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International Baccalaureate
Higher Level Chemistry Notes

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02 x1023)
Mole: amount of substance that contains 6
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2 Chemical Reactions and Equations
Properties of Chemical Reactions
-

In all chemical reactions, bonds in the reactants are broken and bonds in the products are
formed resulting in an energy change between the reacting system and its surroundings
There is a fixed relationship between the number of particles of reactants and products
resulting in no overall in mass

Chemical Equations
-

Reactants are written on the left hand side and the products on the right hand side
The number of moles of each element must be the same on both sides in a balanced
chemical equation

State Symbols
-

-

Because the physical state that the reactants and products are in can affect both the rate of
reaction and the overall energy change, the state symbols are included in the chemical
equation
(s)-solid, (l)-liquid, (g)-gas, (aq)-aqueous

Coefficients and Molar Ratio
-

The coefficients give information on the molar ratio

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31JK-1mol-1

Molar Volume of a Gas and Calculations
Molar Volume of a Gas
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One mole of any gas will occupy the same volume at the same temperature and pressure
At 273K and 1
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24 x10-2m3
P1V1/T1 = P2V2/T2

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2 Mass Spectrometer and Relative Atomic Mass
Mass Spectrometer
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Used to determine the relative atomic mass of an element

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2
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Acceleration: the resulting unipositive ions pass through holes in parallel plates under the
influence of an electric field where they are accelerated
4
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The amount of
deflection depends on the mass of the ion and its charge
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Detection: ions with a particular mass-to-charge ratio are then recorded on a detector which
measures both the mass and the relative amounts of all ions present

Uses of Radioactive Isotopes
-

Many, but not all, isotopes of elements are radioactive as the nuclei of these atoms break
down spontaneously
When they break down these radioisotopes emit radiation which is dangerous to living
things
Uses include nuclear power generation, the sterilisation of surgical instruments, crime
detection, finding cracks and stresses in metals and the preservation of food

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4 Electron Arrangement
Evidence from Ionisation Energies
-

First Ionisation Energy: the energy required to remove one electron from an atom in its
gaseous state
The highest value is for helium as the two electrons are in the lowest level and are held
tightly by the two protons
The graph of first ionisation energies provides evidence that the levels can contain different
numbers of electrons before they become full

Electron Arrangement
-

Known as its electron configuration
The electrons in the highest main energy level are known as valence electrons

Evidence for Sub-Levels
-

Successive ionisation energies for the same element can also be measured
As more electrons are removed the pull of the protons holds the remaining electrons more
tightly so increasingly more energy is required to remove them
A logarithmic scale is usually used
The successive ionisation energies of potassium also provides evidence of the number of
electrons in each main level
By looking where the first large jump occurs in successive ionisation energies, one can
determine the number of valence electrons
If the graph of first ionisation energies is examined more closely then it can be seen that the
graph does not increase regularly and hence providing evidence that the main levels are split
into sub-levels

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1 The Periodic Table and Physical Properties
The Periodic Table
-

In the Periodic Table elements are placed in order of increasing atomic number
Elements with the same number of valence electrons are placed vertically in the same group
Elements with the same outer shell of valence electrons are placed horizontally in the same
period

Atomic Radius
-

-

-

The atomic radius is the distance from the nucleus to the outermost electron
Since the position of the outermost electron can never be known precisely, the atomic
radius is usually defined as half the distance between the nuclei of two bonded atoms of the
same element
As a group descends the outermost electron is in a higher energy level
o This energy level is further away from the nucleus and hence the atomic radius is
greater
Across a period, electrons are being added to the same energy level, but the number of
protons in the nucleus increases
o This attracts the energy level closer to the nucleus and the atomic radius decreases
across a period

Ionic Radius
-

-

Both cations and anions increase in size down a group as the outer level gets further from
the nucleus
Cations
o Cations contain fewer electrons than protons so the electrostatic attraction between
the nucleus and the outermost electron is greater and the ion is thus smaller than
the parent atom
o It is also smaller because the number of electron shells has decreased by one
o Across the period the ions contain an increasing number of protons thus decreasing
the ionic radius
Anions
o Anions contain more electrons than protons so are larger than the parent atom
o Across a period the size decreases because the number of electrons remains the
same but the number of protons increases

Periodicity
-

Elements in the same group tend to have similar chemical and physical properties
There is a change in chemical and physical properties across a period
The repeating pattern of physical and chemical properties shown by the different periods is
known as periodicity

Melting Points
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Depends on both the structure and the type of attractive forces holding the atoms together

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At the left of the period elements exhibit metallic bonding which increases in strength as the
number of valence electrons increases
Silicon in the middle of the period has a macromolecular covalent structure with very strong
bonds resulting in a very high melting point
Elements in groups 5, 6 and 7 show simple molecular structures with weak van der Waals’
forces of attraction between the molecules
The noble gases exist as monatomic molecules with extremely weak forces of attraction
between the atoms
In group 1 the melting point decreases down the group as the atoms become larger and the
strength of the metallic bond decreases
In group 7 the van der Waals’ attractive forces between the diatomic molecules increase
down the group so the melting points increase

