CHAPTER 1 STATES OF MATTER

There are three states of matter: solid, liquid and gas.
State
Volume
Density
Shape
Fluidity
Solid
Fixed volume
High
Definite shape
Doesn’t flow
Liquid
Fixed volume
Moderate to
No definite
Generally flows
high
shape, as it
easily
takes the shape
of the container
Gas
No fixed
Low
No definite
volume, as it
shape, as it
expands to fill
takes the shape
the container
of the container
CHANGES IN STATE
Flows easily

At a certain temperature, a liquid becomes hot enough to become a gas within the liquid and not
just at the surface. This process is called boiling. The specific temperature that it occurs at is called
the boiling point.

Volatile is a term used to describe a liquid that evaporates easily. The property of how easily a liquid
evaporates is called volatility.
PURE SUBSTANCES

A pure substance is something that consists of only one substance without containing any
contaminating impurities, and it has definite melting and boiling points.

Impurities affects the value of the melting or boiling point of a substance. An impure substance
melts or boils at a range of temperatures instead of a precise point. For example, seawater freezes
at a temperature below the freezing point of pure water (0 degrees) and boils at a temperature
above the boiling point of pure water (100 degrees).
HEATING AND COOLING CURVES

A heating curve

A cooling curve

Temperature stays constant during change of state, as energy is either absorbed or given out
during change of state.

If temperature does not stay constant during change of state, the substance is not pure.
CHAPTER 2 ATOMIC STRUCTURE

Elements are substances made up of just one type of atom.

Compounds are substances formed by chemical combination of two or more elements.

Atoms consist of three subatomic particles: protons, neutrons and electrons.
Subatomic particle
Relative mass
Relative charge
Location in atom
Proton
1
+1
In nucleus
Neutron
1
0
In nucleus
Electron
1/1840 (negligible)
-1
Outside nucleus

Structure of an atom (helium atom as an example) and explanation on atomic and mass
number:
Tip: Remember, the bigger number is always the mass number! (also, proton number also
tells the number of electrons as well, so protons = electrons)
ISOTOPES

Isotopes are atoms of the same element that have the same proton number but different
nucleon number, so they have different numbers of neutrons in their nuclei.

The difference between isotopes of the same element is just the number of neutrons in the
atoms. They all have the same properties, as they have the same number of outer shell
electrons. The isotopes are referred to using their mass number, for example, carbon-12,
carbon-13, carbon-14.
ELECTRONS IN SHELLS

Electrons are in orbit around the nucleus, and the orbits are called electron shells. The first shell can
only hold maximum two electrons, and the second shell can hold maximum eight electrons.

Electronic configuration is the electron arrangement that can be given simply in terms of numbers.
For example, the electronic configuration of sodium (11 electrons) is 2, 8, 1.
GROUP AND PERIOD NUMBERS IN THE PERIODIC TABLE

The number of outer shell electron in an atom indicates the group number for the element in the
Periodic Table.

The number of occupied shells in an atom indicates the period (row) number of the element in the
table.

Group 1 is called the alkali metals.

Group 7 is called the halogens.

Group 8 is called the noble gases (stable and unreactive, due to always having full outer shells).
CHAPTER 3 CHEMICAL BONDING

Molecular compounds: atoms are bonded together (sharing of electrons). Ex: water, ammonia,
methane, etc

Ionic compounds: ions (charged atoms) are held together in a regular structure. Ex: sodium chloride
COVALENT BONDING

Covalent bonding is chemical bonding formed by the sharing of one or more pairs of electrons
between two atoms. (between non-metal and non-metal) Ex: water and ammonia

Physical properties of covalent compounds:
 Low melting and boiling points and are often liquids or gases at room temperature
 Poor electrical conductivity
IONIC BONDING

Ionic bonding is a strong electrostatic force of attraction between oppositely charged ions. (between
metal and non-metal) Ex: sodium chloride, magnesium oxide

Physical properties of ionic compounds:
Properties
Reasons
High melting and boiling points
Ions are attracted to each other by strong
electrostatic forces. Large amounts of energy
are needed to separate them.
Crystalline solids at room temperature
There is a regular arrangement of the ions in
a lattice. Ions with opposite charge are next
to each other.
Often soluble in water
Water is attracted to charged ions and
therefore many ionic solids dissolve.
Conduct electricity only when molten or
In the liquid or solution, the ions are free
dissolved in water (not when solid)
move about. They can move towards the
electrodes when a voltage is applied.
GIANT IONIC LATTICE STRUCTURES

Ionic compounds form lattices consisting of positive and negative ions (giant ionic lattice).

Ionic crystals are hard but much more brittle than other types of crystal lattice, because pushing one
later against another brings ions of the same charge next to each other. The repulsion forces the
layers apart.
GIANT COVALENT STRUCTURES

Diamond and graphite
 Both are made up entirely of carbon, but in different structures.
 Properties of diamond and graphite:
Diamond
Properties
Appearance Colourless,
Graphite
Reason
Uses
Properties Reason
-
In jewellery Dark
transparent
and
crystals
ornamental shiny
that
objects
solid
Uses
-
-
There
Pencils
grey,
sparkle in
light
Hardness
Hardest
Each
In drill bits,
Soft – the
natural
carbon
diamond
layers can are weak
and
substance
atom
saws and
slide over
lubricants
forces of
forms
glass
each
attraction
strong
cutters
other –
between
covalent
and solid
the
bonding
has a
layers.
that
slippery
extends
feel
throughout
the whole
structure
Electrical
Does not
All the
conductivity
conduct
outer
electricity
-
Conducts
There
As
electricity
are free
electrodes
electrons
electrons
and for
of the
not used
the
atoms are
for
brushes in
used to
covalent
electric
form
bonding
motors.
covalent
by the
bonds, so
atoms in
there are
the
no free
layers.
electrons.
METALLIC BONDING

Structure of metals:
Metals ions surrounded by a sea of delocalised electrons that are free to move about.

