Formulas

Chemical formula is derived either from the ions or for non-metal compounds a

description within the name

Ions are combined to cancel out overall charge subscriptions indicate multiples for an ion.

Polyatomic ions are put into brackets

Eg. Sodium Chloride Magnesium Chloride Zinc Hydroxide

Na+ Cl- �NaCl Mg2+ Cl- �MgCl2 Zn2+ OH- � Zn(OH)2

Sulfur Dioxide Sulfur Trioxide

SO2 SO3

Equation

An equation shows the path of a chemical reaction. It consists of reactants going to

products. When writing equations the correct formula is written and then it is balanced

with mole coefficients. There are a number of common reaction types:

1. - Acids + Base � Salt + Water

- Bases are -Hydroxides/-Oxides

2. Acids + Carbonates � Salt + Water + Carbon Dioxide

3. Acid + Metal � Metal Salt + Hydrogen

4. Reactions with Oxygen � Oxides/Hydrocarbon + O2 + CO2 + H2O

Matter

- Matter consists of atoms/particles that fill and available space

- There is no matter in a vacuum

- Matter can divide into: - Mixtures: Combinations of particles that are not

chemically joined

- Pure Substances: Made up of only a single type of

particles: - Elements

- Compounds

States

Matter can exist in three main states (phases): - Solids (s)

- Liquid (aq)

- Gases (g)

At an atomic level energy (heat) and movement are proportional. ie more heat � More

movement. As a substance is heated the temperature raises then stops at its melting point

as the energy is used to overcome the force between particles once liquid its temperature

increases again

Properties

Properties are specific attributes for a compound physical propertied include:

- State and Density

- Operative including colour, size, etc.

- Conductivity

- Melting Point (mp) Boiling Point (bp)

- Solubility

Chemical properties include how a substance reacts with common chemicals such as:

- Oxygen, Acids, Bases, Oxidants and Reactants

Separations of Mixtures

A mixture of substance can be ‘easily’ separated by physical means. By using differences

in physical properties of the components they can be separated.

Separation Techniques:

- Filtering insoluble particles trapped in filter paper

- Decanting - Careful pouring off top liquid layer from two different density

substances

- Dissolving - separate soluble from insoluble substances

- Evaporating/Boiling/Crystallization where a soluble solid is recovered by

evaporating the solvent liquid

- Distillation - Used to separate liquid of different boiling points

- The lower boiling point (bp) fraction is evaporated out of the

solution and condensed to recover it

- Centrifuging - Use mass differences in a spinning machine (with centrifugal

force) to separate the heaviest components to the bottom

- Chromatography - Use differences in solubility to separate a mixture in a

moving solvent

Atomic Structure

Atoms make up all things they are tiny indivisible particles consisting of a central nucleus

of protons and neutrons, surrounded by negative electrons arranged in shells

The periodic table tells us information for each element (type of atom)

- Atomic number = # of Protons

- Mass number = # of Protons + # of Neutrons

- # of Electrons = # of Protons

Isotopes are elements/atoms with a different number of neutrons so different mass

number:

Eg. - Carbon-14 (14C) & Carbon-12 (12C)

- Oxygen-16 (16O) & Oxygen-18 (18O)

- Hydrogen (1H)

Electron Configuration

Electrons are arranged in shells surrounding the nucleus. The shells fill in order from

inside first. The max number of electrons is:

1st Shell: - 2 electrons

2nd Shell: - 8 electrons

3rd Shell: - 8 electrons

4th Shell: - 18 electrons

Each shell corresponds to a period on the periodic table. The notation used is element

symbol ten the number of electrons in each shell:

Eg. H: 1

He: 2

O: 2,6

Mg: 2,8,2

K: 2,8,8,1

Ca: 2,8,8,2

Ions

An ion is a particle that has lost or gained electrons to become charged positive ions are

cations, negative ions are anions. Simple monatomic ions form from elements losing or

gaining elections to achieve the stability of a complete outer shell thus, noble gasses do

not form ions. Metals lose elections to become cations. Non-metals gain electrons to

become anions.