Electronegativity
-

A relative measure of the attraction that an atom has for a shared pair of electrons when it is
covalently bonded to another atom
As the size of the atom decreases the electronegativity increases
o The value increases across a period and decreases down a group

First Ionisation Energy
-

The energy required to remove an electron from an atom in its gaseous state
Decreases down each group as the outer electron is further from the nucleus and therefore
less energy is required to remove it
Generally the values increase across a period as extra electrons are filling the same energy
level and the extra protons in the nucleus attract this energy level closer making it harder to
remove an electron

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3 Oxides of the Third Period
Oxides of Period 3 Elements
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-

The oxides of sodium, magnesium and aluminium are all ionic
o This accounts for their high melting points and electrical conductivity when molten
Silicon dioxide has a diamond-like macromolecular structure with a high boiling point
On the right hand side of the period, the oxides are simple covalent molecular structures
with low melting and boiling points
o This results from the small difference in electronegativity between oxygen and these
atoms
The oxides of the electropositive elements are very basic and form solutions that are alkaline
The amphoteric nature of aluminium oxide can be seen from its reaction with hydrochloric
acid and sodium hydroxide
Silicon dioxide behaves as a weak acid, it doesn’t react with water but will form sodium
silicate with sodium hydroxide
The oxides of phosphorus, sulfur and chlorine are all strongly acidic

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5 D-Block Elements
The first Row Transition Elements
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An element that possess an incomplete d sub-level in one or more of its oxidation states
Scandium is not a typical transition metal as its common ion Sc 3+ has no d electrons
Zinc is not a transition metal as it contains a full d sub-level in all its oxidation states
For Cr and Cu it is more energetically favourable to half-fill rather than completely fill the d
sub-level respectively so they contain only one 4s electron

Characteristic Properties of Transition Elements
Variable Oxidation States
-

The 3d and 4s sub-levels are very similar in energy
When transition metals lose electrons they lose the 4s electrons first
All transition metals can show an oxidation state of +2
Some transition metals can form +3 and +4 oxidation states (Fe3+ and Mn4+)
o The ionisation energies are such that up to two d electrons can also be lost
o Cr(+3), Cr(+4), Mn(+4), Mn(+7), Fe(+3), Cu(+1)

Formation of Complex Ions
-

-

D-bock ions attract species that are rich in electrons because of the small size of the d-block
ions
Such species are known as ligands
Ligands: neutral molecules or anions which contains a non-bonding pair of electrons
o These electron pairs can form co-ordinate covalent bonds with the metal ion to form
complex ions
A common ligand is water (Fe(H2O)6)+3
Ligands can be replaced by other ligands
Co-Ordination Number: the number of lone pairs bonded to the metal ion
o Compounds with a CON of 6 are octahedral in shape
o Those with a CON of four are tetrahedral or square planar
o Those with a CON of two are usually linear

Catalytic Behaviour
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Many transition elements and their compounds are very efficient catalysts
This aids industrial processes such as the production of ammonia and sulfuric acid

Coloured Complexes
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In the free ion the five d orbitals are of equal energy
In complexes the d orbitals are split into two distinct levels
The difference in energy between these two levels corresponds to a particular wavelength
or frequency in the visible region of the spectrum
When light falls on the complex, energy of that particular wavelength is absorbed and
electrons are excited from the lower level to the higher level
The amount the orbitals are split depends on the nature of the transition metal, the
oxidation state, the shape of the complex and the nature of the ligand
o This explains why different complexes have different colours

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If the d orbital is completely empty (Sc 3+) or completely full (Zn2+), no transitions within the d
level can take place and the complexes are colourless

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8 for
ionic bonding to occur
Ionic bonding results in the molecule having a much greater melting point

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3 Shapes of Simple Molecules and Ions
VSEPR Theory
-

-

The shapes of simple molecules and ions can be determined by using the Valence Shell
Electron pair Repulsion Theory
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4 Intermolecular Forces and Allotropes
of Carbon
Molecular Polarity
-

Whether a molecule is polar or not, depends both on the relative electronegativities of the
atoms in the molecule and on its shape
If the individual bonds are polar then it does not necessarily mean that the molecule will be
polar as resultant dipoles may cancel out all the individual dipoles

Intermolecular Forces
-

The covalent bonds between the atoms within a molecule are very strong
The forces between the molecules are much weaker
These intermolecular forces depend on the polarity of the molecules

Van der Waals’ Forces
-

Even in non-polar molecules, the electrons can at any one moment be unevenly spread
This produces temporary instantaneous dipoles
An instantaneous dipole can induce another dipole in a neighbouring particle resulting in a
weak attraction between the two particles
Van der Waals’ forces increase with increasing mass

Dipole : Dipole Forces
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Polar molecules are attracted to each other by electrostatic forces
Although still relatively weak the attraction is stronger than Van der Waals’ forces