Physical properties of metals:
 High melting and boiling points. A large amount of energy is needed to overcome the strong
force of attraction between the metal ions and the delocalised electrons.
 Good conductors of electricity. The mobile electrons can move through the structure,
carrying the current.
 Easily bent and shaped (malleable) or stretched into wires (ductile). The positive ions in a
metal are arranged in layers. When a force is applied, the layers can slide over each other.
CHAPTER 4 CHEMICAL FORMULAE AND EQUATIONS (very simplified)
IONIC EQUATIONS

In chemical equations, only some of the ions present actually change their status – by changing
either their bonding or their physical state. The other ions present are simply spectator ions – they
are there in both reactants and products and do not take part in the reaction. (Tip: only write the
ions of stuff that’s (aq)!)

Example:
CHAPTER 5 CHEMICAL CALCULATIONS

One mole of a substance contains 6.02 x 10^23 (the Avogadro constant) atoms, molecules or
formula units, depending on the substance considered.

Calculation triangle for relating number of of substance to mass:
LIMITING REACTANTS

Using the knowledge of mol, we can see how much product we can expect from particular amounts
of reactants.

In a reaction, one of the reactants may be present in excess (there’s more of it than needed), so
some of this reactant will be left over after the reaction. The limiting reactant is the one that is not in
excess. A reaction stops when the limiting reactant is used up.

Steps to find which reactant is limiting:
 First, work out the numbers of moles of each reactant involved.
 Then, divide each of them by its balancing number (coefficient) in the balanced symbol
equation.
 The smallest number indicates the limiting reactant.
PERCENTAGE YIELD AND PURITY

A reaction may not always yield the total amount of product predicted by the equation.

This can be expressed as the percentage yield for a particular experiment.

The percentage purity of a chemical product can also be calculated in a similar way to the
percentage yield.
CALCULATIONS INVOLVING GASES

The molar gas volume of any gas is 24 dm^3/mol, and it is the same for every gas.

Calculation triangle for calculating moles of gases:
CONCENTRATION OF SOLUTION

This is either mass concentration (g/dm^3) or molar concentration (mol/dm^3).

Calculation triangle for working with solution concentrations in dm^3:
CHAPTER 6 ELECTROCHEMISTRY

Electrolysis is the breakdown of ionic compounds, molten or in aqueous solution, by the use of
electricity.

Electrolysis apparatus example:

Positive ions (either metal ions or hydrogen ions) will move towards the cathode, and they are
known as cations.

Negative ions (non-metal ions) will move towards the anode, and they are known as anions.
ELECTROLYSIS OF MOLTEN IONIC COMPOUNDS

When a molten ionic compound is electrolysed:
 The metal is always formed at the cathode (the negative electrode).
 The non-metal is always formed at the anode (the positive electrode).

Different reactions take place at the two electrodes, and these reactions are represented by halfequations.

For example, the electrolysis of molten sodium chloride (NaCl):
 Cathode reaction: Na^+ + e  Na
 Anode reaction: 2Cl^-  Cl2 + 2e
ELECTROLYSIS OF SOLUTIONS

In a solution, there’s not only just the ions of the electrolyte, but also water ions.

Mantra for electrolysis of solutions:
1. Less reactive positive ions will discharge.
2. OH^- will discharge
3. Unless there’s “concentrated” and “halides”

For example, the electrolysis of concentrated sodium chloride solution:
 Cathode reaction: 2H^+ + 2e  H2 (hydrogen is less reactive than sodium)
 Anode reaction: 2Cl^-  Cl2 + 2e (there’s concentrated and halides)

For overall reactions, you must first balance the electrons in both half-equations, which can be done
by multiplying one half-equation with a number to balance. Then, combine both equations to form
an overall reaction.
ELECTROLYSIS OF COPPER(II) SULFATE WITH COPPER ELECTRODES

Electrodes that are not inert, such as copper, can be used for electrolysis, instead of the inert ones
(platinum, graphite and carbon).

In this electrolysis:
 The cathode gains mass as copper is deposited on the electrode.
Cathode reaction: Cu^2+ + 2e  Cu
 The anode loses mass as copper dissolves from the electrode.
Anode reaction: Cu  Cu^2+ + 2e

The idea of dissolving anodes is useful in other processes, such as electroplating.
ELECTROPLATING

Electroplating is a process of electrolysis where a metal object is coated (plated) with a layer of
another metal.

In electroplating:
 The cathode is the object to be plated.
 The anode is made of the metal being used to plate the object.
 The electrolyte is a salt of the same metal used for the anode.

Example of electroplating (silver):
HYDROGEN AS A FUEL

A much more efficient way of changing chemical energy into electrical energy is by using a fuel cell.
Such as cell operates continuously, with no need for recharging.

A hydrogen-oxygen fuel cell uses the reaction between hydrogen and oxygen:
2H2 + O2  2H2O

This reaction releases a large amount of energy, with water being the only product. Hydrogen can
be regarded as a non-polluting fuel, as it is non-toxic and does not emit carbon dioxide.

However, it requires a large fuel tank, hydrogen is hard to store, and its currently expensive.
CHAPTER 7 CHEMICAL ENERGETICS

There are two types of reactions: Exothermic and endothermic

Exothermic reaction:
 Releases heat/energy
 Increases the temperature if the surroundings
 Has a negative enthalpy change (heat change during course of reaction)
 Energy of the product is lower than reactant
 Energy released for bond making is greater than energy absorbed for bond breaking
 Reaction pathway diagram for exothermic reaction:

Endothermic reaction:
 Absorbs heat/energy
 Decreases the temperature of the surroundings
 Has a positive enthalpy change
 Energy of the product is higher than reactant
 Energy released for bond making is less than the energy absorbed for bond breaking
 Reaction pathway diagram for endothermic reaction:
BOND BREAKING AND BOND MAKING

To find the enthalpy change of a reaction, you must find the total energy needed for bond breaking
and bond making.

