Notesale: Turn your study into money

Document Preview

Extracts from the notes are below, to see the PDF you'll receive please use the links above

---

Chemistry
Unit Five

Key Definitions
• Entropy  The quantitative measure of the degree of disorder in a system
...

• Enthalpy change of hydration  One mole of gaseous ions forming aqueous
ions under standard conditions
...

• Enthalpy change of neutralisation  One mole of water is formed from a
reaction with an acid and base under standard conditions
...


Rate of Reaction
•
•

Rate of reaction measures 
How fast reactants are used up and how fast products are made
...


•
•
•

As the reaction proceeds, concentration decreases, leading to 
Fewer collisions per second between reactant particles
...


•

Rate of reaction is determined by measuring the concentration of a substance at time
intervals 
Rate is equal to the slope of the curve
...


•
•
•

Initial rate  the change in concentration of a reactant, of product, per unit time at the start
of the reaction when t=0
...

• pH changes by using a pH meter
...

• The loss in mass of reactants
...

• A colour change
...


•
•
•

First Order 
Rate (s-1) proportional to [B]1
If [B] increases by 2, the rate increases by 2
...


•
•
•

Third Order 
Rate (mol-2 dm6 s-1) proportional to [D]3
If [D] increases by 2, the rate increases by 23 = 8
...

Overall order  sum of individual orders
...

• Zero order 
• Concentration decreases at a constant rate
...

• First order 
• Constant half-life
...

• Second order 
• Concentration decreases rapidly, but the rate of
decrease then slows down
...


Half-Lives
• Exponential decay 
• Many natural processes have a constant half-life – value halves every half-time
...

• Used to measure element stability and in radio-carbon dating
...


•
•

Initial rate, t=0, can be determined by drawing a tangent
...

Example: precipitate appearance, solid disappearance or colour change
...

Carry out a series of clock reactions, varying each reactant in turn – plot graph of
initial rate against initial concentration
...

• Example equation: Rate = K [A]2 [B]
• Fast reaction  large value of K
...

• If the rate increases with increasing temperature when the concentrations are
the same, then the rate constant must increase with temperature
...

• For many reactions, the rate doubles every 10°c
...

• The rate of the reaction is dominated by the slowest step
...

• If a reactant is involved in the rate-determining step, it will be in the rate
equation
...

• The rate-determining step must be followed by further “fast” steps
...


Equilibrium Constant
• A dynamic equilibrium is established in a closed system when 
• The rate of the forward reaction is the same as the reverse reaction
...

• Equilibrium law  the relative proportions of reactants and products present
at equilibrium
...


• Homogeneous equilibrum  all
species making up the reactants and
products are in the same physical states
...


Equilibrium Position
• Magnitude of Kc indicates the extent of a chemical reaction
...

• Kc > 1
• Reaction is product-favoured  equilibrium lies to the right
...

• Increase in temperature shifts equilibrium in the endothermic direction (+ΔH)
...

• Shifts in equilibrium are controlled by Kc
...


Equilibrium Constant & Rate Constant
•
•
•

Pressure 
If pressure changes, equilibrium will restore itself
...


•
•
•

Concentration 
If concentration changes, equilibrium will restore itself
...


•

Catalysts speed up the rate of a chemical reaction, but do not change the position
of equilibrium
...


•
•
•
•

Equilibrium constant  compares the concentrations of reactants and products
present at equilibrium
...

Using compromise conditions increases the value of K enough to get a reasonable
yield, without Kc decreasing the equilibrium yield too much
...

• Bronsted-Lowry base  proton acceptor
...

• Lewis base  electron-pair donor
...

• Arrhenius base  dissociates when dissolved in water to form hydroxide ions
...

• Example: HCl + aq  H+ + Cl-

• Monobasic acid  can release one proton
...

• Tribasic acid  can release three protons
...

• Acid + Metal  Salt + Hydrogen
...


Conjugate Acid-Base Pairs
• The release of a proton from an acid takes place when an acid is added to
water
...