Hydrogen Bonding
-

Occurs when hydrogen is bonded directly to a small highly electronegative element, such as
fluorine, oxygen, nitrogen
As the electron pair is drawn away from the hydrogen atom by the electronegative element,
all that remains is the proton in the nucleus as there are no inner electrons
The proton then attracts a non-bonding pair of electrons from the F, N or O resulting in a
much stronger dipole:dipole attraction
Molecules with hydrogen bonding have much higher boiling points as the hydrogen bonding
is harder to break down
This trend can be seen in the boiling points of the group 5 and 7 hydrides

Allotropes of Carbon
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Allotropes occur when an element can exist in different crystalline forms

Diamond
-

Each carbon atom is covalently bonded to four other carbon atoms to form a giant covalent
structure
All the bonds are equally strong and there is no plane of weakness in the molecule so
diamond is extremely hard
Also, because all the electrons are localised it does not conduct electricity

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Both silicon and silicon dioxide form similar giant tetrahedral structures

Graphite
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Each carbon atom has very strong bonds to three other carbon atoms to give layers of
hexagonal rings
There are only very weak bonds between the layers though
The layers are able to slide over each other, graphite is therefore an excellent lubricant
And, because the electrons are delocalised between the layers it is a good conductor of
electricity

Buckminsterfullerene
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Sixty carbon atoms are arranged in hexagons and pentagons to give a geodesic spherical
structure similar to a football
Since its discovery, many other similar carbon molecules have been isolated, leading to the
new branch of science called nanotechnology

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6 Molecular Orbitals and Hybridization
Combination of Atomic Orbitals to Form Molecular Orbitals
-

Although the Lewis representation is a useful model to represent covalent bonds, it makes
the false assumption that all valence electrons are the same
More advanced model of bonding considers the combination of atomic orbitals to form
molecular orbitals

Sigma (σ) Bonds
-

A sigma bond is formed when two atomic orbitals on different atoms overlap along a line
drawn through the two nuclei
Occurs when two s orbitals overlap, an s orbital overlaps with a p orbital, or when two p
orbitals overlap ‘head on’

Pi (π) Bonds
-

A pi bond is formed when two p orbitals overlap ‘sideways on’
Overlap occurs above and below the line drawn through the two nuclei
A pi bond is made of two regions of electron density

Hybridization
sp3 Hybridization
-

-

Methane is a good example of sp3 hybridization
It contains four C-H bonds pointing towards the corners of a tetrahedron with bond angles
of 109
...
5˚ when they overlap with the s orbitals on the hydrogen
atoms
When the carbon atom bonds in methane one of its 2s electrons is promoted to a 2p orbital
and then the 2s and three 2p orbitals hybridize to form four new hybrid orbitals
These four new orbitals arrange themselves to be as mutually repulsive as possible
(tetrahedrally)
Four equal sigma bonds can then be formed with the hydrogen atoms

sp2 Hybridization
-

-

Occurs in ethene
After a 2s electron on the carbon atom is promoted the 2s orbital hybridizes with two of the
2p orbitals to form three new planar hybrid orbitals with a bond angle of 120˚ between
them
These can form sigma bonds with the hydrogen atoms and also a sigma bond between the
two carbon atoms
Each carbon atom now has one electron remaining in a 2p orbital
These can develop to form a pi bond
Ethene is a planar molecule with a region of high density above and below the plane

sp Hybridization
-

Occurs when the 2s orbital hybridizes with just one of the 2p orbitals to form two new linear
sp orbitals with an angle of 180˚ between them

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The remaining two p orbitals on each carbon atom then overlap to form two pi bonds
Occurs in ethyne

Relationship between Type of Hybridization, Lewis Structure and
Molecular Shapes
-

Molecular shapes can be arrived at by either knowing the VSEPR theory or by knowing the
type of hybridization
Hybridization can take place between any s and p orbital in the same energy level and is not
just restricted to carbon compounds
If the shape and bond angles are known from the Lewis structure then the type of
hybridization can be deduced
Similarly if the hybridization is known, the shape and bond angles can be deduced

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1-Enthalpy Changes
Exothermic and Endothermic Reactions
-

Energy is defined as the ability to do work, that is, move a force through a distance
In a chemical reaction energy is required to break the bonds in the reactants, and energy is
given out when new bonds are formed in the products
If the bonds in the products are stronger than the bonds in the reactants then the reaction is
said to be exothermic, heat is given out to the surroundings
In endothermic reactions heat is absorbed from the surroundings as the bonds in the
reactants are stronger than the bonds in the products
The internal energy stored in the reactants is known as its enthalpy (H)
The absolute enthalpy of neither the reactants nor the products can be measured but the
difference between them can (∆H = Hproducts – Hrea ctants)
In exothermic reactions ∆H has a negative value and the products are more stable than the
reactants
In endothermic reactions ∆H has a positive value and the reactants are more stable than the
products
The standard enthalpy change of a reaction is denoted by ∆Hø

Temperature and Heat
-

Heat: a measure of the total energy in a given amount of substance and therefore depends
on the amount of substance present
Temperature: a measure of the hotness of a substance which represents the average kinetic
energy of the substance, but is independent of the amount of substance present

