the
type of bonding that they exhibit..
9 Predict the solubility of a substance from information about
solute and solvent polarities.
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PART 1 CONTROLLING THE RATE
The Rate of Chemical Reactions
Everyday reactions have different speeds; some are over in a fraction of a second
(fast: like a gas explosion) while others can take years (slow: like the rusting of iron).
Most reactions occur at rates between these two extremes (medium: like a cake
baking).
Collision Theory
For a chemical reaction to occur some important things have to happen:
1. The reacting particles must collide together.
2. Collisions must have sufficient energy to produce a product.
3. The reacting particles must have the correct orientation.
Therefore anything that increases the number of and energy of collisions between
reactant particles will speed up a reaction.
Factors Affecting the Rate of a Reaction
There are three main factors affecting the rate of a chemical reaction:
a) Particle Size:
The smaller the particles, the faster
the reaction. This is because smaller
particles provide more surface area
for collision.
Example – Marble powder reacts faster with acid than marble chips.
b) Concentration:
The higher the concentration, the faster the
reaction. The higher the concentration
of solutions, the more particles you have
crowded into a small volume of liquid.
Hence, the more likely they are to
collide with each other.
Example – 2 mol l-1 hydrochloric acid reacts faster with magnesium ribbon than 1 mol l-1
hydrochloric acid.
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c) Temperature:
Although a higher temperature will cause molecules to move faster, and there may be
more collisions, this is not the main reason why higher temperature increases reaction
rate. The main reason is that more of the collisions which occur will lead to a
successful reaction. This is because at higher temperature, more particles have the
activation energy required for a reaction to happen.
As a rough guide, the rate of reaction doubles when the temperature increase by
10OC.
Example - Benedicts solution reacts faster with glucose solution at 50OC than at
25OC.
Catalysts
Even when particle size is decreased and concentration and temperature are
increased, many chemical reactions are still too slow. How can the rate of these
reactions be increased? This is especially important in today’s competitive market:
companies are constantly trying to produce more cost effective products by increasing
the rate of industrial reactions.
A catalyst is a substance which can be used to increase the rate of a chemical
reaction. The 'amount' of catalyst at the end of the reaction is the same as at the
start, i.e. the catalyst is not used up in the reaction and the catalyst can be
recovered chemically unchanged at the end of reaction. Different reactions require
different catalysts and not all reactions have a suitable catalyst.
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Collision Theory and the Activated Complex
In order to react particles must collide.
A chemical reaction will only occur if the reacting particles collide with enough kinetic
energy. The energy is required to overcome the repulsive forces between the atoms
and molecules and to start the breaking of bonds.
The minimum kinetic energy required for a reaction to occur is called the activation
energy (EA).
When the reactant particles collide with the required activation energy they form an
activated complex. This unstable intermediate breaks down.
E,g, The reaction of hydrogen and bromine
Sometimes the collisions do not result in a reaction, despite having the
minimum kinetic energy.
This is thought to be because the particles have not collided with the
correct geometry (angle) to allow the activated complex to be formed.
In the above reaction of hydrogen and bromine the particles collide side
on but if they collided end on…
H-H + Br-Br H----H-----Br----Br
no reaction occurs as the activated complex cannot be formed if only 2 of
the atoms come into contact with one another.
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Collision Theory and Concentration
The straight line graph means rate is directly proportional to the
concentrations of the reactants, i.e. double the concentration and you
double the rate. This is true of many reactions.
The faster rate is due to the increased number of collisions which must
occur with higher concentrations of reactants.
Collision Theory and Particle Size
The smaller the particle size, the faster the reaction as the total
surface area is larger so more collisions will occur.
Note
The steeper the curve the faster the reaction
The same volume of gas will be produced if the same number of
moles of reactants are used.
Concentration
(mol l-1)
Rate = 1/t
(s-1)
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The Effect of Concentration Changes on Reaction Rate
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Duncanrig Secondary School cfe Higher Chemistry Chemical Changes and Structures
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KINETIC Theory and Temperature
Temperature is a measure of the average kinetic energy of the particles
of a substance.
At any given temperature, the particles of a substance will have a range
of kinetic energies and this can be shown on an energy distribution
graph.
NB The maximum height of T2 is always lower than T1
The graph above shows the kinetic energy distribution of the particles of
a reactant at two different temperatures.