Eg. Mg: 2,8,2 will lose 2 outer electrons to be Mg2+: 2,8

N: 2,5 will gain 3 extra electrons to fill its outer shell N3-: 2,8

Elements in the same group from the same charge ions:

Eg. Group 1 all +1

Group 2 all +2

Group 17 all -17

Group 16 all -16

Bonding

Elements undertake chemical bonding to enable a stable electron arrangement to form

There are 3 forms of bonding available:

1. Metallic bonding occurs between metal atoms only:

A metal atom’s valence or outer shell electrons are delocalized into the

metal 3D structure

Metallic Bonding

The structure is a 3D lattice (network) of positive nuclei kernel in a ’sea’

of delocalized valence electrons.

The structure of substance leads to its properties:

- Strong substances, they have high melting points and boiling

points due to the 3D lattice structure and strong attraction between

positive kernels and surrounding electrons

- Conduct electricity due to mobile electron in structure

- Malleable (beaten into sheet) ductile (drawn into wire)

- Shiny when scratched

2. Ionic Bonding:

Occurs between metal and non-metal elements. Electrons are transferred

from metals forming cations to non-metals forming anions. Then a 3D

network of oppositely charged ions with strong electrostatic attraction

forms

Ionic Compound properties:

- High melting and boiling points due to 3D network lattice structure, it

requires high energy to overcome the electrostatic attractions surrounding

an ion. Eg. NaCl = 801°c

- Do not conduct as solids, but do as liquids/molten or aqueous solutions

as the ions are free to move and carry charge.

- Hard but brittle substance due to lattice imperfections creating cleavage

planes

3. Covalent Bonding

This occurs between non-metal elements and involves sharing electrons to

enable atoms to have access to a full valence shell for stability

Covalent Bonding can form two different types of substances:

1. Network or extended lattice structures:

- Strong bonds in 2 or 3 dimensions

- Very strong, high melting and boiling points, hard

substances

- Insoluble and usually do not conduct (except

graphite)

2. Covalent Molecules

Covalent bonding can form individual discrete molecules. These

can be non-polar or polar

Non-Polar Bonding:

As a Substance:

- Very weak force between non-polar molecules

- Very low melting and boiling points due to weak

forces

- Non-conductors - no free ions or electrons to

carry charge

- Insoluble and immiscible in water

- Individual discrete molecules up of a few covalently

bonded atoms

Polar Molecules:

When two different elements covalently bond, the difference in

electronegativity causes an unequal sharing of bonding electrons,

so a polar bond

Electronegativity - an atom’s ability to attract and hold electron

from a bond

A polar bond can create a polar molecule. However, several polar

bonds may cancel out in a symmetrical molecule to make the

overall non-polar

- Melting and boiling points rose due to increased attractions

- Small molecules are weakly soluble

Lewis Diagrams

Lewis or Electron Dot Diagrams show how a molecule is bonded and can be used to a

certain shape

- Only show valence electrons

To Draw:

1. Put elements with the least number of electrons in the centre

2. Draw valence electrons on each four points of compass, single then

pairs

3. Form bonding pairs from leftover single electrons

4. Resolve diagram for neatness and even spacing between bonds and lone

electron pairs, and check for octet rule:

- elements need 8 electrons except Hydrogen

Types of Reactions

There are four types of chemical reactions:

- Precipitation

- Decomposition

- Oxidation - Reduction (Redox)

- Neutralization (Acid-Base)

Eg. Acid + - Hydroxide

- Oxide

- Carbonate

Precipitation (ppt)

A precipitation is a solid that is formed from a reaction of two aqueous solutions. Things

form a precipitation if a product from a reaction is insoluble.

Eg. Copper Sulfate + Sodium Hydroxide � Copper Hydroxide + Sodium Sulfate

Solubility Rules:

1. All group 1 (Na+, K+, Li+), Nitrate (NO3-) and Ammonium (NH4

+) compounds

are soluble

2. All Chlorides are soluble, except for Silver Chloride (AgCl) and Lead Chloride

(PbCl2+)

3. All Sulfates are soluble, except for Barium Sulfate (BaSO4), Lead Sulfate

(PbSO4) and Calcium Sulfate (CaSO4)

4. All Carbonates are insoluble except for ones in Rule 1

5. All Hydroxides are insoluble except for ones in Rule 1

Precipitation Equation

The precipitate that forms from mixing two solution is the reaction that occurring. The

other ions are spectator ions - they are not involved therefore an ionic equation is written

that omits spectators.