To find the values, you must get the bond energies of each bond in the reaction and total them
together.

Then, you minus the bond making value from the bond breaking value.

BB – BF (BABY – BOYFRIEND)
CHAPTER 8 RATES OF REACTION

Factors that affect the rate of a reaction:
1. SURFACE AREA OF SOLIDS
Reactions involving solids take place on the surface area of the solids. A solid has a much larger
surface area when it is powdered than when in large pieces, as more of the surface is exposed,
so they are in greater contact with other reactants and are more likely to collide and react.
Larger surface area – faster rate of reaction
2. CONCENTRATION OF SOLUTIONS
A more concentrated solution means that there are more reactant particles in a given volume.
Therefore, collisions will occur more often, and the higher the chance the particles have to react.
Higher concentration – faster rate of reaction
3. PRESSURE OF GASES
Increasing the pressure of gases pushes the gas particles closer together, so collisions are
more frequent and they can react more readily.
Higher pressure – faster rate of reaction
4. TEMPERATURE
When temperature is raised, the particles have more kinetic energy and move faster. This
means that collisions will happen more often and with more energy, which gives more chance of
reaction and bond breaking.
Higher temperature – faster rate of reaction
5. CATALYST
A catalyst is a substance that increases the rate of a chemical reaction and remains chemically
unchanged at the end of the reaction. It does this by reducing the amount of energy needed to
break the bonds (activation energy).
Reaction pathway diagrams for a reaction with and without a catalyst:
Rate of reaction graph with changes:
CHAPTER 9 REVERSIBLE REACTIONS AND EQUILIBRIUM

Reversible reaction: a chemical reaction that can go either forwards or backwards, depending on
the conditions.

Example of format: A + B⇌C + D
CHEMICAL EQUILIBRIUM

When two chemical reactions, one the reverse of the other (reversible reaction), takes place at the same
time, where the concentrations of the reactants and products remain constant because the rate of the
forward reaction is equal to the backward reaction.

Equilibrium only happens in a closed system, a system where none of the reactants or products can escape
the reaction mixture or the container where the reaction is taking place.

The position of equilibrium tells us how far a reaction has gone in favour of reactants or products. If the
concentration of products is greater, then we say the position of equilibrium is to the right – it favours the
products.

So, the four particular features of an equilibrium reaction are:
 It is dynamic: reactants are continuosly being changed to products and products are continuosly being
changed back to reactants.
 The forward and backward reactions occur at the same rate.
 The concentrations of reactants and products remain constant.
 It requires a closed system.
TEMPERATURE ON THE POSITION OF EQUILIBRIUM


High temperature favours endothermic reactions. So, the reaction will move in the endothermic direction.
Lower temperature favours exothermic reactions. So, the reaction will move in the exothermic direction.
PRESSURE ON THE POSITION OF EQUILIBRIUM




Only happens for gases.
Increased pressure moves the equilibrium to the side with less molecules (mols).
Decreased pressure moves the equilibrium to the side with more molecules (mols).
If both sides have the same mols, no change in position of equilibrium.
CONCENTRATION ON POSITION OF EQUILIBRIUM


Product increase (+)  move to the left side
Product decrease (-)  move to the right side


Reactant increase (+)  move to the right side
Reactant decrease (-)  move to the left side
HABER PROCESS

Formula:

Le Chatelier’s principle: When a change is made to the conditions of a system in dynamic equilibrium, the
system moves to oppose that change.
How the reaction system is changed to produce more ammonia (move the equilibrium to the right):
 Changing the pressure: More pressure shifts the system to the side with less molecules, so it will shift
to the right (4 mols on the left side and 2 mols on the right side). So, high pressures will increase the
yield of ammonia (20000 kPa or 200 atm).

 Changing the temperature: Forward reaction is exothermic, so backward reaction is endothermic. So,
higher temperatures will reduce the ammonia produced, as it favours endothermic reactions.
Lowering the temperature will favour ammonia production. However, the rate of reaction will be so
slow as to be uneconomical, so an optimum temperature is used (450 degrees celcius).
 Reducing concentration of ammonia: If the system was at equilibrium and then some of the
ammonia was removed, more ammonia would be produced to replace what had been taken away.
 Use of a catalyst: A catalyst can be used to increase the rate of reaction. Finely divided iron is used
as the catalyst.

So, the essential conditions used in the Haber process are:
 N2 and H2 are mixed in a ratio of 1:3
 An optimum temperature of 450 degrees celcius
 A pressure of 20000 kPa (200 atm)
 A catalyst of finely divided iron
CONTACT PROCESS


Used to produce sulfuric acid.
Main reaction that converts sulfur dioxide (SO2) to sulfur trioxide (SO3) is reversible:

Essential conditions used in the Contact process:
 An optimum temperature of 450 degress celcius
 A catalyst of vanadium(V) oxide (V2O5)
 An operating pressure of 200 kPa
Steps in the contact process:
1) Sulfur burnt to form sulfur dioxide.
S (s) + O2 (g)  SO2 (g)
2) Gases mixed and cleaned by electrostatic precipitation
3) Mixture of gases reacted
2SO2 (g) + O2 (g)  2SO3 (g)
Yield: 98% SO3
4) SO3 dissolved in 98% H2SO4

5) Concentrated sulfuric acid (oleum) diluted when needed.
FERTILISERS


Ammonium salts and nitrates can be used as fertilisers.
NPK fertilisers are used to provide the three important elements for plant growth:
 Nitrogen (N): important for producing the proteins needed for plant growth
 Phosphorus (P): important for healthy roots
 Potassium (K): important for the production of flowers and fruit.
BONUS: CHEMICAL TEST FOR WATER