• Example: HCl + H2O  H3O+ + Cl• Acid-Base pair  a pair of two species that transform into each other by gain or
loss of a proton
...

• A high pH value  a small concentration of H+
...


•
•
•

Strong acids dissociate completely in aqueous solution
...


•
•
•

Weak acids only partially dissociate
...

Equilibrium lies well to the left
...


•
•

Weak acid: HA ↔ H+ + A-
...


• pKa = -log Ka
• Ka = 10-pKa

Ionisation of Water
• Pure water ionises into H+ and OH- ions
...

• Equilibrium lies well to the left
...


• Kw controls the balance between [H+ ] and [OH-] in all aqueous solutions
...

•
•
•
•

NaOH is a strong base and completely dissociates
...

NH3 is a weak base and only partially dissociates
...


• Strong bases in solution are alkalis
...
g
...


•
•
•
•

A buffer solution is a mixture of 
A weak acid, HA
...

A buffer can be made from a weak acid and the salt of a weak acid
...

The salt dissociates completely: CH3COONa  CH3COO- + Na+
...

Therefore, the concentration of H+ ions is very small
...

The weak acid (HA) removes added alkali
...


Buffers – Marking Points
1
...


2
...


When acid is added  Equilibrium moves to the left
...


The protons from the acid react with the conjugate base (A-) to form HA
...


When alkali is added  Equilibrium moves to the right
...


The OH- ions react with the H+ of the buffer solution to form water
...


pH Values of Buffer Solutions
• Method One 


• [HA] (at equilibrium) = [HA] (undissociated)
...

H + = protons
...


Neutralisation – Titration Curves
•

Equivalence point  the point in a titration at which the volume of one solution
has reacted exactly with the volume of the second solution
...

Within 1-2 cm3 of the equivalence point, pH increases more quickly
...

The equivalence point is at the centre of this vertical section
...


•

Most indicators change colour over a range of about two pH units
...

End point  the point in a titration at which there are equal concentrations of the
weak acid and conjugate base forms of the indicator – this shows a colour change
...


•
•

Lattice enthalpy key points 
Exothermic change – energy is given out when ionic bonds are formed from gaseous
ions
...

Covalent substances do not have lattice enthalpies
...


•
•
•

•

Hess’ law  if a reaction can take place by more than one route and the initial and
final conditions are the same, the total enthalpy change is the same for each route
...

It is usually exothermic for ionic compounds
...

It is usually exothermic for ionic compounds
...

Endothermic as bonds are broken
...

Endothermic
...

Exothermic
...

Exothermic
...

Endothermic – electron repelled by 1- ion
...

• This is known as the datum line
...

• Exothermic  ΔH points downwards
...


Born-Haber Cycles
(Formation & Lattice Enthalpy)
•

Formation = Atomisation (for both elements) + Ionisation + Electron Affinity + Lattice Enthalpy
...


2nd Ionisation

1st

EA

2nd EA

1st Ionisation
Lattice Enthalpy
Atomisation
Atomisation
Formation

Born-Haber Cycles
(Enthalpy Change of Solution)
• Example:

Enthalpy Change of Solution
• Standard enthalpy change of solution  one mole of a compound is dissolved
in water under standard conditions
...

•
•
•
•

Ionic solid dissolves 
Breakdown of ionic lattice into gaseous ions
...

The value of enthalpy change depends on the balance between these two
processes
...

• Hydration  hydration of gaseous ions – they bond with water molecules
...


Hydration & Lattice Enthalpies
•

Lattice enthalpies  large exothermic values mean a large electrostatic attraction
...

As the ionic radius increases 
Attraction between ions decrease
...

Ionic charge 
Most negative lattice enthalpies are those with small, highly charged ions
...


•
•
•
•
•
•

Hydration enthalpies 
Ionic size 
As ionic radius decreases, enthalpy is more exothermic
...

Ionic charge 
As ionic charge increases, it has a greater attraction for water molecules – results in
a more exothermic reaction
...


•

All substances possess some degree of disorder because particles are always in
constant motion  S is always positive
...