Eg. NiCl2(aq) + 2KOH(aq) � 2KCl(aq) + Ni(OH)2(s)

Ni2+(aq) + 2OH(aq) � Ni(OH)2(s)

Complexes

Complexes are soluble ions made from addition of excess of a reagent (chemicals). They

are used of for confirmation testing of unknown ions

Complex ions to remember:

1. [Cu(NH3)4]2+ � Copper tetra amine - Royal Blue

2. [Zn(OH)4]2+ � Zincate - Colourless

3. [Zn(NH3)4]2+ � Zinc tetra amine - Colourless

4. [Al(OH)4] � Aluminate - Colourless

5. [FeSCN]2+ � Iron (III) Thioscyanate - Blood Red

Unknown Identification

Solutions where the ions are unknown can be identified using precipitation reactions.

Each cations and anion must be identified separately. To bring about known, precipitation

reactions reagents are used.

Common Reactions:

- Sodium Hydroxide - Forms characteristic coloured precipitations

- Some complexes with cations

- Ammonia Solution - For making amino complexes for confirmation

- Barium Nitrate/Chloride - Identify Sulfates

- Silver Nitrate - Identify Chlorides

- Dilute Acid (HNO3) - Identify Carbonates

Decomposition (Thermal)

Decomposition is when one compound breaks into several compounds

A�B + C (D, etc)

Thermal decomposition is when this process occurs during heating. The products of

decomposition depend upon what it is

Types:

Metal Carbonate:

Forms CO2 + Metal Oxide Products

Eg. CaCO3 � CO2 + CaO

MgCO3 � CO2 + MgO

Na2CO3 � CO2 + Na2O

The limewater test confirms CO2

Metal Bicarbonate/Hydrogen Carbonate:

Forms Metal Carbonate + H2O + CO2

Eg. 2NaHCO3 � Na2CO3 + H2O + CO2

Ca(HCO3)2 � CaCO3 + H2O + CO2

Metal Hydroxide:

These decompose to form the Metal Oxide + H2O

Eg. Mg(OH)2 � MgO + H2O

2NaOH � Na2O + H2O

Metal Sulfates:

Decompose to form Metal Oxide and Sulfur Dioxide (which

usually gets further oxides to form Sulfur Trioxide)

Eg. MgSO4 � MgO + SO2

Oxidation - Reduction (Redox)

These are chemical reactions that involve the transfer of elections. Both oxidation and

reduction must occur at the same time.

L - Loss

E - Electrons

O - Oxidation

The Lion Goes

G - Gain

E - Electrons

R - Reduction

Oxidation may also be seen by:

- The gain of Oxygen element

- The loss of Hydrogen element

Eg. C + O2 � CO2 the C is oxidized

Reduction may also be seen by:

- The loss of Oxygen element

- The gain of Hydrogen element

Eg. 2PbO + C � 2Pb + CO2

Anytime there is a metal as its element in the reactants

or products of a reaction then redox has occurred

Eg. 2Mg + O2 � 2MgO Mg is oxidized as it:

- gained Oxygen

- lost Electrons

O2 is reduced as it changed to O2- so gained electrons

Oxidations is the addition of Oxygen to an element or compound

Reduction is the addition of Hydrogen to an element or compound

Oxidation is the removal of electron(s) from a particle in a reaction

Reduction is the addition of electron(s) to a particle in a reaction

The Mole

In chemistry very large amounts of particles are reacting are in a given reaction. Even

very accurate balances (scales) will be weighing trillions of atoms. The mole is a name

given to a large quantity of 6.023x1023 =1mol (Avogadro’s Number)

1mol of Fe = 6.023x1023 Atoms

0.5mol of Fe = 3.011x1023 Atoms

1mol CO2 = 6.023x1023 Molecules of CO2 ≡ 3mols of atoms (1C + 2O)