Calorimetry
-

The enthalpy change of a reaction can be measured experimentally by using a calorimeter
In an exothermic reaction all the heat evolved is used to raise the temperature of a known
mass of water
In an endothermic reaction the heat transferred from the water to the reaction can be
calculated by measuring the lowering of temperature of a known mass of water

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18kJkg-1K-1, that is, it requires 4
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3- Bond Enthalpies and Hess’s Law
Bond Enthalpies
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Enthalpy changes can also be calculated from bond enthalpies
For bond formation the value is negative and for bond breaking the value is positive
o ∆H = ∑Hreactants - ∑Hproducts

Limitations of Using Average Bond Enthalpies
-

Average bond enthalpies can only be used if all the reactants and products are in the
gaseous state
Average Bond Enthalpies: the average energy required per mole to break a bond in the
gaseous state obtained from a range of similar compounds

Hess’s Law
-

The enthalpy change for a reaction depends only on the difference between the enthalpy of
the products and the enthalpy of the reactants
It is independent of reaction pathway
This law is a statement of conservation of energy
It can be used to determine enthalpy changes, which cannot be measured directly

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5 Entropy and Spontaneity
Disorder
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Systems naturally tend towards disorder
Greatest increase in disorder is found where the number of gaseous particles increases
Change in disorder of a system is known as the entropy change, ∆S
The more disordered the system becomes, the more positive ∆S becomes

Absolute Entropy Values
-

Standard entropy change of a substance is the entropy change per mole that results from
heating the substance from 0K to the standard temperature of 298K
∆Sø = Sø(products) - Sø(reactants)

Spontaneity
-

A reaction is spontaneous if it causes a system to move from a less stable to a more stable
state
Depends upon the enthalpy change and the entropy change
Expressed as the Gibbs energy change, ∆G
∆Gø = ∆Hø - T∆Sø
For a reaction to be spontaneous it must be able to do work
Therefore, ∆Gø must be negative

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1- Rates of Reaction and Collision Theory
Rate of Reaction
-

Chemical kinetics is the study of the factors affecting the rate of a chemical reaction
Reaction Rate: the increase in the concentration of one of the products per unit time or the
decrease in the concentration of one of the reactants per unit time (mol dm-3 s-1)
The change in concentration can be measured by using any property that differs between
the reactants and the products
When the concentration of the reactant/product is graphed, the rate at any given moment
in time is the gradient of the graph at that point in time
Rates of reaction usually decrease with time as the reactants are used up

Maxwell-Boltzmann Distribution
-

The moving particles in a liquid or gas do not all trave with the same velocity
The faster they move the more kinetic energy they possess
The Maxwell-Boltzmann Distribution illustrates the distribution of the kinetic energies of
particles
As the temperature increases the area under the curve does not change as the total number
of particles remains constant
However more particles have a higher average kinetic energy which leads to a broadening of
the curve
The peak of the curve shifts to the right and decreases in height

Collision Theory
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-

For a reaction to occur three conditions must be met:
o The particles must collide
o They must collide with the appropriate geometry or orientation so that the
respective parts of the particles come into contact with each other
o They must collide with sufficient energy to bring about the reaction
Activation Energy: the minimum amount of energy required for a reaction to occur
Any factor which either increase the frequency of the collisions or increases the energy with
which they collide will make the reaction go faster

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3- Order of Reaction and Half-Life
Rate Expressions
-

The rate of reaction between two reactants can be followed experimentally
The rate will be found to be proportional to the concentration of A raised o some power and
also the concentration of B raised to some power
Rate = k[A] x[B]y
K is the constant of proportionality and is known as the rate constant
x is known as the order of reaction with respect to A
y is known as the order of reaction with respect to B
The overall order of reaction = x + y

Units of Rate Constant
-

The units of the rate constant depend on the overall order of reaction
k = Rate/[A] x[B]y
First order: s-1
Second order: mol-1 dm3 s-1
Third order: mol-2 dm6 s-1

Deriving a Rate Expression by Inspecting Data
-

By examining experiments in which only one of the reactants has changed concentrations
and the corresponding rate, the rate of reaction with respect to that reactant can be
determined

Half-Life, t1/2
-

Half-Life: the time taken for the concentration of a reactant to half its original value
First order reactions have a constant half-life

Graphical Representations of Reactions
-

-

-

Zero Order:
o Concentration-time graph: straight negative line
o Rate-concentration graph: constant
First Order:
o Concentration-time graph: exponential line
o Rate-concentration graph: straight positive line from the origin
Second Order:
o Concentration-time graph: steeper exponential line
o Rate-concentration graph: exponential line from origin

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1- The Equilibrium Law
Dynamic Equilibrium
-

Most chemical reactions do not go to completion
Once some products are formed the reverse reaction can take place and reform the
reactants
In a closed system the concentrations of all the reactants and products will eventually
become constant
At a dynamic equilibrium the forward and reverse reactions continue to occur but at
equilibrium the rate of the forward and reverse reactions are equal
Dynamic equilibrium also occurs when physical changes take place

Closed System
-

A closed system is one in which neither matter nor energy can be lost or gained from the
system
The macroscopic properties remain constant
If the system is open some of the products from the reaction could escape and equilibrium
would never be reached