Presence of water can be detected by using anhydrous copper (II) sulfate or anhydrous cobalt (II) chloride.
Anhydrous copper (II) sulfate: from white to blue
Anhydrous cobalt (II) chloride: from blue to pink
CHAPTER 10 REDOX REACTIONS
A redox reaction is a reaction involving both reduction and oxidation.
Oxidation:
 A reaction in which oxygen is added to an element or compound
 A reaction involving the loss of electrons from an atom, molecule or ion
 A reaction in which the oxidation state (charge) of an element is increased
Reduction:
 A reaction in which oxygen is removed from a compound
 A reaction involving the gain of electrons by an atom, molecule or ion
 A reaction in which the oxidation state of an element is decreased
Examples:
Gain and loss of oxygen
Change in oxidation number
Gain and loss of electrons



Rules for working out the oxidation number of an element:
 The oxidation number of the uncombined element is zero (0)
Ex: H2 = 0 and Cl2 = 0
 The oxidation number of an ion is the charge in the ion
Ex: Zn2+ = +2 and O2- = -2
 The oxidation number of hydrogen is +1 (H = +1)
 The oxidation number of oxygen is -2 (O = -2)
 The sum of the oxidation numbers in a compound is zero (0)
Ex: H2O = 0, H=+1 and O=-2, so [2 x (+1)] + [1 x (-2)] = 0
 The sum of the oxidation numbers in a compound ion is the charge on the ion
Ex: Manganate (VII) ion  (MnO4-) = -1, Mn = +7, so [1 x (+7)] + [4 x (-2)] = -1
Tip: Roman numerals are used to indicate the oxidation number of an element in the formula of a
compound.
Oxidising agent: a substance that oxidises another subtance and is itself reduced during the reaction.
Reducing agent: a substance that reduces another substance and is itself reduced during the reaction.
CHAPTER 11 ACIDS AND BASES



Acid: a substance that dissolves in water, producing H+(aq) ions – a solution of an acid turns litmus red and
has a pH below 7. Acids act as proton donors (discussed later).
Base: a substance that neutralises an acid, producing a salt and water as the only products. They are mainly
insoluble in water. Bases act as proton acceptors (discussed later).
Alkali: soluble bases that produce OH-(aq) ions in water – a solution of an alkali turns litmus blue and has a
pH above 7. (Tip: All alkalis are bases, but not all bases are alkalis.)
INDICATORS


Three commonly used indicators: litmus, thymolphthalein, methyl orange
Indicator
Colour in acid
Neutral colour
Litmus
Red
Purple
Thymolphthalein
Colourless
Colourless
Methyl orange
Red
Orange
Colour in alkali
Blue
Blue
Yellow
Another commonly used indicator: Universal indicator.
Useful because it gives a range of colours (a spectrum) depending on the relative strength of the acid or alkali
added.
METAL AND NON-METAL OXIDES

There are four main types of oxides:
 Non-metal oxides:
 Acidic oxides  can react with a base (Ex: CO2, SO2, NO2, etc)
 Neutral oxides  does not react with acids or bases (Only three: NO, N2O, CO)
 Metal oxides:
 Basic oxides can react with an acid (Ex: CaO, MgO, CuO, etc)
 Amphoteric oxides can react with both acids and bases (Only three: ZnO, PbO, Al2O3)
REACTIONS OF ACIDS




Acid + metal  salt + hydrogen
Ex: Magnesium + nitric acid  magnesium nitrate + hydrogen
Mg (s) + 2HNO3 (aq)  Mg(NO3)2 (aq) + H2 (g)
Tip = Salt made depends on the acid:
 Hydrochloric acid  chloride
 Sulfuric acid  sulfate
 Nitric acid  nitrate
 Phosphoric acid  phosphate
 Ethanoic acid  ethanoate
Acid + base  salt + water
Ex: hydrochloric acid + sodium hydroxide  sodium chloride + water
HCl (aq) + NaOH (aq)  NaCl (aq) + H2O (l)
Acid + metal carbonate  salt + water + carbon dioxide
Ex: hydrochloric acid + zinc carbonate  zinc chloride + water + carbon dioxide
2HCl (aq) + ZnCO3 (s)  ZnCl2 (aq) + H2O (l) + CO2 (g)
Amphoteric oxide + acid/alkali  salt + water
Ex: Zinc oxide + hydrochloric acid  zinc chloride + water
ZnO (s) + 2HCl (aq)  ZnCl2 (aq) + H2O (l)
Zinc oxide + sodium hydroxide  sodium zincate + water
ZnO (s) + 2NaOH(aq)  Na2ZnO2 (aq) + H2O (l)
STRONG AND WEAK ACIDS


Strong acids are completely ionised/dissociated when dissolved in water, so it produces the highest possible
concentration of H+ (aq) ions in solution. (Ex: hydrochloric acid)
Weak acids are partially dissociated when dissolved in water, so it produces a low concentration of H+ (aq)
ions in solution. (Ex: ethanoic acid)
PROTON DONOR AND ACCEPTOR



In some reactions, acids provide hydrogen ions to react with hydroxide ions. In turn, the base is supplying
hydroxide ions to accept the H+ ions and form water.
A hydrogen ion (H+) is simply a proton.
So, an acid is a proton donor (it gives a proton, H+ ion, to a base), and a base is a proton acceptor (it accepts
a proton, H+ ion, from an acid).
CHAPTER 12 PREPARATION OF SALTS

Salts are compounds formed from an acid by the replacement of the hydrogen in the acid by a metal or by
ammonium ions. Essentially, when a positive (+) ion from a metal or ammonium combines with a negative (-)
ion from a non-metal. (Ex: copper (II) sulfate, potassium nitrate)
SOLUBILITY OF SALTS


All sodium, ammonium, potassium salts and nitrates are soluble.
All sulfates (SO4 2-) except Ba, Pb and Ca are soluble.