It shows that at the higher temperature (T2), many more molecules have
energies equal to or greater than the activation energy (Ea). This
leads to an increase in the rate of successful collisions and hence reaction
rate.
A small rise in temperature can cause a large increase in the number of
particles having the activation energy and so can result in a large increase
in reaction rates.
For some reactions, a 10oC rise in temperature can double the reaction
rate.
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The Effect of Temperature Changes on Reaction Rate
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Although most chemical reactions follow this pattern there are other
possibilities.
Photochemical Reactions
Photochemical reactions are speeded up by the presence of light.
In these reactions, the light energy helps to supply the activation energy,
i.e. it increases the number of particles with energy equal to or greater
than the activation energy.
Examples of photochemical reactions are:
Photosynthesis
Alkane with bromine water
Chlorine and hydrogen gases
H2(g) + Cl2(g) 2HCl(g)
Catalysts and Reaction Rate
A catalyst is a substance which changes the speed of a chemical reaction
without being permanently changed itself.
Catalysts speed up chemical reactions by providing an alternative
reaction pathway which has a lower activation energy.
Explosive reaction Enzyme controlled reaction Enzyme reaction
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There are 2 main types of catalyst:
Homogeneous Catalysts Heterogeneous Catalysts
Homogeneous catalysts are in the same state as the reactants.
Heterogeneous catalysts are in a different state to the reactants.
e.g. Decomposition of hydrogen peroxide (solution) using manganese (IV)
oxide (solid) as a catalyst.
Manganese (IV) oxide (s)
2H2O2 (aq) 2H2O (l) + O2(g)
Common Catalysed Reactions
Catalyst Reaction catalysed Type of Catalyst
iron Haber process-ammonia manufacture Hetrogeneous
platinum Ostwald process - oxidation of
ammonia to make nitric acid
Hetrogeneous
nickel Hydrogenation of vegetable oils to
make margarine
Hetrogeneous
aluminium oxide Cracking of long chain hydrocarbons Hetrogeneous
titanium( IV) oxide Addition polymerisation -
manufacture of poly(ethene)
Hetrogeneous
Conc. sulphuric acid Esterification - making esters from
alcohol/carboxylic acid
Homogeneous
Heat
Very little
reaction
occurs
Fast reaction,
solution turns
green, gases
evolve rapidly
Solutions of
potassium
sodium tartrate
and hydrogen
peroxide (colourless)
CoCl2(aq)
+
Reaction
complete,
solution turns
pink again
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How Heterogeneous Catalysts Work
This type of catalyst is called a surface catalyst.
It works by adsorbing the reacting
molecules on to active sites and holding
them with weak bonds on its surface.
This not only causes the bonds within
the molecule to weaken but also helps
the collision geometry.
The reaction occurs on the surface with
less energy needed to form the
activated complex (lower activation
energy).
The products are formed and leave the
catalyst surface free for further
reactions
Catalyst Poisoning
A surface catalyst can be poisoned when another substance attaches
itself to the ‘active sites’. This is very often irreversible so prevents
reactant molecules from being adsorbed onto the surface.
For this reason, catalysts have to be regenerated or renewed.
E.g. Lead and its compounds are poisons of transition metal catalysts.
This is why unleaded petrol must be used in cars with catalytic
converters.
Catalysts can also be made ineffective by side-reactions. E.g. Iron used
in the Haber Process eventually rusts.
Active sites
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Enzymes
Enzymes catalyse the chemical reactions which take place in living cells.
Enzymes are complex protein molecules which are very specific- they
usually only speed up one particular reaction and work best at specific
temperatures and pH (optimum).
Examples in nature are:
Amylase – breaks down starch during digestion.
Catalase – breaks down hydrogen peroxide
Many enzymes are used in industry:
Invertase – used in chocolate industry for the hydrolysis of sucrose to
form fructose and maltose.
Zymase - converts glucose into alcohol in the brewing industry.
Protease (and others) – used in biological washing powders to dissolve
natural stains like protein .
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Energy Changes in Chemical Reactions
Chemical reactions involve a change in energy which often results in the
loss or gain of heat energy (exothermic/endothermic reactions)
The heat energy stored in a substance is called its Enthalpy
( H ).