Equilibrium Constant
-

aA + bB  cC + dD
Equilibrium constant, Kc = [A]a[B] b/[C]c[D]d
Homogeneous reactions are ones in which all of the reactants and products are in the same
phase

Magnitude of the Equilibrium Constant
-

The magnitude of the equilibrium constant is related to the position of equilibrium
When the reaction goes nearly to completion, the equilibrium constant is very large
When the reaction hardly proceeds, the equilibrium constant is very small

Le Châtelier’s Principle
-

If a system at equilibrium is subjected to a small change the equilibrium tends to shift so as
to minimise the effect of the change
Provided the temperature remains constant the value for Kc must remain constant
If the concentration of the reactants is increased, or one of the products is removed from
the equilibrium mixture then more of the reactants must react in order to keep Kc constant
o Equilibrium shifts to the right

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3 Equilibrium Calculations and Phase
Equilibrium
Units of the Equilibrium Constant
-

The units of Kc depend on the powers of the concentrations in the equilibrium expression
𝐾𝑐 =

[𝑁𝐻3 ] 2
[𝐻2 ] 3 𝑥 [𝑁2 ]

=

𝑐𝑜𝑛𝑐𝑒𝑛𝑡𝑟𝑎𝑡𝑖𝑜𝑛2
𝑐𝑜𝑛𝑐𝑒𝑛𝑡𝑟𝑎𝑡𝑖𝑜𝑛4

= 𝑐𝑜𝑛𝑐𝑒𝑛𝑡𝑟𝑎𝑡𝑖𝑜𝑛 −2 = 𝑑𝑚6 𝑚𝑜𝑙 −2

If the concentration is the same on the top and bottom, Kc has no units

Phase Equilibrium
-

-

Dynamic equilibrium between a liquid and its vapour occurs when the rate of vaporization is
equal to the rate of condensation
Vapour Pressure of a liquid is the pressure exerted by the particles in the vapour phase
It is independent of the surface area of the liquid or of the size of the container
Although, in bigger containers it may take longer for equilibrium to establish
Vapour pressure of any liquid depends on both the strength of the molecular forces holding
the liquid particles together and on the temperature
The stronger the intermolecular forces the lower the vapour pressure at a particular
temperature
Enthalpy of vaporisation is the enthalpy change required to overcome the intermolecular
forces
Eg
...
1- Theories of Acids and Bases and Salt Hydrolysis
The Ionic Theory
-

An acid is a substance which produces hydrogen ions in aqueous solution
A base is a substance that can neutralise an acid
An alkali is a base that is soluble in water

Brønsted-Lowry Acids and Bases
-

-

A Brønsted-Lowry acid is a substance that can donate a proton
A Brønsted-Lowry base is a substance that can accept a proton
Conjugate acids and bases are two substances on different sides of a reaction that differ by
one H+ ion
o A conjugate base of an acid is the species remaining after the acid has lost a proton
o A conjugate acid of a base is the species remaining after the acid has gained a
proton
Substances that can act as both an acid and a base are called amphiprotic

Lewis Acids and Bases
-

A Lewis Base is a substance which can donate a pair of electrons
A Lewis Acid is a substance which can accept a pair of electrons
In the process a co-ordinate covalent bond is formed between the base and the acid

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3- The pH Scale
The pH Scale
-

-

pH is defined as being equal to minus the logarithm to the base ten of hydrogen ion
concentration
o pH = -log10[H+]
pH of pure water is 7
Acids have pH lower than 7 and bases have a pH greater than 7
pH scale runs from 0 to 14
A change in one unit on the pH scale corresponds to a tenfold change in the hydrogen ion
concentration

The Log10 Scale and p-Scale
-

Normal Scale: distance between each number is equal
Log10 Scale: distances between powers of ten are equal
p-Scale: equal to minus the power of ten in the logarithmic scale
o Sometimes used by chemists to express equilibrium constants and concentration

Determination of pH
-

The pH of a solution can be determined by using a pH meter of by using ‘universal’ indicator
Universal indicator is red for acidic solutions and progresses to blue for alkaline solutions

Strong, Concentrated and Corrosive
-

Strong: completely dissociated into ions
Concentrated: a high number of moles of solute per litre of solution
Corrosive: chemically reactive
Weak: only slightly dissociated into ions
Dilute: a low number of moles of solute per litre of solution

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00 𝑥 10−14 𝑚𝑜𝑙 2 𝑑𝑚 −6 at 298K, where Kw is the ionic
product of water
The dissociation of water into its ions is an endothermic process, the value of Kw will increase
as the temperature increases
For pure water [ 𝐻 +( 𝑎𝑞 )] = [ 𝑂𝐻 − ( 𝑎𝑞 )] = 1
...
5 Calculations with Weak Acids and
Bases
Weak Acids
-

The dissociation of a weak acid, HA, in water can be written:
o 𝐻𝐴 ( 𝑎𝑞 ) ⇌ 𝐻 + ( 𝑎𝑞 ) + 𝐴− (𝑎𝑞)
The equilibrium expression for this reaction is:
o

