Wednesday, 8 April 2009

iGCSE Chemistry - Chemistry of the Elements

The Periodic Table

 

  • Periodic table arranges elements by their atomic number.
  • Groups are vertical columns.
  • Periods are horizontal rows.
  • Group 1 elements are all alkali metals.
  • Group 2 elements are all alkaline earth metals.
  • Group 7 elements are all halogens.
  • Group 8 elements are all noble gases.
  • Metals are on the left and middle of the periodic table.
  • Non-metals are on the right of the periodic table.

 

Properties of Non-metals

Properties of Metals

Low melting points

See above, in “Metallic Crystals” section

Poor conductors of electricity

Poor conductors of heat

Brittle

Dull

 

  • Metals become more reactive as you go down the group.
  • Non-metals become less reactive as you go down the group.

 

Group 1 Elements

 

  • All react with water to produce alkaline solutions.
  • One electron in their outer shell.
  • Reactivity increases as you move down the group because the outer electrons are further from the nucleus and so are easier to be lost.

 

Properties of Group 1 elements

Soft to cut

Shiny when cut, but soon oxidise

Very low melting points compared to other metals

Very low densities compared to other metals

React very easily

 

 

 

Element

Colour of flame when burned in air

Lithium

Red

Sodium

Orange/yellow

Potassium

Lilac

 

·        When reacted with water, the metals react vigorously, float on the surface, moving around rapidly. The metal forms a sphere and gives off hydrogen.

·        When reacted in chlorine, the metals produce a white solid.

·        Group 1 nitrates decompose upon heating.

·        Lithium Nitrate decomposes differently to the other nitrates as it produces 3 products rather than 2 (i.e. it decomposes further)

·        Lithium carbonate is the only group 1 carbonate that decomposes under a Bunsen.

·        Compounds of group 1 metals are usually colourless crystals and formed by ionic bonding, and thus are usually soluble in water

 

Group 2 Elements

 

  • Two electrons in their outer shell.
  • Less reactive than group 1 as 2 electrons rather than 1 has to lost.
  • Reactivity increases as you move down the group because the outer electrons are further from the nucleus and so are easier to be lost.

 

Properties of Group 2 elements

Harder than group 1

Shiny when cut, but soon oxidise (slightly slower oxidation than group 1)

Higher melting points than group 1

Higher densities than group 1

React very easily, but less than group 1

 

Element

Colour of flame when burned in air

Calcium

Brick-red

Strontium

Crimson

Barium

Apple-green

 

  • Compounds of group 2 elements are usually white and crystalline.
  • The oxides react with water to form hydroxides.
  • The carbonates decompose when heated.
  • The nitrates decompose when heated.

 

Group 7 Elements

 

  • Group 7 elements are halogens.
  • Halogens react with most metals to make salts.
  • Seven electrons in outer shell.
  • Very reactive.
  • Reactivity of elements decreases as you move down the group.
  • All halogens exist as diatomic molecules.
  • They undergo displacement reactions.

 

Element

Appearance

Fluorine

Pale yellow gas

Chlorine

Yellow-green gas

Bromine

Brown liquid

Iodine

Black shiny solid

 

  • To obtain chlorine from hydrochloric acid, it is heated with manganese dioxide.
  • The test for chlorine is that it causes damp indicator paper turn white.

 

Oxygen and Oxides

 

  • The air contains these main gases:

 

Gas

Amount in air

Nitrogen

78.1%

Oxygen

21%

Argon

0.9%

Carbon Dioxide

Trace

 

  • Elements form oxides when heated in oxygen.
  • Oxides can be classified as:

              Basic oxides: when dissolved in water form alkalis

              Acidic oxides: when dissolved in water form acids

              Neutral oxides: when dissolved in water form neutral                           solutions

  • Bases and alkalis react with acids to form salts in neutralisation reactions.
  • Reactions in oxygen:

Element

Reaction with oxygen

Magnesium

White powder formed

Iron

Black oxide formed

Copper

Black oxide formed

Carbon

Colourless gas formed

Sulphur

Colourless gas formed

Methane

CO2 and water formed

 

  • A chemical reaction where oxygen is added is called oxidation
  • A chemical reaction where oxygen is removed is called reduction
  • To find percentage of oxygen in the air:
    1. Set up two syringes connected by a tube containing copper turnings. Ensure the syringes contain 100cm3 in total of air.
    2. Heat from below with a Bunsen.
    3. Push air back and forth between syringes.
    4. The copper reacts with the oxygen to form copper dioxide
    5. When all of the oxygen has been used up, 70cm3 will be left in the syringes.
    6. Thus, oxygen makes up 21% of the air.
  • Sulphur dioxide is formed when sulphur reacts with oxygen.
  • Sulphur dioxide reacts with water to form sulphurous acid.
  • Carbon dioxide can be formed by reacting hydrochloric acid with calcium carbonate (marble chips).
  • The test for carbon dioxide is to bubble it through limewater. If the limewater goes cloudy, carbon dioxide is present.
  • Carbon dioxide freezes at -78°C, forming dry ice.
  • Carbon dioxide reacts with water to form carbonic acid.
  • Carbon dioxide reacts with alkalis to form carbonates.
  • Nitrogen can react with oxygen to form nitrogen monoxide and nitrogen dioxide, which in turn dissolve in rain water to cause acid rain.
  • The process by which iron is corroded is called rusting, and it is sped up if electrolytes (such as sodium chloride [salt]) are present.
  • For rusting to occur, both air and water must be present.
  • Galvanisation is the process by which iron is covered in a more reactive metal, such as zinc, to stop it rusting. The zinc is corroded instead.
  • If metal oxides are heated with carbon they will be reduced if the metal is less reactive than carbon.