All halides (group 7) except Ag, Pb are soluble.
All carbonates (CO3 2-) except Na, K and NH4+ are insoluble.
SNAP (Soluble Salts):
 Sodium
 Potassium
 Ammonium
 Nitrates
WATER OF CRYSTALLISATION

 When crystallised, many salts contain water
that is chemically combined in the crystal
structure. Such salts are called hydrated salts.
The water chemically combined in these hydrated salts is known as water of crystallisation.
When preparing crystals of hydrated salts, we must be careful not to heat them too strongly when drying
them, as it can evaporate the water of crystallisation and turn them into anhydrous powder, not crystals.
PREPARING SOLUBLE SALTS


Method A – acid plus solid metal, base or carbonate
 Step 1: An excess of the solid is added to the acid and allowed to react. An excess of the solid is used
to make sure all the acid is used up.
 Step 2: The excess solid is filtered out.
 Step 3: The filtrate is gently evaporated to concentrate the salt solution.
 Step 4: When crystals can be seen forming, heating is stopped and the solution is left to crystallise.
 Step 5: The concentrated solution is cooled to let the crystals form. The crystals are filtered off,
washed with distilled water, and dried carefully between filter papers.
Method B – acid plus alkali by titration
 Step 1: The acid solution is poured into a burette. A known volume of alkali solution is placed in a
conical flask using a volumetric pipette. A few drops of an indicator (e.g. thymolphthalein or methyl
orange) are added to the flask.
 Step 2: The acid solution is run into the flask a few drops at a time from the burette until the
indicator just changes colour. The conical flask must also be swirled after each portion of acid to
ensure everything is mixed and the reaction is complete. The volume of acid added is noted. The
experiment is repeated without using the indicator with the same volume of acid noted run into the
flask.
 Step 3: The salt solution is evaporated and cooled to form crystals as described in method A.
PREPARING INSOLUBLE SALTS BY PRECIPITATION

Soluble salt 1 + soluble salt 2  Inhysr salt 3 + salt 4
Ex:


Soluble
Soluble
Insoluble
Soluble
Soluble salt 1 + Acid 1  Insoluble salt 2 + Acid 2
After precipitation: filter it off, wash with distilled water and dry in a warm oven.
CHAPTER 13 THE PERIODIC TABLE
Short Review


Group: No. of electrons in outer shell
Period: No. of shells
 Example  Na = 2.8.1 (Group = 1; Period = 3)
GROUP 1 (ALKALI METALS)






Soft metal
Highly reactive and stored in oil to prevent them from reacting with the oxygen and water vapour in air.
Reactivity increases down the group
Low Melting Points, Boiling Points and Densities
Melting and Boiling points decrease down the group; Densities increase down the group
Ion = +1
GROUP 7 (HALOGENS)







Non-metal
Reactivity decreases down the group
Melting and Boiling points increase down the group
Elements change from gases to solids as you go down the group
Exist as diatomic molecules (Cl2, Br2, I2)
Ion = -1; called Halides (e.g. Chloride/Cl-)
Colours of Halogens:
 Fluorine (F): Pale yellow
 Chlorine (Cl): Greenish yellow
 Bromine (Br): Reddish brown
 Iodine (I): Grey black
 Astatine (At): Black
GROUP 8 (NOBLE GASES)





Gas
Low MP, BP, Density
Full electrons in outer shell
 He  2
 Ne  2.8
 Ar  2.8.8
Unreactive
Monoatomic (not diatomic like He2, Ne2)
TRANSITION METALS





Hard metal
High MP, BP, density
Variable Oxidation number
Coloured Compounds
Often act as catalysts
CHAPTER 14 METALLIC ELEMENTS AND ALLOYS
PHYSICAL PROPERTIES OF METALS

Good conductors of electricity (due to free-moving electrons in its structure)

Good conductors of heat (due to metal atoms vibrating and colliding, transferring heat)

Malleable and ductile (the layers of metal atoms can slide over each other when force is applied)


Hard and strong (except for group 1 i.e. sodium, potassium)
Sonorous (can make a ringing sound)
USES OF METALS

Aluminium:
 Aeroplane bodies (Light and strong, with a low density)
 Overhead power cables (Low density and good electrical conductivity)
 Saucepans (Good heat conductor)
 Food containers (Resistant to corrosion, due to a stable layer of aluminium oxide which forms on its
surface, stopping the aluminium from reacting, so it can’t react with natural acids in food)

Copper:
 Electrical wiring (Very good electrical conductor and ductile)
 Water pipes (Non-toxic, strong but malleable)
 Pots and pans (Very good conductor of heat, unreactive, malleable)
 Surface in hospitals (Antibacterial properties)
ALLOYS



Alloys are mixtures of a metal with other elements.
Properties and uses of alloys:
 Brass - an alloy of copper and zinc and is much stronger than either metal. Used in musical
instruments, ornaments and door knobs.
 Bronze – made of copper and tin, used in statues, bells and machine parts.
 Mild steel – made of iron and carbon, used in car bodies
 Stainless steel – an alloy of iron and other elements, for example, chromium, nickel and carbon.
Used in cutlery as it is hard and resistant to corrosion.
 Solder – made of tin and lead, used in making electrical connections due to lower melting point.
Alloys often have properties that are different to the metals they contain. For example, they can be stronger,
harder and resistant to corrosion. These enhanced properties can make alloys more useful than pure metals.
STRUCTURE OF ALLOYS

Metals have a regular arrangement of ions. Alloys, on the other hand, has an irregular arrangement of ions,
due to the different sized atoms that make the lattice structure less regular.
CHAPTER 15 REACTIVITY OF METALS

Reactivity series of metals:
Potassium (K)
Sodium (Na)
Calcium (Ca)
Magnesium (Mg)
Aluminium (Al)
Carbon (C)*
Zinc (Zn)
Iron (Fe)
Hydrogen (H)*
Copper (Cu)
Silver (Ag)
Gold (Au)
*non-metals but important
Memory aid: Please Send Camels, Monkeys And Corny Zebras In Huge Cages (made of) Silver & Gold
REACTION OF METALS WITH COLD WATER

The more reactive metals (Potassium, Sodium and Calcium) will react with cold water to form a metal
hydroxide and hydrogen gas.