-

𝐾𝑎 =

[ 𝐻+ ] 𝑥 [𝐴 −]
[𝐻𝐴]

, Ka is the acid dissociation constant

If the acids are quite weak the equilibrium concentration of the acid can be assumed to be
the same as its initial concentration
...
6 Salt Hydrolysis and Buffer Solutions
Salt Hydrolysis
-

-

Sodium chloride is neutral in aqueous solution
o It is the salt of a strong acid and a strong base
Salts made from a weak acid and a strong base (sodium Ethanoate) are alkaline in solution
o Ethanoate ions will combine with hydrogen ions from water to form mainly
undissociated ethanoic acid, leaving excess hydroxide ions in solution
Similarly, salts derived from a strong acid and a weak base will be acidic in solution
The acidity of salts depends on the size and charge of the cation
o 𝐴𝑙𝐶𝑙3 ( 𝑠) + 3𝐻2 𝑂 ( 𝑙 ) → 𝐴𝑙(𝑂𝐻)3 ( 𝑠) + 3𝐻𝐶𝑙(𝑎𝑞)
o The +3 charge is spread over a very small ion, which gives the Al 3+ ion a very high
charge density
o The lone pair of one of the six water molecules surrounding the ion will be strongly
attracted to the ion and the water molecule will lose a hydrogen ion in the process
o This process continues until aluminium hydroxide is formed

Buffer Solutions
-

-

-

-

A buffer solution resists changes in pH when small amounts of acid or alkali are added to it
An acidic buffer solution can be made by mixing a weak acid together with the salt of that
acid and a strong base
o Ethanoic acid and sodium Ethanoate
o The weak acid is only slightly dissociated in solution, but the salt is fully dissociated
into its ions, so the concentration of Ethanoate ions is high
o 𝑁𝑎𝐶𝐻3 𝐶𝑂𝑂 ( 𝑎𝑞 ) → 𝑁𝑎+ ( 𝑎𝑞 ) + 𝐶𝐻3 𝐶𝑂𝑂 −(𝑎𝑞)
o 𝐶𝐻3 𝐶𝑂𝑂𝐻 ( 𝑎𝑞 ) → 𝐶𝐻3 𝐶𝑂𝑂 −( 𝑎𝑞 ) + 𝐻 + (𝑎𝑞)
If an acid is added the extra H+ ions coming from the acid are removed as they combine with
Ethanoate ions to form undissociated ethanoic acid, so the concentration of H + ions remains
unaltered
If an alkali is added the hydroxide ions from the alkali are removed by their reaction with the
undissociated acid to form water, so again the H+ ion concentration stays constant
An alkali buffer with a fixed pH greater than 7 can be made from a weak base together with
the salt of that base with a strong acid
o 𝑁𝐻4 𝐶𝑙( 𝑎𝑞 ) → 𝑁𝐻4 + ( 𝑎𝑞 ) + 𝐶𝑙 +(𝑎𝑞)
o 𝑁𝐻3 ( 𝑎𝑞 ) + 𝐻2 𝑂 ( 𝑙 ) ⇌ 𝑁𝐻4 +( 𝑎𝑞 ) + 𝑂𝐻 −(𝑎𝑞)
If H+ ions are added they will combine with OH- ions to form water and more of the
ammonia will dissociate to replace them
If more OH- ions are added they will combine with ammonium ions to form undissociated
ammonia
In both cases the hydroxide ion concentration and the hydrogen ion concentration remain
constant

8
...
At this point, which
equates to the half equivalence point when ethanoic acid is titrated with sodium
hydroxide, the pH of the solution will equal the pKa value of the acid

9
...
In an ionic compound between two elements the oxidation number of each element is equal
to the charge carried by the ion
a
...
For covalent compounds assume the compound is ionic with the more electronegative
element forming the negative ion
a
...
The algebraic sum of all oxidation numbers in a compound is equal to zero
a
...
The algebraic sum of all the oxidation numbers in an ion is equal to the charge on the ion
a
...
Elements not combined with other elements have an oxidation number of zero
a
...
Oxygen when combined always has an oxidation number of -2 except in peroxides (H2O2)
where it is -1
7
...
2 Reactivity Series
Reactivity
-

Reactivity with water decreases down group 1
In all of these reactions the metal is losing electrons
o It is being oxidized and in the process it is acting as a reducing agent
A reactivity series of reducing agents can be deduced by considering the reactivity of metals
with water and acids
The more readily the metal losses its outer electrons the more reactive it is
Metals higher in the series can displace metal ions lower in the series from the solution
The series can be extended for oxidizing agents, the most reactive oxidising agent will be the
species that gains electrons the most readily
Oxidising agents lower in the series gain electrons from species higher in the series