The difference between the enthalpy of the reactants and the enthalpy
of the products in a reaction is the Enthalpy Change
(∆H):
∆H is measured in kJ per mole (kJ mol-1)
Potential Energy Diagrams
We can show the energy changes involved in exothermic and endothermic
reactions by using potential energy diagrams.
A chemical reaction can be regarded as a series of bond breaking and
bond making steps.
Consider the following reaction:
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Exothermic Reactions
Reactions which give out heat energy are called exothermic reactions.
The products have less enthalpy (potential energy) than the reactants
and the temperature of the surroundings increases.
E.g. the combustion of fuels
From this diagram we can work out:
The activation energy (Ea) which is the energy needed to start
the reaction.
The change in enthalpy between the reactants and products
(∆H)
Potential energy
or Enthalpy
kJ mol-1
Reaction pathway
∆H is always negative for exothermic reactions
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Endothermic Reactions
Reactions which absorb heat energy from the surroundings are called
endothermic reactions.
The products have more enthalpy than the reactants and the
temperature of the surroundings decreases.
From this diagram we can work out:
The activation energy (Ea)
The change in enthalpy between the reactants and products (∆H)
∆H is always positive for endothermic reactionsPotential energy
or Enthalpy
kJ mol-1
Reaction pathway
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Activation Energy
The Activation Energy is the ‘energy barrier’ which must be overcome
before the reactants can change into products.
The size of the Activation Energy will control how fast or slow a reaction
is. The higher the Activation Energy (or ‘barrier’) the slower the reaction.
If the activation energy is high, very few molecules will have enough
energy to overcome the energy barrier and the reaction will be slow.
e.g. Combustion of Methane
CH4 + 2O2 CO2 + 2H2O
This is a very exothermic reaction. At room temperature, no reaction
occurs as too few reactant molecules have sufficient energy to react
when they collide. Striking a match provides the molecules with enough
energy to overcome the barrier- it supplies the Activation Energy. Once
started, the energy given out by the reaction keeps it going.
Energy
Products
Reactants
Ea
Large Activation Energy
Slow Reaction
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The Activated Complex
When particles collide with the required Activation Energy (& geometry),
the activated complex is formed.
The activated complex is an unstable intermediate arrangement of
atoms formed as old bonds are breaking and new bonds are forming.
Energy is needed to form the activated complex as bonds in the reactants
may need to be broken, or charged particles brought together.
As the activated complex is very unstable it exists for a very short
period of time. From the peak of the energy barrier the complex can lose
energy to form either the products or the reactants again.
The higher the enthalpy change (∆H), the more unstable the activated
complex.
ReactantsActivated Complex Products
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Catalysts and Activation Energy
A catalyst provides an alternative pathway for the reaction with a lower
activation energy.
N.B The catalyst has no effect on the enthalpy change, ∆H
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PART 2 TRENDS IN THE PERIODIC TABLE
Development of the Periodic Table
• The periodic table was invented by Dimitri Mendeleev (1869).
• He arranged elements in order of increasing atomic
mass, and noted that their properties e.g. Melting point,
boiling point and density were periodic in nature (repeating
patterns existed). .
• Those elements with similar properties were placed below one
another in groups and gaps were left for unknown elements.
• The modern periodic table is based on an elements atomic number,
and this removed a number of the anomalies in the original version.
Trends in Physical Properties of the Elements
Melting points and boiling points
• Melting points and boiling points show periodic properties. This
means that they vary in a regular way or pattern depending on their
position in the Periodic Table.
• Melting points and boiling points depend on the strength of forces
which exist between the particles which make up a substance.
• Going down group 1 the alkali metals M.pt. decrease so there must
be a decrease in the strength of the force of attraction between
the particles.
• Going down group 7 the halogens m.pt. and b.pt. increases so there
must be a increase in the strength of the force of attraction
between the particles
Density
• The density of a substance is its mass per unit volume, usually in
gcm-3. They are found on pg 5 of the data booklet.
• Elements with densities greater than 0.5 gcm-3 are solids and
generally lie to the left-hand side of the Periodic Table.
• Elements with densities less than 0.5 gcm-3 are gases and generally
lie to the right-hand side of the Periodic Table.
In general,
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Covalent Radius - Atomic size
The size of an atom is difficult to measure because atoms do not have a
sharp boundary. However an X-ray technique can be used to measure the
distance between the nuclei of two covalently bonded atoms - this
distance is called the bond length.