To relate to mass the relative Atomic masses (Ar) or Molar masses (M) on the periodic

table become 1mole of an element

Eg. 1mol of K will weigh 39.0g

31.6g of Cu will be 5mol

For compounds the molar mass of molecular mass must be found first by summing all the

masses for each element

M(CO2) = 12 + 2(16) = 44

M(Fe) =55.5

Mole Calculations

The relationship between molar mass (M), mass (m) and number of moles or amount of a

substance (n) are:

M � Molar mass (gmol-1) � M=m/n

m � Mass in grams (g) � m=Mxn

n � Number of moles in mol � n=m/M

Eg. 18.2g of Magnesium

n(Mg)=m/M=18.2/24.3 =0.749mol

1.6g of SO2

n(SO2)=m/M=4.2/136.4 =0.0308

1.63mol of Ammonium Sulfate

m((NH4)2S)=mxM=1.63x68 =111g

0.00489mol of Magnesium Sulfate Heptahedrate

m(MgSO4 • 7H2O)=mxM=0.00489x246.3 =1.20g

Mass to Mass

To calculate find an unknown mass for a species in a reaction from the mass of a different

species, mole is used

Eg. H2SO4 + Mg � MgSO4 + H2

If 17.2g of Mg is used what mass of MgSO4 is formed

3 steps:

1. Find the number of moles of the known species n(known)

n(Mg)=m/M=17.2/24.3=0.708mol

2. Use the equation’s mole ratio to fin the number moles for unknown n(unknown)

n(MgSO4)=n(Mg)=0.708mol

3. Find the mass of your unknown m(unknown)

m(MgSO4)=nxM=0.708x120.4=85.2g(3sf)

What mass of Sulfuric Acid is required to make 27.4g of Zinc Sulfate from Zinc Metal

H2SO4 + Zn � ZnSO4 + H2

1. n(ZnSO4)=m/M=27.4/161.4=0.170mol

2. n(H2SO4)=n(ZnSO4)=0.170mol

3. m(H2SO4)=nxM=0.170x98.1=16.7g

11.3g of coal (C) is combusted, what mass of Carbon Dioxide is produced

C + O2 � CO2

1. n(C)=m/M=11.3/12=0.942mol

2. n(CO2)=n(C)=0.942mol

3. m(CO2)=nxM=0.942x44=41.4g

24.6g of Sodium Bicarbonate is reacted with what mass of Hydrochloric Acid

2HCl + 2NaHCO3 � 2NaCl + H2O + 2CO2

1. n(NaCHO3)=m/M=246/84=0.293mol

2. n(HCl)=n(NaCHO3)=0.293mol

3. m(HCl)=nxM=0.2293x36.5=10.7g (3sf)

Non 1:1 Ratios

When the mole ratio in question is not 1:1 we multiply the n(known) by unknown/known

mole ratio in Step 2

Eg. What mass of Magnesium will react with 14.3g of Hydrochloric Acid

Mg + 2HCl � MgCl2 + H2

1. n(known)

n(HCl)=m/M=14.3/36.5=0.392mol

2. n(unknown) from the equation ratio

n(Mg)=1/2xn(HCl)=1/2x0.392=0.196mol

3. m(unknown)=nxM=0.196x24.3=4.76g (3sf)

If 8.63g of Sodium Sulfate is produced what mass of Sodium is used

H2SO4 +2Na � Na2SO4 + H2

1. n(Na2SO4)=m/M=8.63/142=0.0608mol

2. n(Na)=2/1xn(H2SO4)=2x0.0608=0.122mol

3. m(unknown)=nxM=1.22x23=2.80g (3sf)

What mass of Carbon Dioxide is produced from 1.63g of Nitric Acid with

Magnesium Carbonate

MgCO3 + 2HNO3 � Mg(NO3)2 + CO2 +H2O

1. n(HNO3)=m/M=1.63/63= 0.026mol

2. n(CO3)=1/2xn(HNO3)=1/2x0.026=0.0129mol

3. m(unknown)=nxM=0.0129x44=0.569g (3sf)