Simple Voltaic Cells
-

-

A half-cell is simply a metal in contact with an aqueous solution of its own ions
Voltaic cells consist of two half-cells, connected together to enable the electrons transferred
during the redox reaction to produce energy in the form of electricity
Connected by an external wire and a salt bridge which allows free movement of ions
The electrons will flow from the metal which is higher in the reactivity series to the one
which is lower
The voltage produced by a voltaic cell depends on the relative difference between the two
metals in the reactivity series
o The greater the difference the greater the voltage produced
In a voltaic cell:
o The anode is the negative electrode where oxidation occurs
o The cathode is the positive electrode where reduction occurs

9
...
00V) at 1 atm, 298K and 1
...
4 Electrochemical Series
Electrochemical Series
-

Found in the Data Booklet
The standard electromotive force (emf) of any cell E øcel l is simply the difference between the
standard electrode potentials of the two half-cells

Electron Flow and Spontaneous Reactions
-

By using the standard electrode potentials it is easy to determine what will happen when
two half-cells are connected together
The electrons will always flow from the more negative half-cell to the more positive half-cell
Positive Eøcel l values give a negative ∆Gø values as the reaction can provide electrical energy
(i
...
do work, and the reaction is spontaneous)

10
...


2
...

4
...
1 carbon = methb
...
3 carbons = propd
...
5 = pent-, 6 = hex-, 7 = hept-, 8 = octIdentify the type of bonding in the chain or ring:
a
...
Double bonds = -ene
c
...
May become at the beginning or the end of the name
Numbers are used to give the positions of groups or bonds along the chain

Homologous Series
-

Have the same general formula with the neighbouring members of the series differing by –
CH2
The chemical properties of the individual members of an homologous series are similar and
they show a gradual change in physical properties

Common Functional Groups
Name
Alkane

Formula
R–H

Alkene
Alkyne
Alcohol

R – OH

Amine
Halogenoalkane

R – NH2
R–X

Aldehyde
Ketone
Carboxylic Acid
Ester

R – COH
R – COR’
R – COOH
[RCOOR’]

Ending Name
-ane
-ene
-yne
-ol (position determines 1,2 or
3)
-amine
Fluoro-, chloro-, bromo-, iodo-anal
-anone
-anoic acid

Examples
Methane, Butane, 2methylpropane

Ethanol, propan-1-ol, propan-2ol
Ethylamine, 2-aminobutane
Bromoethane, 1,2dichloroethane
Ethanal, propanal
Propanone, pentan-2-one
Methanoic Acid, Propanoic Acid
Ethyl Ethanoate, propyl
methanoate

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...
3 Structural Isomers
Structures of Hydrocarbons
-

Structural Isomers: have the same molecular formula but a different structural formula
Normally have similar chemical properties but slightly different physical properties

Isomers of Alkanes
-

Each carbon atom contains four single bonds
There is only one possible structure for methane, ethane and propane
There are two possible structures for butane
Three structural formulas of pentane

Structures of Alkenes
-

Ethene and propene only have one possible structure
Butane has three structural isomers

Classification of Alcohols and Halogenoalkanes
-

Depends on how many other carbon atoms the carbon atom containing the functional group
is bonded to
Primary alcohol: carbon containing functional group is bonded to one other carbon
Secondary alcohol: carbon containing functional group is bonded to two other carbon atoms
Tertiary alcohol: carbon containing functional group is bonded to three other carbon atoms

Naming Structural Isomers
-

The IUPAC names to distinguish between structural isomers of alcohols, aldehydes, ketones,
carboxylic acids and halogenoalkanes containing up to six carbon atoms are required

10
...
+ Cl
...
 H3C
...

Since a new radical is produced this stage is called propagation
The methyl radical is also highly reactive and reacts with a chlorine molecule to form the
product and regenerate another chlorine radical
This is a further propagation step and enables a chain reaction to occur as the process can
repeat itself
CH3
...

Termination occurs when two radicals react together
Cl
...
 Cl2
Further substitution can occur when chlorine radicals react with the substituted products
The substitution can continue even further to produce trichloromethane and then
tetrachloromethane
The overall mechanism is called free radical substitution

10
...
6- Alcohols
Combustion
-

-

Ethanol is used both as a solvent and as a fuel
In a plentiful supply of oxygen it combusts completely to give carbon dioxide and water
o C2H5OH(l) + 3O2(g)  2CO2(g) + 3H2O(l)
o General equation for complete combustion C n H2n+1OH + (2n-1)O2  nCO2 + (n+1)H2O
When is partially combusts in oxygen it produces either carbon monoxide or carbon and in
both cases produces water

Oxidation of Ethanol
-

-

Ethanol can be readily oxidised by warming with an acidified solution of potassium
dichromate(VI)
During the process the orange dichromate(VI) ion Cr2O72- is reduced from an oxidation state
of +6 to the green Cr3+ ion
Simple breathalyser tests use this process:
o A motorist who is suspected of having exceeded the alcohol limit blows into a bag
containing crystals of potassium dichromate
o The crystals then change colour from orange to green but the magnitude of alcohol
present in the breath changes the magnitude of the colour change
Ethanal does not have hydrogen bonding and so has a much lower boiling point than ethanol
and ethanoic acid
To stop the reaction at the aldehyde stage the ethanal can be distilled from the reaction
mixture as soon as it is formed
If the complete oxidation to ethanoic acid is required, then the mixture can be heated under
reflux so that none of the ethanal can escape