The covalent radius of an atom is half the distance between the nuclei of
two of its covalently bonded atoms.
Trends in covalent radii
a. Across a period - covalent radii decreases
Across a period the protons are being added to the nucleus and electrons
are being added to the same shell. The increasing positive charge on the
nucleus pulls the outer electrons more closely and the covalent radius
decreases.
b. Down a group - covalent radii increases
Down a group each member has an
extra shell of electrons so the
covalent radius increases. The
positive charge on the nucleus
increases which tends to pull the
electrons closer but the effect of
adding an extra shell outweighs this.
eg. the bond length in a chlorine
molecule is 198 pm
(pm = picometre: 10-12 metre)
So the covalent radius of a
chlorine atom = 198/2 = 99 pm
These are found on pg 7 of the
data booklet.
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Covalent radius is an example of a Periodic Property as elements in the
same group appear at the same position on the wave.
First Ionisation Energy
This is defined as "the amount of energy required to remove one mole of
electrons from one mole of atoms in the gaseous state”
M (g) M+(g)
+ e 1st ionisation energy
The outermost electron will be the most weakly held and is removed first.
The ionisation energy is an enthalpy change and therefore is measured
per mole. Its units are kJmol-1 (kilojoules per mole). This is always an
endothermic process because energy is required to overcome the
attraction between the electron and the positive nucleus.
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Trends in Ionisation Energy
a. Across a period - ionisation energy increases
Across a period atomic size decreases. This means that the outer
electrons to be removed are closer to the nucleus and are held tightly.
Across a period the size of the positive charge on the nucleus also
increases attracting the electrons more strongly and making them more
difficult to remove.
b. Down a group - ionisation energy decreases
Down a group atomic size increases. This means that the outer electrons
to be removed are further from the nucleus and held less tightly.
However, the size of the nuclear charge is also increasing and you would
expect that the electrons would be strongly attracted to the bigger
nuclear charge. This is not the case. The outer electrons are shielded by
inner energy levels and do not feel the full attraction of the positive
nucleus. This is known as the screening effect. Outer electrons are
easier to remove.
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The second ionisation energy of an element is the energy required to
remove a second mole of electrons.
M+ (g) M2+
(g) + e 2nd ionisation energy
The total energy required to remove 2 moles of electrons is the sum of
the 1st and 2nd ionisation energies.
M (g) M2+
(g) + 2e
The second ionisation energy of an atom is always larger than the first as
it involves removing an electron from a species that is already positively
charged.
Ionisation energies are shown on pg 11 of the data booklet.
Note
1. The large increase from the 1st to the 2nd ionisation energy of lithium. The
1st ionisation energy removes the single outer electron from the lithium atom.
The 2nd ionisation energy requires breaking into a stable electron arrangement -
this requires a lot more energy.
2. With beryllium the largest increase comes between the 2nd and 3rd
ionisation energy. The 3rd ionisation energy requires breaking into a stable shell
of electrons.
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Electronegativity
Electronegativity is a measure of an atom’s attraction for the shared
pair of electrons in a bond.
Which atom would have a greater attraction for the electrons in this
bond and why?
Linus Pauling
Linus Pauling, an American chemist (and winner of two Nobel prizes!) came
up with the concept of electronegativity in 1932 to help explain the
nature of chemical bonds.
Since fluorine is the most electronegative element (has the greatest
attraction for the bonding electrons) he assigned it a value and compared
all other elements to fluorine.
Today we still measure electronegativities of elements using the Pauling
scale.
Values for electronegativity can be found on pg 11 of the data booklet.
Trends in Electronegativity
a. Across a period - electronegativity increases.
This is because the nuclear charge increases, attracting the electrons
more strongly to the nucleus. As a result, the electronegativity increases.
b. Down a group - electronegativity decreases.
Going down the group, the nuclear charge increases but the number of
electron shells also increases. As a result of ‘shielding’ and the increased
distance the outer shell is from the nucleus, electronegativity decreases.
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PART 3 BONDING AND STRUCTURE
Bonds are electrostatic forces (attractions between positive and
negative charges) which hold atoms together.
Atoms form bonds to become more stable - by losing, gaining or sharing
electrons.
