Non- Metal Chemistry

The 6 Non-metals you will be learning about are: oxygen, sulfur, nitrogen,

carbon, chlorine and bromine.

Metals vs Non-metals

• Metals and Non-metals• Metals are found on the left hand side of the

periodic table and the non-metals are found on the right. Elements physical properties are used to characterise them as either a non-metal or metal.

• The differences in the chemical properties of metals and non-metals are shown in the table below.

Metals Non-metalsAll except Mercury are solids at room

temperatureThey can be found as solids, liquids

and gases at room temperature.

Their appearance is shiny

Do not conduct electricity

These are strong and bendy. Most can be hammered into different shapes( malleable)

Can be drawn out into wires (they are ductile)

Can not be drawn out into wires

Most have low melting and boiling points except carbon.

They are good conductors of heat

Metals Non-metals

All except Mercury are solids at room temperature

They can be found as solids, liquids and gases at room temperature.

Their appearance is shiny Most are dull when solid

Are good conductors of electricity Do not conduct electricity

These are strong and bendy. Most can be hammered into different shapes( malleable)

These can not be shaped

Can be drawn out into wires (they are ductile)

Can not be drawn out into wires

Most have high melting and boiling points.

Most have low melting and boiling points except carbon.

They are good conductors of heat Do not conduct heat very well.

Nitrogen

Nitrogen

Gas at roomtemp

Found in amino Acids, proteins

And nitrates

No lab test For nitrogen

Has the sameDensity as air

Insoluble in water

Colourless,Tasteless and

Odourless

Makes up 78% ofThe atmosphere

Used to make Ammonia via the

Haber process

Used to create anInert atmosphere &Used in refrigerant

Has many Associated

oxides

Commercial Production via

Fractional Distillation.

The Nitrogen Cycle

Nitrogen is essential for living things because it is used to make proteins. Proteins are the building blocks that make up all living things.

Summary

Uses of Nitrogen

• Artificial fertilisers contain nitrogen compounds to help plants grow.

• Liquid nitrogen. Nitrogen gas turns into a liquid at -195.8°C which makes it a good coolant. It is used to freeze warts and other biological material such as embryos for storage.

• Several nitrogen compounds are highly explosive. E.g nitro-glycerine and TNT.

• Nitrous oxide (N2O) is used as an anaesthetic.

Normally we do not absorb nitrogen from the atmosphere but when the air we breathe is pressurised, nitrogen gas can dissolve in your blood. N2 bubbles form in the blood. This is known as ‘the bends’.

Chemical Properties of Nitrogen

Nitrogen gas comprises of two nitrogen atoms __________ bonded together, forming a triple bond. The bond is very _________ and requires a large amount of _________ to break. Only at very high ______________ or if an electrical spark is passed through nitrogen gas will it react with ____________. This does occur in the internal combustion engine.

Nitrogen Compounds

Nitrogen reacts withOxygen to produce

Nitrogen dioxideNO2

Nitrogen monoxideNO

DinitrogenMonoxide

N2O

Nitrogen cycle

NitratesNO3

-Nitrites

NO2-

Nitric acidHNO3

Ammonia NH3 and Ammonium compounds

NH4+

Amino acids which are theBuilding blocks for all

Proteins.NH2-CH2-COOH

Nitrogen Dioxide NO2

Nitrogen dioxide is a toxic gas with a choking smell. It has an irritating effect on humans, affecting the nose, throat and eyes.

When it reacts with water it produces an acidic solution causing acid rain. It also is involved in the complex series of reaction, with other pollutants in the atmosphere producing photochemical smog.

Nitrous oxide N2O

It is also called dinitrogen oxide or laughing gas. It is a colourless slightly sweet smelling gas that is soluble in water producing a neutral solution. It has a slightly anaesthetic properties.This is a covalent compound that is produced when ammonium nitrate is heated.

Nitric oxide NO

This is also known as nitrogen monoxide or nitrogen oxide. It is a colourless gas that is soluble in water. It is a covalent compound made when copper reacts with 50% nitric acid. It then reacts with oxygen to form nitrogen dioxide.

2NO + O2 → 2NO2

Nitrogen compounds are formed when cars burn fuel. It is these compounds that are responsible for pollution in the air. To reduce the pollution that is caused by these nitrogen oxides, cars which run on unleaded petrol can be fitted with a catalytic converter. This uses a catalyst and high temperature in the exhaust pipe to remove oxides by reacting it with carbon monoxide.

Photochemical smog

Photochemical smog is a form of local pollution caused by the internal combustion engine. Some of the chemicals produced and emitted by the engine (e.g unburnt fuel, oxides of nitrogen and sulfur dioxide) react in the presence of sunlight with other chemicals in the atmosphere. This results in a blue-brown haze which is photochemical smog.

Ammonia

Structure

NH4

Properties

Colourless gas

Less dense thanAir.

Sharpe, pungent Smell.

Very soluble in water

When it is in solutionIt has basic properties

Used in cleaningproducts

Made by the Haber process

Preparation:Calcium hydroxide +Ammonium chloride

Reacts with Conc. HCl

Used to makefertiliser

Lab Preparation

Mixture of calcium hydroxide and ammonium chloride.

Mixture of calcium hydroxide + ammonium chloride

Fountain ExperimentConcentrated ammonia solutionIs heated to fill the flask with ammoniaGas. This also increase the pressureIn the flask.