Oxidation of Alcohols
-

Primary Alcohols: alcohol group joined to a carbon atom which is only joined to one other
carbon atom
Secondary Alcohols: alcohol group joined to a carbon atom which is joined to two other
carbon atoms
Tertiary Atoms: alcohol group joined to a carbon atom which is joined to three other carbon
atoms
Primary alcohols can be oxidized from alcohols to aldehydes and then carboxylic acids
Secondary alcohols can only be oxidized from alcohols to ketones but not carboxylic acids
Tertiary alcohols cannot be oxidized from alcohols into either aldehydes/ketones or
carboxylic acids
o Tertiary alcohols cannot be oxidised as they have no hydrogen atoms attached
directly to the carbon atom containing the alcohol group
o It is not true to say that tertiary alcohols can never be oxidised, as they burn readily,
but when this happens the carbon chain is destroyed

10
...
It
shows where they come from and where they move to

Mechanism of Nucleophile Substitution
Primary Halogenoalkanes
-

The reaction between bromoethane and warm dilute sodium hydroxide
The proposed mechanism involves the formation of a transition state which involves both of
the reactants
Because the molecularity of this single-step mechanism is two it is known as an SN2
mechanism (bimolecular Nucleophilic substitution)
To explain the process through diagrams:
o Draw the halogenoalkane and the sodium hydroxide separately
o On these diagrams, draw curly arrows from the non-bonding pair of the sodium
hydroxide ion to the carbon atom and an arrow from the bond between the carbon
atom and the halogen atom
o The next diagram portrays the sodium hydroxide and halogen atom connected to
the carbon atom with a dashed line and an overall charge on the molecule of -1
o The final diagram has the sodium hydroxide fully connected to the carbon atom and
the halogen ion separately

Tertiary Halogenoalkanes
-

Reaction between 2-bromo-2-methylpropane and warm dilute sodium hydroxide
A two-step mechanism is proposed that is consistent with the rate expression
The mechanism is known as an SN1 reaction as it is a Unimolecular Nucleophilic substitution

Secondary Halogenoalkanes
-

The mechanism for the hydrolysis of secondary halogenoalkanes is more complicated as
they can proceed by either SN1 or SN2 pathways or a combination of both

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...
9- Nucleophilic Substitution
Nucleophilic Substitution
-

The reaction between halogenoalkanes and a warm dilute aqueous solution of sodium
hydroxide is a Nucleophilic substitution reaction
The nucleophiles are attracted to the positive dipole on the carbon atom and substitute the
halogen atom in the halogenoalkane
Primary halogenoalkanes = SN2
Tertiary halogenoalkanes = SN1
There are several factors which affect the rate of the substitution reaction

The Nature of the Nucleophile
-

The effectiveness of a nucleophile depends on its electron density
Anions tend to be more reactive than the corresponding neutral species
Among species with the same charge a less electronegative atom carrying a non-bonded pair
of electrons is a better nucleophile than a more electronegative one
This is because the less electronegative atom can more easily donate its pair of electrons as
they are held less strongly
CN- > OH- > NH3 > H2O

The Nature of the Halogen
-

For both the SN1 and SN2 reactions the iodoalkanes react faster than bromoalkanes which in
turn react faster than chloroalkanes
This is due to the relative bond energies as the C-I bond is much weaker than the C-Cl bond
and therefore breaks more rapidly

The Nature of the Halogenoalkane
-

Tertiary halogenoalkanes react faster than secondary halogenoalkanes, which in turn react
faster than primary halogenoalkanes
The SN1 route, which involves the formation of an intermediate carbocation, is faster than
the SN2 route, which involves a transition state with a relatively high activation energy

Substitution with Ammonia and Potassium Cyanide
-

-

-

In addition to forming alcohols when water or hydroxide ions are used as the nucleophile,
halogenoalkanes can react with ammonia to form amines and with cyanide ions to form
nitriles
Bromoethane reacts with ammonia to form ethylamine
...
Even the tertiary amine is still a nucleophile and can
react further to form the quaternary salt
...
10- Elimination and Condensation Reaction
Elimination Reactions of Halogenoalkanes
-

-

-

-

The reactions of halogenoalkanes with hydroxide ions provide an example of how altering
the reaction conditions can cause the same reactants to produce completely different
products
With dilute sodium hydroxide solution the OH- ion acts as a nucleophile and substitution
occurs to produce an alcohol
o HO-: + R – Br  R – OH + BrWith hot alcoholic sodium hydroxide solution elimination occurs and an alkene is formed
o CH3CH2Br + OH-  C2H4 + H2O + BrIn this reaction the hydroxide ion reacts as a base
...
11- Condensation Polymerization and Reaction Pathways
Condensation Polymerization
-

Involves the reaction between two molecules to eliminate a smaller molecule
If each of the reacting molecules contain two functional groups that can undergo
condensation, then the condensation can continue to form a polymer
Amines can also condense with carboxylic acids to form an amide link (peptide bond)

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