 

Sulphur

 

·        Sulphur is released by volcanoes.

·        When sulphur dioxide is dissolved in water it forms sulphurous acid.

·        The salts of sulphurous acids are called sulphites

·        When sulphites are reacted with dilute acid, sulphur dioxide is released.

 

Nitrogen and Ammonia

 

  • Nitrogen is obtained through the fractional distillation of liquid air.
  • Ammonia is a compound of nitrogen.
  • Ammonia is used to produce fertilisers, and is industrially produced during the Haber process.
  • Ammonia is a colourless gas that is less dense than air, as well as being very soluble and being the only common alkaline gas.

 

 

Hydrogen

 

  • Metal + Acid à hydrogen + salt
  • The test for hydrogen is that it burns and causes a pop sound.
  • The test for water is that it causes anhydrous copper sulphate to turn from white to blue.
  • To test whether water is pure, its boiling point needs to be measured accurately. If it is 100°C exactly, the water is pure.
  • Hydrogen reacts with chlorine when heated and hydrogen chloride is formed. This may cause an explosion.

 

 

The Transition Metals

 

  • Transition metals are found in the middle of the periodic table.
  • All transition metals have more than one electron in their outer shell, making them much less reactive and thus more common.
  • Transition metals are often used as catalysts.

 

 

Property of Transition Metals

High melting point

High boiling point

Coloured compounds

Slow/no reaction with water or acid

 

  • Iron reacts very slowly with water but reacts quickly with steam. Iron oxide and hydrogen are produced.
  • When iron is reacted with hydrogen chloride¸ Iron (II) Chloride and hydrogen are produced.
  • When iron is reacted with chlorine, Iron (III) Chloride is formed.
  • Iron can form two different hydroxides, Iron (II) hydroxide and Iron (III) hydroxide.
  • Iron (II) hydroxide is white when pure, but often appears to be a pale green colour.
  • Iron (III) hydroxide is orangey-brown.
  • Copper compounds are usually blue or green.
  • Copper (II) oxide is formed when copper is heated.

 

Compound

Colour

Copper (II) Oxide

Black

Copper (II) Chloride

Green

Copper (II) Nitrate

Blue

Copper Sulphate

Intense Blue

 

  • Some copper compounds decompose on heating.
  • Important copper compounds are copper (II) compounds, but copper (I) compounds exist, such as copper (I) oxide, which is red-brown.
  • When ammonia is added to copper (II) ions in solution, copper (II) hydroxide forms.
  • If excess ammonia is added, the copper (II) hydroxide dissolves, and a complex ion is formed:

 

[Cu(H2O)2(NH3)4]2+

 

Transition metal properties

Variable valency

Formation of coloured compounds

Formation of complex ions

 

Reactivity Series

 

1.       Potassium

2.     Sodium

3.     Lithium

4.     Calcium

5.     Magnesium

6.     Aluminium

7.     Carbon

8.     Zinc

9.     Iron

10.  Tin

11.   Lead

12.  Hydrogen

13.  Copper

14.  Silver

15.  Gold

16.  Platinum

 

·        Displacement reactions can be used to determine the reactivity series.

·        Iron can be stopped from rusting by being coated in a layer of zinc (galvanisation).

·        This is because zinc is more reactive than iron.

·        The zinc is called a sacrificial anode.

 

 

Preparing and Analysing

 

 

  • Flame tests for cations:

 

Ion

Colour in flame

Lithium

Bright red

Sodium

Golden yellow/orange

Potassium

Lilac

Calcium

Brick red

 

  • The test for ammonium ions is:
    1. Add dilute sodium hydroxide solution
    2. Test with damp red litmus paper
    3. If litmus paper goes blue ammonium ions are present.

·        Test for aqueous solutions of cations:

 

Ion

Test

Result

Copper (II)

Add sodium hydroxide

Light blue precipitate

Iron (II)

Add sodium hydroxide

Green precipitate

Iron (III)

Add sodium hydroxide

Red/brown precipitate

 

  • Test for anions in solution:

 

Ion

Test

Result

Chloride

1. Add dilute Nitric acid

 

2. Add silver nitrate solution

White precipitate

Bromide

Cream precipitate

Iodide

Yellow precipitate

 

  • To test for the carbonate anion, dilute hydrochloric acid is added and the gas given off is passed through limewater. If the limewater goes cloudy (i.e. carbon dioxide present) then carbonate anion is present.
  • Tests for gases:

 

Gas

Test

Result

Hydrogen

Put in a flame

Pop heard

Oxygen

Put in glowing splint

Splint relights

Carbon dioxide

Pass through limewater

Limewater goes cloudy

Chlorine

Put in damp blue litmus paper

Paper goes red

Ammonia

Put in damp red litmus paper

Paper goes blue

Sulphur dioxide

Put in damp potassium dichromate paper

Paper goes green

 

  • Solubility of salts:

          All common sodium, potassium + ammonium salts are soluble

          All nitrates are soluble

          Common chlorides are soluble, except silver chloride

          Common sulphates are soluble, except barium and calcium

          Common carbonates and hydroxides are insoluble, except                      sodium, potassium and ammonium

  • Insoluble salts can be formed as precipitates of a reaction of two soluble salts.

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