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  • Created by: Elena
  • Created on: 02-04-13 20:33

History Of The Atom

Early 1800s - John Dalton - all atoms of the same element are the same

Late 1800s - J. J. Thomson - discovered the electron

1911 - Ernest Rutherford - discovered that the atom had a dense centre called the nucleus

1912 - Niels Bohr - predicted that electrons occupy orbitals

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Structure of an Atom

An atom has a central nucleus surrounded by shells of negatively charged electrons.

The nucleus is made up of protons and neutrons. The nucleus is positively charged but the atom has no overall charge.

An atom has no overall charge because it has the same number of protons and neutrons, so the charges cancel each other out.

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Elements and Compounds

An element is made of one type of atom. Elements can't be chemically broken down.

A compound is a substance made of two or more elements that are chemically combined. You can identify the elements in a compound from its formula, using the periodic table.

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Mass and Atomic Number

The mass number is the total number of protons and neutrons in an atom.

The atomic (proton) number is the number of protons in an atom.

The elements in the periodic table are arranged in increasing atomic number:

The group number is the same as the number of electrons in the outer shell of an element's atom. The period number is the same as the number of occupied shells that an element's atom has.

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Isotopes are atoms of the same element that have the same atomic number but different mass number.

Each isotope has the same number of protons and electrons, but a different number of neutrons. For example carbon has three main isotopes: carbon-12, carbon-13 and carbon-14.

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Ions and Ionic Bonding

An ion is a charged atom or group of atoms. A positive ion is made when an atom loses one or more electrons. A negative ion is made when an atom gains one or more electrons.

Ionic Bonding:

  • the metal atoms loses all outer-shell electrons to become a positive ion
  • the non-metal atom gains electrons to fill its outer shell and becomes a negative ion
  • the positive and negative ions are attracted to each other. This is an ionic bond.
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Sodium Chloride and Magnesium Oxide

NaCl and MgO form giant ionic lattices in which positive and negative ions are strongly attracted to each other. This means that they:

  • have high melting points as there is a strong attraction between oppositely charged ions
  • can conduct electricity when molten or aqueous because the charged ions are free to move
  • don't conduct electricity when solid because the ions are held in place and can't move

But MgO has a higher melting point tha NaCal as the ionic bonds are stronger and need more energy to be broken.

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Groups and Periods

Groups - A vertical column of elements in the periodic table. Elements in the same group have similar chemical properties because they have the same number of electrons in their outer shell. This outer number of electrons is the same as their group number.

Periods - a horizontal row of elements in the periodic table. The period for an element is related to the number of occupied electron shells it has.

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Covalent Bonding

Covalent bonding is when non-metals combine by sharing pairs of electrons. Simple covalently bonded molecules have low melting points. They don't conduct electricity because there aren't any free electrons.

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Development of the Periodic Table

Dobereiner was the first to suggest a Law of Triads, where he grouped similar elements into set of three with similar properties. The middle element would have the average mass of the other two elements. However, not all the elements were known and the pattern did not work for every known element.

John Newlands was the first scientist to make a table of elements, which he called the Law of Octaves, where every eighth elements behaved the same. But he included some compounds which he thought were elements.

Mendeleev was the author of the modern periodic table. He left gaps in his table for the unknown elements and made predictions about their properties. His predictions were later proved correct. Also, investigations on atomic structure agreed with his ideas.

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Group 1 - The Alkali Metals

The alkali metals are found in group one of the periodic table. The first three elements in the group are lithium, sodium and potassium. They all have one electron in their outer shell so all share similar properties.

Alkali metals are stored under oil because they react with air and react vigorously with water.

Alkali metals react with water to produce hydrogen and a hydroxide. Alkali metal hydroxides are soluble and form alkaline solutions, which is why they are known as the alkali metals.

The alkali metals react more vigorously as you go down the table - lithium reacts gently; sodium reacts more violently than lithium; potassium reacts more violently thatn sodium (it melts and burn with a lilac flame).

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Flame Tests

Lithium, sodium and potassium compounds can be recognised by the colours they make in a flame test:

  • A piece of clean nichrome wire is dipped in water
  • The wire is dipped in the solid compound, the wire is then put into a Bunsen flame
  • Each compound will produce a different coloured flame

Lithium - red flame

Sodium - yellow flame

Potassium - lilac flame

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Properties of the Alkali Metals

Rubidium is the fourth elements in group 1. Rubidium's reaction with water is:

  • very fast
  • exothermic (gives out energy)
  • violent

Density increases as you go down the group (except potassium). Caesium has the greatest density, and lowest melting and boiling points.

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Trends in Group 1

Alkalie metals have similar chemical properties because as they react, each atom loses one electron from its outer shell. So, a positive ion with a stable electronic structure is made.

The alkalie metals become more reactive as you go down the group because the outer shell gets further away from the positive attraction of the nucleus. This makes it easier for an atom to lose an electron from its outer shell.

Oxidation involves the loss of electrons by an atom. If the ionic equation for a reaction shows that an electron has been lost, an oxidation reaction took place.

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Group 7 - The Halogens

The five non-metals in group 7 are known as the halogens. They have all seven electrons in their outer shell so they have similar chemical properties. Fluorine, chlorine, bromine and iodine are halogens; at room temperature:

  • Chlorine is a green gas - used to sterilise water and make pesticides and plastics
  • Bromine is an orange liquid
  • Iodine is a grey solid - used as an antiseptic to sterilise woounds
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Displacement Reactions

The reactivity of the halogens decreases as you go down the group, so fluorine is the most reactive and iodine is the least.