The type of bond formed in a substance depends on the elements
involved and their position in the periodic table.
Metallic Bonding
Metallic bonding occurs between the atoms of metal elements.
Metals have little attraction for their outer electrons. These electrons
are free to move so are delocalised.
Electrons can move freely between these partially filled outer shells
creating what is called a ‘sea’ or ‘cloud’ of electrons around positive metal
cores.
The metallic bond is the electrostatic force between positively charged
core and delocalised outer electrons.
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Covalent and Polar Covalent Bonding
Covalent bonding occurs in non-metal elements.
A covalent bond is the electrostatic force of attraction between
positively charged nuclei and negatively charged outer electrons.
In the diatomic element chlorine both atoms have the same
electronegativity so the electrons are shared equally. This is called a pure
or non-polar covalent bond.
In the compound hydrogen iodide the bonded atoms have different
electronegativities. The iodine atom has a bigger attraction for the
shared electrons than the hydrogen atom. As the electrons are attracted
closer to the iodine it becomes slightly negative (δ-) and the hydrogen
atom becomes slightly positive (δ+).
This is called a polar covalent bond.
Other examples are
Ethanol Propanone
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Polar Molecules and Permanent Dipoles
Not all substances with polar covalent bonds will have ‘polar molecules’.
If there is a symmetrical arrangement of polar bonds, the polarity
cancels out over the molecule as a whole and the molecule is non polar.
e.g.
Carbon dioxide Tetrachloromethane
If the bonds are not symmetrical, the molecule has an overall polarity
and is said to have a permanent dipole, i.e. each end has a different
charge.
e.g. Hydrogen Chloride Water
Other examples are; Trichloromethane, ammonia, carbon monoxide
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Covalent Structure
Covalent and polar covalent substances are usually made up of discrete
molecules, but a few have giant covalent network structures.
e.g. Carbon dioxide – discrete molecules
Silicon Dioxide – covalent network structure
(images from BBC Higher Bitesize Chemistry)
Another example of an covalent network compound is silicon carbide.
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The Bonding Continuum
The greater the difference in electronegativity between two elements,
the less likely they are to share electrons, i.e. form covalent bonds.
To judge the type of bonding in any particular compound it is more
important to look at the properties it exhibits rather than simply the
names of the elements involved.
Ionic Bonding
Ionic bonds are formed between metal and non-metal elements with a
large difference in electronegativity.
The non-metal element with the high electronegativity gains the
electrons to form a negative ion:
e.g. Cl + e- Cl-
The element with the low electronegativity loses electrons to form a
positive ion:
e.g. Na Na+ + e-
Both the positive and negative ion will have the same electron
arrangement as a noble gas.
Ionic bonding is the electrostatic force of attraction between
positively and negatively charged ions.
Pure
Covalent
Bond
Polar
Covalent
Bond
Ionic Bond
Increasing ionic character
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Structure of Ionic Compounds
The forces of attraction between the oppositely charged ions results in
the formation of a regular structure called an ionic or crystal lattice.
E.g. Sodium chloride
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Attractions Between Covalent Molecules
There are attractive forces between covalent and polar covalent
molecules which can affect their properties.
These attractions between molecules are called van der Waals or
intermolecular forces (or attractions).
There are 3 types:
1. London Dispersion Forces
2. Dipole-dipole Attractions (permanent dipole-permanent dipole)
3. Hydrogen Bonds are a special type of dipole-dipole attraction which
is particularly strong.
1. London Dispersion Forces
This is the weakest form of intermolecular bonding and it exists between
all atoms and molecules.
Dispersion forces are caused by uneven distributions of electrons.
The atom or molecule gets slightly charged ends known as
a temporary dipole. This charge can then induce an opposite charge in a
neighbouring atom or molecule called an induced dipole. The oppositely
charged ends attract each other creating the intermolecular force. The
relative strength of the force depends on the size of the atoms or
molecules.
Dispersion forces increase with increasing atomic and molecular size.
Uneven distribution
of electrons in Helium
Temporary
Dipole
Induced
Dipole
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2. Permanent Dipole-Permanent Dipole Attractions
A polar molecule is one which has permanently charged ends (permanent
dipole).
Polar-Polar attractions (permanent dipole-permanent dipole) are the
intermolecular force of attraction between the oppositely charged ends
of the polar molecules.