Metal + water  Metal hydroxide + hydrogen

Example: 2K + 2H2O  2KOH + H2
REACTION OF METALS WITH STEAM



Metals just below calcium in the reactivity series do not react with cold water but will react with steam to
form a metal oxide and hydrogen gas.
Metal + H2O (g)  Metal oxide + hydrogen
Example: Mg (s) + H2O (g) → MgO (s) + H2 (g)
REACTION OF METALS WITH ACID

Only metals above hydrogen in the reactivity series will react with acids. Unreactive metals below hydrogen,
such as copper, silver and gold, do not react with acids.



The more reactive the metal, the more vigorous the reaction will be.
Metal + acid  Salt + hydrogen
Example: Ca + 2HCl  CaCl2 + H2
REACTION OF METALS WITH OXYGEN

Some reactive metals, such as the alkali metals, react easily with oxygen.



Silver, copper and iron can also react with oxygen but much more slowly. Gold does not react with oxygen.
Metal + oxygen  Metal oxide
Example: 2Cu (s) + O2 (g) → 2CuO (s)
THERMAL DECOMPOSITION

Metal carbonate  Metal oxide + Carbon dioxide
All metals can react other than Sodium and Potassium.
Ex: CaCO3  CaO + CO2

Metal hydroxide  Metal oxide + steam
All metals can react except Sodium and Potassium.
Ex: Zn(OH)2  ZnO + H2 (g)

Metal nitrate  Metal nitrite + Oxygen (Only for Sodium and Potassium)
Metal nitrate  Metal oxide + Nitrogen dioxide (brown in colour) + Oxygen
 Ex: 2KNO3  2KNO2 + O2
 Ex: 2Pb(NO3)2  2PbO + 4NO2 + O2
CHAPTER 16 EXTRACTION AND CORROSION OF METALS

The methods for extraction of metals from ores depend on the postion of the metal in the reactivity series.
EXTRACTION OF IRON



Major ore of iron: Hematite (Fe2O3)
Done in a blast furnace (called that as it uses hot air blasted in to raise the temperature)
Important materials:
 Coke (Carbon/C)  to increase the temperature of the blast furnace and produce the reducing agent
(carbon dioxide)
 Limestone (Calcium carbonate/ CaCO3)  to remove acidic impurities from the ore



Zone 1
 Coke burns in the hot air forming carbon dioxide. This reaction is exothermic so it gives off
heat, heating the furnace.
 Carbon + Oxygen → Carbon dioxide
 C + O2  CO2
Zone 2
 At the high temperatures in the furnace, more coke reacts with carbon dioxide forming carbon
monoxide (this CO is the reducing agent in the reaction with iron (III) oxide, hence why coke
produces the reducing agent)
 Carbon + Carbon dioxide → Carbon monoxide
 C + CO2  2CO
Zone 3
 Carbon monoxide reduces the iron (III) oxide in the iron ore to form iron. This will melt and collect at
the bottom of the furnace, where it is tapped off.
 Iron (III) oxide + carbon monoxide → iron + carbon dioxide
 Fe2O3 + 3CO  2Fe + 3CO2
 Limestone (calcium carbonate) is added to the furnace to remove impurities in the ore. The calcium
carbonate in the limestone thermally decomposes to form calcium oxide.
 Calcium carbonate → Calcium oxide + Carbon dioxide
 CaCO3  CaO + CO2
 The calcium oxide formed reacts with the silicon dioxide, which is an impurity in the iron ore, to form
calcium silicate. This melts and collects as a molten slag floating on top of the molten iron, which is
tapped off separately.
 Calcium oxide + Silicon dioxide → Calcium silicate (slag)
 CaO + SiO2  CaSiO3 (slag)
EXTRACTION OF ALUMINIUM


Major ore of aluminium: Bauxite (Al2O3)
Done by electrolysis (refer to chap 6). This is because aluminium is higher than carbon in the reactivity series,
so it can’t be reduced by carbon.

After the Bauxite is purified with sodium hydroxide, the purified aluminium oxide is dissolved in molten
cryolite. This is because aluminium oxide has a melting point of over 2000 °C which would use a lot of energy
and be very expensive. The cryolite lowers the melting point of the solution without interfering with the
reaction.

The mixture is placed in an electrolytic cell, made of steel and lined with graphite. The graphite lining acts as
the cathode, with several large graphite blocks as the anodes.

At the cathode:
 Aluminium ions gain electrons (reduction). This molten aluminium forms at the bottom of the cell. It
is siphoned off from time to time and fresh aluminium oxide is added to the cell.
 Al3+ (l) + 3e- → Al (l)

At the anode:
 Oxide ions lose electrons (oxidation), so oxygen is produced.
 2O2- (l) → O2 (g) + 4e-
 At the high temperature of the cell, the oxygen reacts with the carbon in the electrodes forming
carbon dioxide. The carbon anodes slowly burn away and have to be replaced frequently.
 C (s) + O2 (g)  CO2 (g)
CORROSION OF METALS

When a metal is attacked by air, water or other surrounding substances, it is said to corrode. In the case of
iron and steel, the corrosion process is known as rusting.
RUST PREVENTION

Barrier methods
 This is coating the iron or steel with material to prevent them from coming into contact with water
and oxygen.
 Examples of this include painting, oiling and greasing, plastic coatings and electroplating (refer to
chap 6).