What mass of Iron (III) Oxide is required to react with 186g of Hydrochloric Acid

Fe2O3 + 6HCl � 2FeCl3 + 3H2O

x/159.6 = 186/219(=36.5x6)

x=29685.6/219

x=135.55

x=136 (3sf)

1. n(HCl)=m/M=186/36.5=5.10mol

2. n(Fe2O3)=1/6xn(HCl)=1/6x5.10=0.8493mol

3. m(unknown)=nxM=0.8492x159.8=136 (3sf)

Percent for Composition

The percent of composition for a compound tells us the contribution to the total mass

from each element

Step 1. Find the molar mass for the compound

2. Divide each elements contribution by the total and then multiply by 100

Eg. BaCl2 • 2H2O

M(BaCl2 • 2H2O)= 137.3 + 35.5(2) + 1(4) + 16

=228.3

Ba=137.3/228.3 x100= 60.14

Cl=71/228.3 x100= 31.1

H=4/228.3 x100= 1.75

O=16/228.3 x100= 7

99.998%

Empirical Formulas

An Empirical formula is the simplest ratio of elements in a compound. A molecular

formula is the actual ratio

Eg. Empirical Formula � Molecular Formula

CH3 � C2H6

MgSO4 � MgSO4

CH2O � C6H12O6

The molar mass and percentage of composition of unknown compounds can be found

from mass spectrometers. This information is used to find empirical and molecular

formulas

Step 1. Assure you have 100g of the substance, then all percentages becomes mass

2. Divide each element’s mass by its molar mass

3. Divide all answers from 2 by the smallest one

4. Judge the numbers to a whole number.

5. Find the ratio of empirical formula Mr to given Mr to get a multiplier from the

molecular formula

Eg. An unknown has Mr=180, 40% C, 6.7% H and 53% O

40g C, 6.7g H, 53g O

C = 40/12=3.33 � 3.33/3.31 ≈ 1

H = 6.7/1=6.7 � 6.7/3.31 ≈ 2

O = 53/16=3.31 � 3.31/3.31 ≈ 1

M(CH2O)=30, Ratio =180/30=6=Multiplier

So the molecular formula is C6H12O6 (=6[C1H2O1])

Water Crystallization

When a solution is evaporated to form crystals, water molecules become trapped in the

spaces within the crystal lattice. The amount of waters to the salt is in ratio in a hydrated

salt. (Eg. CuSO4 • 5 H2O has 5 waters per CuSO4)

By using gravimetric analysis (accurate weighing) we can determine what that ratio is and

deduce the formula of the hydrated salt.

Step 1. Find the amount of Water

2. Find the amount of anhydrous salt

3. Divide the answers by the smallest to get a whole number ratio

Eg. A hydrated sample of Barium Chloride is heated to determine the waters of

crystallization

Mass BaCl2 • xH2O=4.862g

Mass after heating =4.146g

1. m(H2O)= 4.862-4.146=0.716g

n(H2O)=m/M=0.716/18=0.0398mol

2. m(BaCl2)=4.16g

n(BaCl2)=m/M=4.146/208.3=0.0199mol

3. 0.0199/0.0199:0.0398/0.0199=1:2

So the hydrated formula is:

BaCl2 • 2H2O

Non-Metals

Non-metals elements are located on the right of the periodic table’s zigzag. They have

different properties to metals:

- Generally do not conduct (except graphite)

- Have much lower melting and boiling points, except diamond and quartz (SiO2)

- Are more often molecular structures rather than extended lattices

- Often have different colours and are non-shiny

- They form acidic compounds (metals then to form alkali compounds)

Oxygen

Oxygen forms a diatomic gas as an element. It has a melting point of -206°c and a boiling

point of -186°c. Oxygen’s density = 0.0032 gcm3. An allotrope (different physical form of

same element in same state) of oxygen is Ozone (O3)

Ozone is a dangerous respirator irritant in low levels. Often produce by electrical spark or

UV light. Its pale blue and sweet sickly smelling.