Dipole to dipole attractions are the main forces of attraction between
polar molecules.
Effect of dipole-dipole attractions
Propanone Butane
Formula Mass 58 58
Structure
Main intermolecular permanent dipole- London force
permanent dipole Dispersion
Boiling Point 56oC 0oC
Polar molecules have higher boiling points than non-polar molecules of a
similar mass due to the permanent dipole-permanent dipole interactions.
Permanent dipole-permanent dipole interactions are stronger than London
Dispersion forces.
C
O
CCH
H
H
H
H
H
C C C C
H
H
H
H
H H
H
H
H
H
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3. Hydrogen Bonding
Hydrogen bonds are dipole- dipole interactions found between molecules
which contain highly polar bonds.
They are found in molecules where hydrogen is bonded to very
electronegative atoms like nitrogen, oxygen or flurine (NOF).
Other examples include ammonia, carboxylic acids and alcohols.
Hydrogen bonds are stronger than permanent dipole-permanent dipole
attractions and London dispersion forces but weaker than covalent
bonds.
When Hydrogen bonds are present, the compound will have a much higher
melting point (m.pt) and boiling point (b.pt) than other compounds of
similar molecular size.
Summary of intermolecular attractions
Van der Waals intermolecular forces are much weaker than covalent, ionic
and metallic bonding. The order of strength of the van der Waals forces
is summarised below.
Summary of van der Waals intermolecular forces
London Dispersion Dipole to Dipole Hydrogen
Weakest Strongest
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Bonding and Properties of Elements 1-20
Monatomic Elements - Noble Gases
All consist of single, unbonded atoms.
Only have London Dispersion forces between the atoms.
Properties
Low densities, m.pts and b.pts
Non conductors of electricity as no freely moving charged particles.
B.pts increase as the size of the atom increases.
This happens because more energy is required to break the stronger
London Dispersion forces.
b.p / oC
He
Ne
Ar
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Covalent Molecular Elements (in 1-20)
All consist of discrete molecules of varying size.
Fairly low m.pts, b.pts and densities.
Non-conductors of electricity.
Diatomic elements – H2, N2 , O2 , F2 , Cl2
As the size of the halogen atom increases, so does the
strength of the London dispersion forces. Therefore the boiling point
increase as more energy is required to separate the molecules.
b.p./oC
0
-160
--120
--80
--40
0
40
80
120
160
200
F Cl Br I
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Low melting point solids.
Phosphorus – P4
Sulphur - S8
Higher m.pt. because there are stronger London Dispersion forces
between larger molecules.
Fullerenes (Carbon)
Buckminster fullerene C60 (Bucky Balls) discovered in the 1980’s
Due to the large molecules , fullerenes have stronger dispersion forces
between their molecules than smaller
molecules.
NB – they are molecules not covalent networks
Nanotubes
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Covalent Network Elements (in 1-20)
Giant network structures containing millions of
atoms.
E.g. Carbon exists in 2 main forms…
Diamond
4 bonds per carbon atom – tetrahedral
structure
Non-conductor of electricity as no free
electrons (delocalised).
Hardest natural substance as many strong
bonds to break so used for drills, cutting
tools, etc.
Graphite
3 bonds per carbon atom – layered structure
with London dispersion forces between the
layers
Conductor of electricity due to one
delocalised electron per atom able to move
between the layers – used in electrodes.
Very soft – the layers break away easily
due to weak dispersion forces so good as
a lubricant and for drawing (pencils).
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Metallic elements (Revision of Nat 5)
All have metallic lattice structure.
The metallic bond is the attraction between
the positive metal cores and the outer
delocalises electrons.
Metallic bonds are generally strong. The strength of the bond will
depend on the number of outer electrons that each atom contributes to
the delocalised pool. The greater the number of outer electons, the
stronger the metallic bond.
They conduct electricity when solid or liquid due to free moving
delocalised outer electrons.
Positive core
Delocalised electrons
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The first twenty elements
The first twenty elements of the Periodic Table can be grouped according
to the type of bonding and structure.
Covalent molecular
A covalent molecular structure consists of discrete molecules held
together by London dispersion forces of attraction. Some elements
normally exist as solids, others exist as diatomic gases.
Covalent network
A covalent network structure consists of a giant lattice of covalently
bonded atoms.