Entropy can tell us how 
A gas spreads through a room, heat from fire spreads through a room, ice melts in a
warm room and how salt dissolves in water
...

Entropy increases from solid to liquid to gas
...

ΔS is positive  change makes a system more random
...


•
•

Free Energy
• Spontaneous processes lead to lower energy and increased stability 
• Many exothermic reactions take place spontaneously at room temperature –
enthalpy content decreases an excess energy is released to the surroundings –
increases stability
...


ΔG

Make sure to change ΔS from
joules to kilojoules
...


Redox
• Oxidation  loss of electrons
...

• Oxidising agent  is reduced and accepts electrons
...


•
•
•
•
•

Balancing redox equations 
Complete half equations
...

In acidic conditions, add H+
...


Cells & Half Cells
•

Electrochemical cells control the electron transfer to produce electrical energy – this
happens in cells and batteries
...

Simplest half cell  metal in aqueous solution of ions
...


•

An electrochemical cell is made by connecting two half cells with different
electrode potentials 
One releases electrons
...

Voltmeter  measures difference in electrode potentials
...

Voltmeter has high resistance and is used to minimise the flowing current
...

A simple salt bridge can be made from filter paper soaked in an aqueous solution of
an ionic compound that does not react with either half cell
...

• Example: 2H+ (aq) + 2e- ↔ H2 (g)
• An inert platinum electrode is used
...

H2 (100kPa) is used
...


•
•
•
•

Half cells can use the same metal in different oxidation states
...

An inert platinum electrode is used
...
m
...
)  voltage produced by a cell when no current flows
...
m
...


•

The tendency of different half cells to release or accept electrons is measured and
compared when the half cells are combined separately with a hydrogen half cell
...
m
...
of a hydrogen half cell is called the standard electrode
potential (0V)
...
m
...
of a half cell compared with a standard
hydrogen half cell, measured at 298 K with solution concentrations of 1 M and a gas
pressure of 100 kPa
...
m
...
between two half cells under standard conditions –
this is the difference between the standard electrode potentials of each half cell
...


Storage & Fuel Cells
• Electrochemical cells are used as modern-day cells and batteries
...

• Rechargeable cells – the chemicals in the cell react, providing electrical energy
– recharging reverses the cell reaction
...


Storage & Fuel Cells
• Fuel cells 
• The hydrogen-oxygen fuel cell uses energy from the reaction of a fuel with
oxygen to create a voltage
...

• Fuel cells can operate continuously so long as the fuel and oxygen continue to
flow into the cell
...

• Hydrogen –Oxygen Fuel Cell Equations 
• H2 + 2OH-  2H2O + 2e• ½ O2 + 2H2O +2e-  2OH• Overall Equation: H2 + ½ O2  H2O

Hydrogen for the Future
• Fuel-cell vehicles (FVC’s) 
• Uses hydrogen gas or hydrogen-rich fuels
...
g
...

•
•
•
•

Using methanol instead of hydrogen gas 
Liquid fuel is easier to store than gas
...

Carbon dioxide is waste product – greenhouse gas
...

• Greater efficiency – petrol engines are less than 20% efficient in converting
chemical energy – hydrogen fuel-cell vehicles are 40-60% efficient
...

Can be adsorbed onto the surface of a solid material
...


•
•

Limitations of hydrogen fuel-cells 
Large-scale storage and transportation – cost-effective and energy-efficient
infrastructure is needed
...

Current adsorbers and absorbers have a limited lifetime
...

Fuel-cells use toxic chemicals in their production
...

Logistical problems in the handling and maintenance of hydrogen systems
...


Transition Metals
• Transition metals  D-block elements that form at least one ion with an
incomplete D-subshell
...

• Sc3+ ion  D-orbitals are empty
...

• In chromium and copper, it is suggested that electron repulsions between outer
electrons are minimised
...

• Transition element atoms lose electrons to form positive ions – lose 4S
electrons before 3D
...


Transition Metals

Properties of Transition Metal Compounds
•
•
•
•

Transition metals 
Shiny, high density, high melting and boiling points
...