The ozone layer filters dangerous UVA and UVB light by absorption. Excessive CFCs

has caused a depletion of O3, thus creating the ‘Ozone Hole.’

Commercial production of Oxygen and Nitrogen gas

Air is cooled and compressed several times followed by an adiabatic expansion.

Air consists of a number of gasses including N2, H2, O2, H2O, CO2 and Nobel gasses Ar,

Ne etc.

The liquid air is allowed to warm slightly to boil off the O2 and recondensed, thus N2 and

O2 are obtained by the fraction distillation of liquid air, O2 production in the lab.

O2 is made in the lab from decomposition of various substances

1. Hydrogen Peroxide with Manganese Dioxide Catalyst

2H2O2 � 2H2O + O2

2. Heating Potassium Permanganate

2KMnO4 � 2KMnO3 + O2

Nitrogen

N2 is a clear colourless gas with a melting point of -210°c and a boiling point of -196°c.

Its density = 0.808 gcm3. It consists of a diatomic molecule with a strong triple bond. This

makes it generally uncreative as a substance and is often used as an iner carriergass or as

a liquid for cryogenics and super cooling.

Nitrogen is made commercially by the fractional distillation of liquid air. In the lab it’s

made by heating Sodium Nitrate with Ammonium Chloride.

NH4Cl + NaNO2 � NaCl + NH4 NO2

The Ammonium Nitrate then decomposes

NH4NO2 � N2 + 2H2O

Nitrogen is an essential element in plant (and animal) growth for making proteins. To

obtain nitrogen plants require soluble nitrates (NO3-) which can be uptake from the soil.

Nitrogen is ‘fixed’ to nitrates by: - Lightning

- Nitrogen fixing bacteria in legumes plants

- Addition of Ammonium or Nitrate based fertilizer

- Decaying of dead organisms releases some

Nitrates and also denitrifying bacteria release N2 to

the air

Nitrogen Oxides

Nitrogen car reacts with oxygen with the presence of an electrical discharge (spark).

Mostly in lightning strikes (N2 fixing) but also large amounts inside car engines.

Commonly known as NOx there are three main forms:

1. Nitrous Oxide (N2O).

N2O is also known and used as laughing gas, or a commonly use

anaesthetic. Can ve made in the lab by decomposing Ammonium Nitrate

NH4NO3 � N2O + 2H2O

2. Nitric Oxide (NO)

Colourless gas that immediately oxidises in air to NO2, made from

reactions of Dilute Nitric Acid with Copper Metal

3Cu + 8HNO3 � 3Cu(NO3)2 + 2NO + 4H2O

3. Nitrogen Dioxide (NO2)

Dense brown pungent gas that sinks into low lying areas. Made from

Concentrated Nitric Acid and Copper Metal

Cu + 4cHNO3 � Cu(NO3)2 + 2 NO2 + 2H2O

NO2 is an acid gas that reacts by hydrolysis in water

2NO2(g) + H2O � HNO2 + HNO3

HNO2 = Nitrous Acid

HNO3 = Nitric Acid

In areas with significant vehicle congestion, then NO2 causes acid rain and

Photochemical Smog

Acid Rain: - Damages buildings and structures

- Causes skin irritations in animals

- Destroys leaf foliage on plants

- Lowers pH of soil and waterways, often killing life

Photochemical Smog: - Brown haze in valleys or low lying areas

- NO2 will react with Oxygen molecules, catalysed by

sunlight to produce ozone

O3: - Respiratory irritant

- Damages eye tissue

- Some crop/plant damages

However - NO2 can be useful to plants as it produces nitrates (from HNO3) which

are needed to make protein for growth

Ammonia

Ammonia NH3 is a pungent colourless gas. That is denser than air. Ammonia acts as a

base and reacts with water to form Ammonium Hydroxide:

NH3 + H2O � NH4OH

Ammonia is detected by: - Damp Red Litmus � Turns Blue

- White smoke when reacted with HCl vapour

NH3 + HCl � NH4Cl (white)

Ammonia is produced in the lab by heating an Ammonium Salt with a strong Alkali

NH4Cl + Ca(OH)2 � CaCl2 + 2H2O + 2NH3

Ammonia vapour is very soluble and will generate a vacuum over water

Haber Process

Ammonia is produced commercially by the Haber Process. The raw materials are

Methane gas and Water.