Metallic
A metallic structure consists of a giant lattice of positively charged ions
and delocalised outer electrons.
Monatomic A monatomic structure consists of discrete (separate) atoms
held together my van der Waals' forces of attraction.
Complete the following table by including the appropriate type of bonding
and structure
A
B
C(i)
C(ii)
D
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Bonding and Properties of Compounds
Compounds can be split into 3 main groups, depending on their bonding,
structure and properties:
1. Ionic Lattice Structures
2. Covalent Network Structures
3. Covalent Molecular Structures
1. Ionic Lattice Structures
All ionic compounds are solids at room
temp so have high melting and boiling
points. This is because the ionic bonds
holding the lattice together are strong
and a lot of energy is required to
break them.
The stronger the ionic bond the higher
the melting point.
Ionic compounds conduct electricity
when dissolved in water or when molten as the ions are free to move.
Electrolysis of an ionic solution or melt causes a chemical change at the
electrodes.
They do not conduct when solid as the ions are ‘locked in the lattice and
cannot move to carry the current.
2. Covalent Network Structures
Covalent networks have very high melting and boiling points as many
strong covalent bonds need to be broken in order to change state. They
can also be very hard.
Silicon Carbide (SiC) – carborundum, has a similar structure to diamond.
It has a high melting point (2700oC). and it is used
as an abrasive.
CovalentBondTetrahedral
shape
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The 4 carbon atoms are available to bond with
another 4 silicon atoms resulting in a
covalentnetwork.
Covalent network structures are usually non-
conductors of electricity as they have no free
moving charged particles.
2. Covalent Molecular Structures
Usually have low melting and boiling points as there is little attraction
between their molecules.
E.g. Carbon dioxide CO2: m.pt -57oC (non-polar)
Compounds with polar molecules may have slightly higher m.pts and b.pts
than non-polar molecules due to permanent dipole-permanent dipole
attractions
e.g. Iodine chloride Bromine
I - Cl Br – Br
b.pt 97oC b.pt 59oC
Effects of Hydrogen bonding on properties.
1. Melting and boiling points
When hydrogen bonds are present, the compounds will have a much
higher m.pt and b.pt than other compounds of similar molecular size as
more energy is required to separate the molecules.
Water has a much higher b.pt than similar compounds containing hydrogen
= Carbon
= Silicon
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Hydrogen bonding explains why water, HF and NH3 have a b.pt higher than
expected.
Similarly HF b.p. 19 oC Whereas: HBr –68 oC and HI –35 oC
2. Viscosity
Viscosity is how thick a liquid is. As molecules get bigger the viscosity
increases.
Hydrogen bonding also affects the viscosity. The more hydrogen bonding
between molecules (ie the more OH groups present) the more viscous the
liquid.
Substance ethanol water glycerol
Molecular mass 46 18 92
Structural Formula
No of –OH groups 1 2 3
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3.Miscibility
Miscible liquids mix thoroughly. Ethanol and water would be described as
miscible but water and oil are immiscible as the oil forms a visible layer
on water.
Hydrogen bonding aids miscibility and ethanol and water both contain
hydrogen bonds so mix easily.
4.Density
Why do pipes burst when water freezes and why does ice float on
water?
As matter is cooled, it normally contracts and becomes more dense.
However, as water freezes it expands (at about 4oC) because the strong
hydrogen bonds between the molecules force them into an open lattice
structure.
This makes the solid ice less dense (takes
more space) than the liquid so ice floats
on water and pipes burst when water
freezes.
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Bonding, Solubility and Solutions
Ionic lattices and polar covalent molecular compounds tend to be:
Soluble in water and other polar solvents, due to the attraction
between the opposite charges.
Insoluble in non-polar solvents, as there is no attraction between
the ions and the solvent molecules.
Non-polar covalent molecular substances tend to be:
Soluble in non-polar solvents like carbon tetrachloride or hexane.
Insoluble in water and other polar solvents as there are no
charged ends to be attracted.
-
++
-
++- ++ -+ -
Water molecule
Ionic lattice
Hydratedions
--
+
+ -+
+
-
+ +
+-+
+
+ -
+
-
+ +
+ve ions attracted to –ve ends of water molecule
-ve ions attracted to +ve ends of water molecule
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Rate of Reactions - Glossary
Word Meaning
Activation energy (Ea)
The minimum amount of energy needed for a reaction to begin.
