4.1-Metals and Non-Metals

Most elements are metals

• Metals are elements which can for positive ions when they react

• They’re towards the bottom and to the left of the periodic table

• Most elements in the periodic table are metals

• Non-metals are at the far right and top of the periodic table

• Non-metals don’t generally form positive ions when they react

The electronic structure of atoms affects how they will react

• Atoms generally react to form a full outer shell

  • They do this via losing, gaining or sharing electrons

• Metals to the left of the periodic table don’t have many electrons to remove

• Metals towards the bottom of the periodic table have outer electrons which are a long way from the nucleus

  • They feel a weaker attraction

  • Both these effects means that not much energy is needed to remove the electrons so it’s feasible for them to either share or gain electrons to get a full outer shell

Metals and non-metals have different physical properties

• All metals have metallic bonding which causes them to have similar basic physical properties

  • They’re strong but can be bent or hammered into different shapes

  • They’re great at conducting heat and electricity

  • They have high boiling and melting points

• As non-metals don’t have metallic bonding, they don’t tend to exhibit the same properties as metals

  • They tend to be dull looking, more brittle, aren’t always solids at room temperature, don’t generally conduct electricity, and often have a lower density

Transition metals can be found between group 2 and group 3

• Transition metals are in the centre of the periodic table

• Transition metals are typical metals and have the properties you would expect of a proper metal

  • They’re good conductors or heat and electricity and they’re very dense, strong and shiny

• Transition metals also have some pretty special properties

  • Transition metals can have more than one iron

    • Copper forms Cu+ and Cu2+ ions

    • Cobalt forms Co2+ and Co3+ ions

  • Transition metal ions are often coloured, and so compounds that contain them are colourful

    • Potassium chromate which yellow and potassium manganate is purple

  • Transition metal compounds often make good catalysts

    • Nickel based catalysts are used in the hydrogenation of alkenes

    • Iron catalyst is used in the haber process for making ammonia

4.2-Group 1 Elements

The group 1 elements are reactive, soft metals

• The alkali metals are lithium, sodium, potassium, rubidium, caesium and francium

• They all have one electron in their outer shell which makes them very reactive and gives them similar properties

• The alkali metals are all soft and have low density

• The trends for the alkali metals as you go down Group 1 include:

  • Increasing reactivity

    • The outer electron is more easily lost as the attraction between the nucleus and electron decreases, because the electron is further away from the nucleus and the further down the group you go

  • Lower melting and boiling points

  • Higher relative atomic mass

Alkali metals form ionic compounds with non-metals

• The Group 1 elements don’t need much energy to lose their one outer electron to form a full outer shell, so they readily form 1+ ions

• It’s so easy for them to lose their outer electron that they only ever react to form ionic compounds.

• These compounds are generally white solids that dissolve in water to form colourless solutions

• Reaction with water

  • When Group 1 metals are put in water, they react vigorously to produce hydrogen gas and metal hydroxide

    • Salts that dissolve in water to produce alkaline solutions

  • The more reactive(lower down in the group) an alkali metal is, the more violent the reaction

  • Sodium + Water - Sodium Hydroxide + Hydrogen

  • 2Na + 2H20 - 2NaOH + H

• Reaction with chlorine

  • Group 1 metals react vigorously when heated in chlorine gas to form white chloride salts

  • As you go down the group, reactivity increases so the reaction with chlorine gets more vigorous

  • Sodium + Chlorine - Sodium Chloride

  • 2Na + Cl - 2NaCl

• Reaction with oxygen

  • The Group 1 metals can react with oxygen to form a metal oxide:

  • Different types of oxide will form depending on the Group 1 metal:

    • Lithium reacts to form lithium oxide

    • Sodium reacts to form a mixture of sodium oxide(Na20) and sodium peroxide(Na2O2)

    • Potassium reacts to form a mixture of potassium peroxide(K202) and potassium superoxide(K02)

Group 1 metals have different properties to transition metals

• Group 1 metals are much more reactive than transition metals-they react more vigorously with water, oxygen or Group 7 elements, for example

• They’re also much less dense, strong and hard than the transition metals, and have much lower melting points

  • e.g. manganese melts at 2000C

  • Sodium melts at 98C

4.3-Group 7 Elements

The halogens are all non-metals with coloured vapours

• Fluorine is a very reactive, poisonous yellow gas

• Chlorine is a fairly reactive, poisonous dense green gas

• Bromine is a dense, poisonous red-brown volatile liquid

• Iodine is a dark grey crystalline solid or a purple vapour

• They all exist as molecules which are pairs of atoms

  

Learn these trends

• As you go down Group 7 the halogens

  • Become less reactive

    • It’s harder to gain an extra electron, because the outer shell’s further from the nucleus

    • Have higher melting and boiling points

    • Have higher relative atomic mass

• All the Group 7 elements react in similar ways.

• This is because they all have seven electrons in their outer shell

Halogens can form molecular compounds

• Halogen atoms can share electrons via covalent bonding with other non-metals so as to achieve a full outer shell

  • For example, HCI, PCI5, HF and CCI4 contain covalent bonds

  • The compounds that form when halogens react with non-metals all have simple molecular structures

Halogens form ionic bonds with metals

• The halogens from 1-ions called halides

  • F-, CI-, Br- and I-

• When they bond with metals

  • Na+Ci or Fe3+Br-3

• The compounds that form have ionic structures

• The diagram shows the bonding in sodium chloride, NaCI

More reactive halogens will displace less reactive ones

• A displacement reaction can occur between a more reactive halogen and the salt of a less reactive one

  • E.g. Chlorine can displace bromine and iodine form an aqueous solution of its salt(a bromine or iodine)

  • Bromine will also displace iodine because of the trend in reactivity

  • Cl2 + 2KI - I2 + 2KCI

  • Pale green - Brown

  • CL2 + 2KBr - Br2 + 2KCI

  • Pale green - Orange

4.4-Group 0 Elements

Group 0 elements are all inert, colourless gases

• Group 0 elements are called the noble gases and include the elements helium, neon and argon(and a few others)

• They all have eight electrons in their outer energy level, apart from helium which has two, giving them a full outer shell

• As their outer shell is energetically stable they don’t need to give up or gain electrons to become more stable

  • This means they are more or less inert-they don’t react with much at all

• They exist as monatomic gases-single atoms not bonded to each other

• All elements in Group 0 are colourless gases at room temperature

• As the noble gases are inert they’re non flammable

  • They won’t set on fire

There are patterns in the properties of the noble gases

• The boiling point of the noble gases increase as you move down the group along with increasing relative atomic mass

  • The increase in boiling point is due to an increase in the number of electrons in each atom leading to greater intermolecular forces between them which need to be overcome

• In the exam you may be given the boiling point of one noble gas and asked to estimate the value for another one

  • Neon is a gas at 25C. Predict what state helium is at this temperature

    • Helium has a lower boiling point than neon as it is further up the group

    • So, Helium must also be a gas at 25C