Step 1 - CH4 and Water are steam reformed to make synthesis gas

CH4 + H2O � CO + 3H2

Synthesis Gas

Step 2 - Air (mix of O2 and N2) is added to the mixture. This further reacts with

the CO and some of the H2

CO + O2 � CO2 - Removed for later use

2H2 + O2 � H2O - Recycled back into Step 1

Now reaction mixture contains leftover H2, and unreacted N2

Step 3 - The ‘Haber Step’, the N2 and H2 is brought together in 1:3 ratio over an

Iron III Oxide catalyst

N2(g) + 3H2(g) (Fe2O3)� 2NH3(g)

Ammonia is mostly used for fertiliser

Eg. - Reacted with CO2 from Step 2

2NH3 + CO2 � NH2CONH2 + H2O

- Bubbled through Sulfuric Acid

2NH3 + H2SO4 � NH4SO4 - Ammonium Sulfate

- Bubbled through Nitric Acid

NH3 +HNO3 � NH4NO3 - Ammonium Nitrate

Ammonia through Ammonium Nitrate is also used to make explosives such as artillery

shells and bombs

NH3 also used: - For a Refrigerant

- As a household cleaner

- Manufacture of dyes and nylon

Sulfur

Sulfur is a yellow powder at room temperature. It is insoluble in water but will dissolve in

non-polar solvents such as Carbon Disulfide. It has a mp=112°c and bp=445°c. Sulfur

exists as an S8 molecule in a ‘puckered’ ring shape

It has two allotropes: - Rhombic Sulfur up to 96°c

- Monoclinic Sulfur from 96-112°c

And a third from of ‘Plastic Sulfur’ when molten, where the rings open and form

long entangled chains. This is flexible like plastic

Sulfur Extraction

Sulfur occurs naturally around volcanic areas and as impurities in fossil fuels

It can be extracted: - By open mining

- Frasch process which uses super heated steam to drive liquid S

from underground deposits

- Obtained by reducing H2S impurities in natural gas.

Sulfur Reactions:

- Sulfur reactions with Metals to form Metal Sulfides

Eg. 2Ag + S � Ag2S - Silver Tarnish

2Fe + 3S � Fe2S3

Zn + S � ZnS - Yellow Powder

Metals Sulfides easily react with Acids to make dangerous H2S, rotten egg smells toxic

gas

- Sulfur burns with a blue flame to form Sulfur Dioxide molecules in Oxygen

S + O2 � SO2

This is reactive as it’s unstable, but partly stabilised by resonance

Sulfur Dioxide

Sulfur Dioxide SO2 can be prepared in the lap from reactions pf a Sulfur Salt

Na2SO3 + 2HCl � 2NaCl + H2O + SO2

As it is more dense than air it is collected by upwards displacement of air

Large amounts of SO2 are produced in industries from impurities in fossil fuels that are

burnt for energy

SO2 is soluble in water to produce Sulfurous Acid

SO2 +H2O � H2SO4 - Sulfurous Acid

This causes Sulfur based acid rain in areas with large industry

This causes - Damage to buildings/concrete

- Skin irritations to animals/humans

- Foliage damage to plants

- Acidifying soil and waterways

There is no gain from Sulfur based acid rain unlike Nitrogen based acid rain

Uses of SO2 - As a sterilant for bottles and food containers, often used in the

form of Metabisulfite

- To prevent food spoiling in transport, SO2 kills bacteria and also

is a Reductant, so prevents oxidation of fruit etc

- Bleaching in some paper mills

- For the manufacturing of Sulfuric Acid

Sulfuric Acid

H2SO4 is a strong diprotic inorganic acid. It is at maximum strength at approximately

70% when there is full Disassociation (splits into ions)

H2SO4 � H+ + HSO-4 � 2H+ + SO-

4

(�) = H2O

So overall: H2SO4(l) + 2H2O(l) � 2H3O+

(aq) + SO2-4(aq)