Catalyst A chemical which speeds up a chemical
reaction without being used up itself and which can be removed
chemically unchanged at the end of the reaction.
Catalytic converter A catalyst found in the exhaust of cars. It changes harmful
gases into less harmful gases.
It is usually made of platinum.
Chemical reaction An interaction between substances (chemicals) in which their
atoms re-arrange to form new substances.
Concentration The amount of particles in a given volume.
Enzyme A biological catalyst (ie. a catalyst found in living things).
Products The substances (chemicals) at the end of a chemical reaction.
Rate of reaction How quickly a reactant is used up OR how quickly a product is
created.
Surface area Total area of a substance which is exposed to the surroundings.
Heterogeneous The reactants are in a different state from the catalyst (the
catalyst is generally a solid).
Homogeneous The reactants and the catalyst are in the same state
Kinetic Energy The movement energy of particles.
Potential Energy The stored energy in reactants or products.
Activated Complex
a very unstable intermediate arrangement of atoms where the old
bonds are being broken and the new bonds are forming
Exothermic Giving out heat to the surroundings (feels hot). ∆H= -kJmol-1
Endothermic Taking heat in from the surroundings (feels cool). ∆H= +kJmol-1
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Trends in the Periodic Table - Glossary
Word Meaning
Covalent radius Half the distance between the nuclei of two bonded atoms of an
element.
Density The density of a substance is its mass per unit volume, usually in
gcm-3.
Electronegativity A measure of the attraction that an atom involved in a covalent
bond has for the electrons of the bond.
First ionisation energy The amount of energy required to remove one mole of electrons
from one mole of atoms in the gaseous state.
Periodicity The occurrence of patterns in the Periodic Table.
Screening (shielding)
effect
When inner electrons shield an outer electron from the
attractive effect of the nucleus and less energy is needed to
remove the outer electron as a result.
Bonding and Structure - Glossary
Word Meaning
Chemical bonding is the term used to describe the mechanism by which atoms are
held together.
Chemical structure describes the way in which atoms, ions or molecules are arranged.
Covalent bond a covalent bond is formed when two non-metal atoms share a pair
of electrons .
Covalent radius is half the distance between the nuclei of two bonded atoms of
an element Delocalised Delocalised electrons, in metallic bonding, are free from
attachment to any one metal ion and are shared amongst the
entire structure. Dipole an atom or molecule in which a concentration of positive charges
is separated from a concentration of negative charge.
Fullerenes are molecules of pure carbon constructed from 5- and 6-
membered rings combined into hollow structures. The most stable
contains 60 carbon atoms in a shape resembling a football.
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Word Meaning
Hydrogen bonds are electrostatic forces of attraction between molecules
containing a hydrogen atom bonded to an atom of a strongly
electronegative element such as fluorine, oxygen or nitrogen, and
a highly electronegative atom on a neighbouring molecule. Intermolecular forces Forces of attraction between molecules (or atoms)
Ionisation energy The first ionisation energy is the energy required to remove one
mole of electrons from one mole of gaseous atoms (i.e. one
electron from each atom). The second and subsequent ionisation
energies refer to the energies required to remove further moles
of electrons. Isoelectronic means having the same arrangement of electrons. For example,
the noble gas neon, a sodium ion (Na+) and a magnesium ion (Mg2+)
are isoelectronic. Lattice A lattice is a regular 3D arrangement of particles in space. The
term is applied to metal ions in a solid, and to positive and
negative ions in an ionic solid. London Dispersion
Forces are the intermolecular forces of attraction which result from the
electrostatic attraction between temporary dipoles and induced
dipoles caused by movement of electrons in atoms and molecules. Miscible fluids which mix with or dissolve in each other in all proportions.
Polar covalent bond a covalent bond between atoms of different electronegativity,
which results in an uneven distribution of electrons and a partial
charge along the bond. Non-polar (pure)
covalent bond a covalent bond between atoms of the same electronegativity,
which results in an even distribution of electrons .
van der Waals forces Is the general name given to all intermolecular attractions
including London dispersion, permanent dipole to permanent dipole
to dipole attractions and hydrogen bonding. Viscosity The thickness of a liquid.