Sacrificial protection
 This is preventing iron from rusting by attaching blocks of a metal more reactive than iron (such as
zinc and magnesium) to the structure. The more reactive metal will lose electrons and be corroded
in preference to the iron.
 Zinc is more reactive than iron therefore will lose its electrons more easily than iron and is oxidised
more easily.
Zn → Zn2+ + 2e-

Galvanising
 This is protecting iron or steel by coating it with a layer of zinc, a more reactive metal. It’s advantage
over barrier methods is that the protection still works when the zinc layer is badly
scratched/damaged.
 While the zinc layer is unbroken, it will be a form of barrier method of protection. When the zinc
layer is damaged, it is corroded away in preference to the iron, as a form of sacrificial protection.
CHAPTER 17 CHEMISTRY OF OUR ENVIRONMENT
AIR

Air is mostly made up of nitrogen and oxygen, with other gases like noble gases and carbon dioxide present
in very small quantities.
AIR POLLUTION

Sulfur dioxide
 Sources:
o Combustion of fossil fuels containing sulfur compounds
S + O2 → SO2
o Power stations
 Dissolves in rain to form acid rain, causing corrosion to metal structures, buildings and statues made
of carbonate rocks, damage to aquatic organisms
 Can be removed by desulfurisation (removal of sulfur dioxide from the fumes of power stations)

Oxides of nitrogen
 Sources:
 Reaction of nitrogen with oxygen in the presence of high temperatures, such as in car
engines, high-temperature furnaces and when lightning occurs.
 A product of bacterial action in soil.
 It produces photochemical smog, which is linked to respiratory disease and asthma attacks.
 It also dissolves in rain to form acid rain, causing damage.
 It pollutes crops and water supplies.
 Can be removed by using a catalytic converter, which converts polluting exhaust gases from cars into
less dangerous emissions.
 Carbon monoxide + Nitrogen oxide  Carbon dioxide + Nitrogen
2CO (g) + 2NO (g)  2CO2 (g) + N2 (g)

Carbon monoxide
 Sources:
 Incomplete combustion of carbon-containing fuels such as fossil fuels.
 Ex  incomplete combustion of octane:
Octane + Oxygen  Carbon monoxide + Water
2C8H18 + 17O2  16CO + 18H2O
 Carbon monoxide is toxic to humans as it combines with haemoglobin in the blood and prevents it
from carrying oxygen.

Particulates:
 Carbon particulates are also formed as a result of incomplete combustion of fuel. An example is the
incomplete combustion of octane to form particulates:
Octane + Oxygen  Particulate (soot) + Water
2C8H18 + 9O2  16C + 18H2O
 Particulates can cause respiratory disease and cancer.

Carbon dioxide
 Sources:
 Complete combustion of carbon-containing fuels such as fossil fuels.
 Ex  complete combustion of methane:
Methane + Oxygen  Carbon dioxide + Water
CH4 + 2O2  CO2 + 2H2O
 Carbon dioxide increases global warming and climate change, as it is a greenhouse gas.
 Can be reduced by using renewable energy sources and can be removed by photosynthesis.

Methane
 Sources:
Waste gases from digestive processes of animals
Decomposition of vegetation
Bacterial action in swamps, rice paddy fields and landfill sites
 Methane is also a greenhouse gas, so increases global warming and climate change.
WATER

Presence of water can be detected by using anhydrous copper (II) sulfate or anhydrous cobalt (II) chloride.
Anhydrous copper (II) sulfate: from white to blue
Anhydrous cobalt (II) chloride: from blue to pink

The purity of water can be determined by measuring its melting and boiling point. This is because pure water
has a fixed melting point of 0 degrees celcius and boiling point of 100 degrees celcius, while impure water
doesn’t (refer to chap 1).
PURIFICATION OF WATER



The first step in purification is to remove large insoluble objects like rocks, plastic bags and branches. This is
called screening.
The next step is to filter the water to remove smaller insoluble particles.
This water is then disinfected using chlorine, in order to kill harmful bacteria that can cause disease. Only
then, is the water safe to be distributed to homes and drunken.
Organic chemistry Chapter 18
Hydrocarbons: Compounds that contain only hydrogen and carbon
1. Alkanes:

It is a saturated hydrocarbon, which all carbon-carbon bonds are single covalent bonds

General formula CNH2N+2
Molecular formula: Formula that shows the actual number of atoms present in the compound
Displayed or structural formula: Formula that shows the structure of the compound and its bonds
2. Alkenes:

Unsaturated hydrocarbons, since they have a carbon-carbon doubles bond in its structure

General formula CNH2N where N > 1

It is much more reactive than alkanes as the double bond of carbon can be broken and other
atoms can add on to it

It can be tested for its presence by mixing it with a solution of bromine with water, if there is
presence the bromine loses colour (addition reaction)
Homologous series: A family of similar compounds with similar chemical properties due to the same
functional group (eg, Alkanes, Alkenes)
Functional Group: the atom or group of atoms that gives the series its particular characteristical properties
3. Alcohols

-OH functional group

General formula CnH2n+1OH

It is saturated hydrocarbons
4. Carboxylic acids

Functional group-COOH

General formula CnHn2+1COOH

Structural
formula of the
Homologous
series:
Factors important in naming molecules such as 2-methylpropane:
-
Consider the side chain , in this case it is the methyl group (CH3)
-
Numbering the carbon atom which the side chain is connected to , count it from the smallest
number possible
Structural isomerism: Compounds that have the same molecular formula but different structural formulae,
known as structural isomers
Position isomerism : where position of functional group is moved within the molecule
5. Esters:

Where an alcohol and carboxylic acids combine and form this compound, called
esterification

Ex. CH3COOH + C2H5OH → CH3COOC2H5 + H20

The acid name is the second name of the ester, this case ethanoate

The alcohol or now alkyl is the first name, this case ethyl

To draw the displayed formula, remember that one of the molecules have to be flipped in
order to connect the two compounds, they also lose 2 hydrogen and 1 oxygen atoms
Chapter 19

Alkanes are quite unreactive and cannot take part in addition reactions, not affected by acid or alkali
Alkenes as fuels:
-
Some alkenes are obtained from fractional distillation of petroleum
-
They burn very exothermically
-
Complete combustion of alkanes
-
Incomplete combustion of alkanes, carbon monoxide is produced and can produce soot
particles
-
Substitution reaction for alkanes:
o
Alkanes can react with chlorine to produce an chloroalkane, where a chlorine atom
replaces an hydrogen atom
o
Alkane + Chlorine => Chloroalkane + hydrogen chloride
Chemistry of alkenes:

Alkenes can go through catalytic cracking where heavier fractions can be broken into smaller ones

The hydrocarbons are mixed with a catalyst and burned in high temp

It can give either a long chained alkane and a short chain alkene or two or more alkenes and
hydrogen

They combust similarly to alkanes and can take part in addition reactions such as bromination and
hydrogenation

They can go through catalytic addition of steam where ethene is added to steam and can produce
ethanol when phosphoric acid is used and 300 degrees in 6000kPa of pressure
Chemistry of ethanol:

Can be produced from where yeast produces a waste product of ethanol and carbon dioxide in
anaerobic respiration, although it is self limiting as ethanol is toxic to yeast if in high concentration

Alternatively , ethene can be hydrated with steam to produce ethanol


Ethanol are commonly used as a fuel for cars when blended with petrol, it is a renewable resource
and is a biofuel
Chemistry of carboxylic acids and esters:

Ethanoic acid can be formed from the biochemical oxidation of ethanol using bacteria

Warm acidified potassium manganate can be used also for oxidation as it is an oxidizing agent,
using a vertical condenser, the potassium manganate will turn colourless from purple

Ethanoic acid will react normally with any other reactive metals and alkaline

Ethanoic acid and ethanol can react by esterification to produce the ester ethyl ethanoate and water
with the presence of concentrated sulfuric acid.
Chapter 20
Fractional distillation of petroleum:
-
Petroleum is refined and is separated to smaller fraction
-
The bigger the chain length of the fraction, the more viscous it is and the less volatile
1. Petroleum is preheated to 350-450 degrees and pumped in to the base of the fractioning tower
2. The petroleum will boil and its vapour will rise, passing through bubble caps and cools as it rises,
the different fractions condense and cool at different temps so it will be at different heights
-
There is more demand for lighter fractions than the heavier fractions and supply doesn’t really
meet demand
Polymers:
-
Made of monomers, small repeating units joined together through polymerization
-
There are 2 types of polymers:
-
o
Homopolymers where there is only one monomer
o
Copolymers where there is multiple monomers
The alkenes’ carbon double bonds are broken so they can go through addition reaction where
other molecules join the alkenes’ structure to form a larger molecule
Condensation polymerization:
-
When two molecules join together to become a larger molecule with a side product of usually
water
-
Reaction between a carboxylic acid and an amine(NH2) forms an amide and water, a large
number of this is called a polyamide
-
Reaction between a carboxylic acid an alcohol produces an ester, a large number of this is a
polyester
-
In the polyamide , an amide link is formed between the acid and amine, combining it
together,and H2O is produced
Natural polymers:
-
Are proteins built from amino acid monomers
-
Each amino acid contain an amino group(NH2_) and a carboxylic acid group
-
The difference of each amino acids lies in the different side chains (-R) ex. Glycine has R=-CH3)
and alanine (-H)
-
When two amino acids combine, it is a condensation reaction where water is produced and a
dipeptide with an amide link
-
When repeated many times, a short polymer is formed called peptides, a longer chain is called
polypeptides or proteins
Plastics:
-
A group of polymeric materials characterized by their plasticity
-
Non-biodegradable
-
Very strong but also low density
-
Can be recycled to make items such as soft-drink bottles
-
Ester linkage joining the monomer units in polyester can be broken down by preysis and can be
re-polymerized to another polymer
-
Can be in a form of microplastics like microbeads
Additional:

A insoluble and soluble salt is formed when two soluble salts are mixed by disstation
SNAP (Soluble Salts):
 Sodium
 Potassium
 Ammonium
 Nitrates
How sulfur dioxide is obtained:

Burning sulfur with air

Roasting sulfide ores
Features of an equilibrium:

Forward and backward reaction are equal

Reactant and product are in same concentration
-
Lower temperature favours exothermic reaction because, the system counteracts the temp
decrease by releasing more energy
-
Higher temp favours endothermic reaction because the system will counteract the increase in temp
by cooling and absorbing more energy
-
Three commonly used indicators: litmus, thymolphthalein, methyl orange
Indicator
Colour in acid
Neutral colour
Litmus
Red
Purple
Thymolphthalein
Colourless
Colourless
Methyl orange
Red
Orange
Colour in alkali
Blue
Blue
Yellow
-
Substitution- Where one atom in an alkane is replaced by another atom
-
Ultraviolet light in a photochemical reaction can break down bonds ex. Chlorine
-
Addition reaction with steam means a monomer is being added with H2O
o The double bond is broken to add the OH and H ions
-
Compounds with ionic bonding have high melting point due to the electrostatic attraction between positive
and negative ions, this allows a strong attraction and bonds
-
Catalysts are substances added to a reaction to increase its rate, it remains unchanged at the end
-
Rate of reaction is highest at the start due to it having the highest concentration of the substance
-
Moles account for the total number of particles in the reaction, no need to multiply it again
-
Electrolysis is the breakdown of ionic compounds in the form of molten or aqueous by electricity
-
Particle responsible for transfer of charge in conducting wires in electrolysis is electrons
-
Particle responsible for transfer of charge the electrolyte is the ions
-
Product at the anode is the products after it loses the electrons, its not OH but the product of it which is O
-
Cryolite reduces the operating temperature and acts as the solvent for aluminium oxide to dissolve in
-
Overall reaction in hydrogen-oxygen fuel cell is 2H2 + O2=2H2O
-
Fizzling and bubbling can occur when sodium carbonate reacts with carboxylic acids
-
Metals have a sea of electrons which allows it to conduct electricity when solid
-
A chemical change can be indicated by a newly formed precipitate
-
Energy change is not the same and enthalpy change
-
Cobalt chloride is used to indicate the presence of water
-
Ammonia is the only gas that is alkaline .
-
Waste gases should not be released into the lab as it can be flammable
-
It is better to use a burette rather than a measuring cylinder as it is more accurate
-
A reaction can start if we use a pipette to add a liquid one drop at a time
-