The flask is removed from theHeat and the pressure decreasesThis causes the water to rushInto the flask dissolving the Ammonia gas in the processDecreasing the pressure more.

Haber Process

This process produces ammonia gas that is required for the production of fertiliser. The process involves the reaction of nitrogen and hydrogen to produce ammonia.

Nitrogen needed for the process is obtained from the atmosphere via fractional distillation.

• Hydrogen is obtained from the reaction between methane (natural gas) and water.

CH4 (g) + H2O(g) → 3 H2 (g) + CO(g)

This is carried out at 750°C and 30 atmp with a nickel catalyst.

The carbon monoxide is reduced with more unreacted steam to form even more hydrogen.

CO(g) + H2O(g) → H2 (g) + CO2(g)

Nitrogen and hydrogen are then pressurised to approximately 200 atm and passed over a catalyst of iron at between 350°C and 500°C.

The equation for the reaction is N2(g) + 3H2(g) → 2NH3(g)

About 15% ammonia is produced it is collected and removed as a liquid.

Uses of Ammonia• Used to make pharmaceuticals• Commercial cleaning products• Fertiliser• Precursor to nitrogenous compounds• Refrigeration• As a fuel• Dye in woodworking• Antimicrobial agent in food• Stimulant• Prewash for wool.

Properties of Oxygen• Has the formula O2• Colourless, odourless, tasteless, neutral gas• Slightly soluble in water.• mp -218°C and bp -183°C• Makes up 21% of the atmosphere and is removed via

fractional distillation.• Very reactive and forms oxides with most other

elements. • Consists as two allotropes O2 and O3 ozone.• Used by plants and animals for respiration.

Allotropes of Oxygen

There are two allotropes of oxygen:

O2 which is a molecule of oxygen and O3 which is ozone.

Ozone is found in the atmosphere. It is formed when solar radiation breaks the bonds in the oxygen molecules producing two oxygen atoms. These oxygen atoms react with another oxygen molecules to form ozone.

These ozone molecules absorb more radiation and break down into an oxygen molecule and an oxygen atom.

Two oxygen atoms now join to form oxygen

Ultraviolet radiation

Ozone Molecule Oxygen molecule Oxygen atom

Oxygen atom Oxygen molecule

The Ozone Hole

Ozone is constantly being made and broken down but it is happening in balance.

Uses of Oxygen (O2)

• Oxidation reactions- Burning, rusting and cell respiration.

• In medicine, anaerobic bacteria that causes gangrene can be killed with oxygen.

• Pure oxygen helps sick people who struggle to breathe.• Industrial furnaces and gas welding use oxygen to

produce intense heat.• High altitude mountaineers, astronauts and pilots need

to use oxygen.

Uses of Ozone (O3)

• Ozone is known as natures most powerful disinfectant. It is used in air and water purification, deodorization and food sanitation.

• Prevents electromagnetic radiation from reaching the earths surface.

Sulfur

Structure of Sulfur

Sulfur is naturally found as a yellow solid.

It is a molecular solid made up of ‘crown shaped molecule’ each containing 8 atoms.

It is often called a pucked ring.

Facts• Sulfur is often found in volcanic regions. Such as Rotorua and

White island in NZ. USA, Poland, Mexico, Sicily and Japan. • It is insoluble in water and does not conduct electricity.• It has a low melting point (119°C) and boiling point (444°C).• It is extracted by the Frasch process.• It burns in oxygen to produce sulfur dioxide.• It is used to make drugs, pesticides, matches and paper and is

added to rubber to make it string.• Its also used to make sulfuric acid via the Contact process. • It can also be found in high amounts in crude oil and natural

gas.

Heating Sulfur

Heating Sulfur

Stages in the Heating process

The sulfur melts and forms a thin‘runny’ yellow liquid

If it is poured quickly into waterPlastic sulfur is formed.

On further heating the liquid Becomes runny again and then

Boils (444°C)

The sulfur turns dark red and Becomes thick.

Explanation of the stages

The sulfur molecules are gaining Energy so they move freely

The short chains tangle becauseThey are cooled so quickly.

More energy the chains are brokenThey move around freely.

When more energy is added the Molecules break apart and

Become tangled. This causes theLiquid to become thicker

The Contact process

S + O2 → SO2SO2 + O2 → SO3

H2S2O7

SO3 + H2SO4 → H2S2O7

H2O + H2S2O7 → 2 H2SO4

Properties of Sulfuric acid • Colourless liquid• Boiling point 338°C.• Density 1.84 g ml-1 (this is twice as dense as water).

• When concentrated sulfuric acid is added to water heat is released.

• It reacts with water via the following reactions: H2SO4 + H2O → H3O+ + HSO4

-

Then HSO4

- + H2O → H3O+ + SO42-

• It has all the properties of an acid.• Used as a dehydrating agent• Is used as a catalyst in many reactions.

Uses of H2SO4

• Manufacture of fertilisers such as superphosphate, ammonium sulfate and potassium sulfate.

• Making soapless detergents and soap.• Cleaning metals prior to electroplating.• Making paints, dyes, explosives, plastics and

aluminium sulfate.• Making synthetic fibres such as rayon.• Is the electrolyte in lead acid batteries.

• Used in the extraction of metals such as copper, vanadium, manganese and uranium from their ores.

• Drying agent for many gases.• Petroleum refining- production of high octane

petrol.• Manufacture of certain medicines.• Useful bleaching agent in the lab.