Sulfuric Acid at 99% has not dissociated so is non-acidic and can be stored in mild steel

vessels. It is a strong dehydrating agent and will remove water molecules from most

substances

Industrial Manufacture - The Contact Process

Sulfuric Acid is manufactured by the millions of tonnes, mostly for fertilizer

Step 1 - Sulfur is vaporised and combusts instantly on ’contact’ with air

S(g) + O2(g) � SO2(g)

Step 2 - The SO2 is further oxidised over a Vanadium Pent oxide (V2O5) catalyst

bed at 450°c

2SO2 + O2 � 2SO3

Step 3 - The SO3 is dissolved into concentrated Sulfuric Acid to make Oleum

SO3 + H2SO4 � H2S2O7 - Oleum

It cannot be dissolved directly in water as it is exothermic forming an

useable fog

Step 4 - The Oleum is diluted to make 98% H2SO4

H2S2O7 � 2H2SO4

Most H2SO4 is made into Superphosphate by reacting with insoluble rock

Phosphate/Calcium Phosphate. This is done to make a soluble source of

Phosphate

H2O + Ca3(PO4)2 + H2SO4 � CaSO4 • 2H2O + Ca(H2PO4)2

CaSO4 • 2H2O + Ca(H2PO4)2 - Superphosphate

Chlorine

Chlorine is a yellow-green acidic gas that is denser than air. It exists as a diatomic

molecule. Its mp= -101°c, bp= -35°c and density= 1.56 gcm-3. Chlorine is quite soluble in

water to for two acids by hydrolysis

Cl2 + H2O � HClO + HCl - Hypochlorous Acid + Hydrochloric

Cl2 is a very strong oxidant and will react with most substance like, metals including inert

ones, non-metals and common reductants eg SO2-4

Lab Production

It is made by oxidising concentrated HCl with KMnO4 (Potassium Permanganate) or

MnO2 (Manganese Dioxide) with heat

2KMnO4 + 16HCl � 2KCl + 2MnCl2 + 8H2O + 5O2

Cl2 supports combustion eg. 2Fe + 3Cl2 �2FeCl3

Cl2 is produced commercially by electrolysis using The Membrane Cell

Uses of Chlorine (Cl2)

Chlorine is mostly used in the production of Hydrogen Chloride (HCl) by Synthesis

Reaction

H2(g) + Cl2(g) � 2HCl(g)

In water this is Hydrochloric Acid

HCl(g) + H2O � H3O+ + Cl-

Cl2 is also heavily used in the pulp and paper industry to bleach colour from paper pulp

Other uses include: - Purifying water by sterilising and killing bacteria

- Manufacture of ‘Chlorine’ bleaches for commercial

cleaning/bleaching of cloth

- To make plastics such as PVC

- Solvents, pesticides and aerosol propellants (CFC)

Hypochlorite ion (OCl-)

OCl- has similar properties to chlorine but is in a much more usable form. It is present is

chlorine water from the Hypochlorous Acid

HOCl � H+ + OCl-

Also made by bubbling Cl2 through a base

With Sodium Hydroxide � Sodium Hypochlorite

2NaOH + Cl2 � NaCl + NaOCl + H2O

- Household cleaners and bleaching powder

With Calcium Hydroxide � Calcium Hypochlorite

2Ca(OH)2 + 2Cl2 � CaCl2 + Ca(OCl)2 + 2H2O

- Used for pool chlorine and drinking water in low concentrations

Hydrogen Chloride

HCl(g) is a polar molecule, It has a dipole across the bond and therefore over the molecule.

The permanent dipole interactions raise its mp and bp alone that the equidistant non-polar

molecules.

In water it forms Hydrochloric Acid

HCl(g) + H2O � H3O+ + Cl-

It’s very in water (like ammonia) causing a vacuum over water surface and detect for, by

a white smoke forming with one ammonia.

To make in industry � Synthesis

To make in the lab react Concentrated Sulfuric Acid with a Chloride Salt

H2SO4 + NaCl � HCl + NaHSO4

Beam.A