These compounds form coloured solutions when dissolved in water and
frequently catalyse chemical reactions
...


• Coloured compounds 
• When white light passes through a solution containing transition metal ions,
some wavelengths are absorbed
...

• Colour in inorganic chemistry is linked to partially filled D-orbitals
...

Transition metal ions have the ability to change oxidation states – bind to reactant, forming
intermediates as part of a chemical pathway with a lower activation energy
...

Contact process 
2SO2 (g) + O2 (g) ↔ 2SO3 (g)
Use in the manufacture of H2SO4
...

Hydrogenation of alkenes 
Catalyst  nickel metal – lowers temperature and pressure needed to carry out the reaction
...


Precipitation
• Precipitation reactions 
• Soluble ions, in separate solutions, are mixed together to produce an insoluble
compound
...


• Co2+ (aq) + 2OH- (aq)  Co(OH)2 (s)
• Pink solution to blue precipitate (turns beige in presence of air)
...

• Fe3+ (aq) + 3OH- (aq)  Fe(OH)3 (s)
• Pale yellow solution to rusty-brown precipitate
...


•
•

Complex ion  central metal ion surrounded by ligands
...

Coordinate bond  one of the bonded atoms provides both electrons for the shared pair
...


•
•

•

Overall charge of a complex ion  sum of individual charges of the transition metal ion and
ligands
...

Some are neutral – others are negatively charged
...


•
•
•

Complex ions are most commonly octahedral in shape – central ion with six coordinate bonds
...

Four of the ligands are in the same place, one ligand above and one below – bond angle of
90°
...

• Some octahedral complex ions contain two different ligands – four of one and
two of another – can exist as cis and trans isomers
...

• Trans – the two identical ligands at opposite corners – 180° to one another –
the other four opposite each other in pairs
...

Cis – the identical ligands next to each other
...


Bidentate & Multidentate Ligands
• Bidentate ligands  can donate two lone pairs of electrons, on separate atoms,
to the central metal ion to form two coordinate bonds
...

• The complex must contain two identical bidentate ligands and two identical
monodentate ligands
...

• Trans – the identical ligands are opposite each other
...

• Optical isomerism Requirements 
• Complex with three bidentate ligands
...


Ligand Substitution in Complexes
•

Ligand substitution  reaction in which one ligand in a complex ion is replaced by
another ligand
...

Small amounts of ammonia  Cu(OH)2 forms (blue precipitate)
...


•
•
•
•
•

Copper(II) & Concentrated Hydrochloric Acid 
[Cu(H2O)6]2+ + 4Cl- ↔ [CuCl4]2- + 6H2O
Pale blue solution to green to yellow
...

Chloride ligands are larger than water ligands – have stronger repulsions – so fewer
chloride ions fit around the central metal ion – forms a tetrahedral
...


Haemoglobin
• Haemoglobin  complex protein composed of four polypeptide chains
...

• When Fe2+ binds to oxygen it turns red in colour
...

One coordinate bond to the protein globin
...


• Carbon monoxide can bind more strongly than oxygen
...

• Ligand substitution  carbon monoxide replaces oxygen – this is irreversible
...


Stability Constants
• The stability constant (K stab)  the equilibrium constant for an equilibrium
existing between a transition metal ion surrounded by water ligands and the
complex formed when the same ion has undergone a ligand substitution
...

• Water is left out of the equation as all species are dissolved in water – remains
constant
...

• Complex ions with high stability constants are more stable than those with
lower values
...


Redox Titration – Iodine & Thiosulfate
•
•

2S2O3 2- (aq) + I2 (aq)  2I – (aq) + S4O6 2- (aq)
Forms a tetrathionate ion
...

This titration can be used to determine the concentration of an oxidising agent that
reacts with iodide ions to produce iodine
...

They likely want you to work out the concentration, mass, percentage mass etc
...

To get there, just start at the final equation, use the numbers they give you and
work backwards
...

Always double check what they’re asking for once you have an answer

---