A more reactive halogen will displace a less reactive halogen from an aqueous solution of its metal halide:

  • Chlorine will displace bromides and iodides
  • Bromine will displace iodides

If chlorine gas was passed though an aqueous solution of potassium bromide, bromine and potassium chloride would be made in the displacement reaction:

Potassium + Chlorine --> Potassium Chloride + Bromine

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Properties of the Halogens

The physical and chemical properties of the halogens change as you go down the group. Fluorine is the most reactive element in the group; it will displace all of the other halogens from an aqueous solution of their metal halides.

Astatine is a sem-metallic, radioactive element and only very small amounts are found naturall. It's the least reactive of the halogens, and theoretically, it would be unable to displace any of the other halogens (in an aqueous solution of their metal halides).

Astatine is very unstable and difficult to study.

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Trends in Group 7

The halogens have similar chemical properties because, as they react, each atom gains one electron to form a negative ion with a stable electronic structure.

Reduction involves the gain of electrons.

The halogens at the top of the group are more reactive than those at the bottom because the outer shell is closer to the positive attraction of the nucleus. This makes is easier for an atom to gain an electron.

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Transition Metals

The transition metals are a block of metallic elements between groups 2 and 3 on the periodic table. Transition metals have typical properties of metals. Their compounds are often coloured, for example:

  • Copper compounds are blue
  • Iron (II) compounds are light green
  • Iron (III) compounds are orange-brown

Many transition metals and their compounds are catalysts:

  • Iron is used in the Haber process
  • Nickel is used in the manufacture of margarine
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Thermal Decomposition

Thermal decomposition is a reaction where a substance is broken down into two or more substances by heating.

When transition metal carbonates are heated, a colour change happens. They decompose to form a metal oxide and carbon dioxide (the test for carbon dioxide is that it turns limewater milky).

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Identifying Transition Metal Ions

Precipitation is the reaction between solutions that makes an insoluble solid. The insoluble solid is known as a precipitate.

  • Copper (II) - blue precipitate
  • Iron (II) - grey/green precipitate
  • Iron (III) - orange precipitate
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Transition metals have many uses, for example:

  • Iron is used to make steel - which is used to make cars and bridges as it's very strong
  • Copper is used to make electrical wiring because it's a good conductor.

Metals are very useful because of their properties:

  • They're lustrous e.g. gold is used in jewellery
  • They're hard and have a high density
  • They have high tensile strength (able to bear loads)
  • They have high melting/boiling points
  • They're good conductors of heat and electricity
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Structure of Metals

Metal atoms are packed very close together in a regular arrangement. The atoms are held together by metallic bonds.

Metals have high melting/boiling points because lots of energy is needed to break the strong metallic bonds. As the metal atoms pack together, they build a structure of crystals.

Metal crystals are made from closely packed positive metal ions in a 'sea' of delocalised electrons. The free movement of the electrons allows the metal to conduct electricity.

The metal is held together by strong forces called metallic bonds. These are electrostatic attratctions between the metal ions and the delocalised electrons (so they have high melting/boiling points).

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Metals are able to conduct electricity because electrons flow easily through them, moving from atom to atom. At low temperatures, metals can become superconductors.

A superconductor has little, or no, resistance to the flow of electricity. This is useful for:

  • powerful electromagnets e.g. inside medical scanners
  • very fast electronic circuits e.g. in a supercomputer
  • power transmission that doesn't lose energy

The disadvantage of current superconductors is that they only work at temperatures below -200 degrees. This low temperature is very costly to maintain and is impractical for large-scale uses. There is a need to develop superconductors that will work at room temperature.

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The main sources of water in the UK are:

  • rivers
  • lakes
  • reservoirs
  • aquifers (wells and bore holes)

Water is an important resource for industry as well as being essential for drinking, washing etc. The chemical industry uses water as a coolant, a solvent and as a raw material.

In some parts of Britain, the demand for water is higher than the supply so it's important to conserve water.

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Water Treatment

Water has to be treated to purify it and make it safe to drink. Untreated water can contain:

  • insoluble particles
  • pollutants
  • microorganisms
  • dissolved salts and minerals

Tap water isn't pure - it contains soluble materials that aren't removed by the normal water treatment process (these materials could be poisonous). To obtain pure water, it must be distilled but this requires a lot of energy and is expensive.

Water treatment process:

  • Sedimentation - the water settles to allow the insoluble particles to sink
  • Filtration - to remove very fine particles
  • Chlorination - to kill the microorganisms in the water
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Pollutants in Water

Pollutants that can be found in water supplies are often difficult to remove. They include:

  • Nitrates from the run-off of fertilisers
  • Lead compounds from old pipes in the plumbing
  • Pesticides from spraying crops near to the water supply
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Dissolved Ions

The dissolved ions of some salts are easy to identify as they will undergo precipitation reactions.

Sulfates can be detected using barium chloride solution

Silver nitrate solution is used to detect halide ions

With silver nitrate:

  • Chlorides form a white precipitate
  • Bromides form a cream precipitate
  • Iodides form a pale yellow precipitate
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