structure and physical properties

Question 1

what is metallic bondings

Answer 1

• the strong electrostatic attraction between the positive ions and the sea of delocalised electrons
• layers of positive ions surrounded by a sea of delocalised electrons

Question 2

type of structure

Answer 2

giant structures
lattice as they are arranged in a regular pattern

Question 3

properties of metallic bonding

Answer 3

• high melting and boiling points because lots of energy is needed to overcome the strong electrostatic forces of attraction between the positive ions and the delocalised electrons
• can conduct electricity and heat
• malleable - the layers of positive ions can slide over each other
• metals do not dissolve. there is some interaction between polar solvents and charges in the metallic lattices but these lead to reactions, rather than dissolving e.g. sodium and water

Question 4

why does the melting point increase across the groups

Answer 4

• the higher the charge of the positive ion
• the stronger the forces of attraction between the ion and the delocalised electrons
• because you will have more delocalised electrons per positive ion
• so the higher the melting point

Question 5

why can metals conduct electricity and head

Answer 5

the delocalised electrons can move and carry charge

Question 6

what are giant covalent substances

Answer 6

• the bonds between atoms continue indefinitely and a large lattice is formed
• there are no individual molecules and covalent bonding exists between all adjacent atoms

Question 7

what are examples of giant covalent substances

Answer 7

• allotropes of carbon
• graphite
• diamond
• graphene
• silicon dioxide

Question 8

structure of diamond

Answer 8

• made of up one carbon atom strongly bonded by covalent bonds to four other carbon atoms with a bond angle of 109.5°
• formed a structure called a a tetrahedral arrangement - repeated millions of times to create a strong giant covalent lattice
• this structure makes diamond the very hardest natural substance
• diamond has a very high melting and boiling points
• cannot conduct electricity - it has no free moving ions and electrons

Question 9

structure of graphite

Answer 9

• made up of one carbon atom covalently bonded with **three other atoms **with a bond angle of 120°
• this forms interlocking rings of six atoms, which create a giant structure made of many layers
• these layers are held together by weak forces of attraction - intermolecular forces
• in these layers delocalised electrons are found that are free to move
• this means that graphite is soft and slippery because the layers can slide over each other
• it can also conduct electricity because the free electrons between the layers can carry a charge
• high melting and boiling point - giant structure

Question 10

what is the only giant or simple covalent structure to conduct electricity

Answer 10

graphite

Question 11

structure of silicon and silicon dioxide

Answer 11

• both silicon and silicon dioxide have diamond-like structure
• in silicon dioxide, each silicon atom is surrounded tetrahedrally by 4 oxygen atoms so that each silicon atom is bridges to its neighbouring silicon atom by an oxygen atom
• silicon dioxide is hard, has a high melting point and does not conduct electricity
• it is insoluble in water and in organic solvents
• each silicon is shared by four oxygens and each oxygen is shared by two silicons
• gives the empirical formula of SiO2

Question 12

properties of giant covalent substances

Answer 12

• GCL have very high melting and boiling points as these compounds have a large number of covalent bonds linking the whole structure. A lot of energy is required to break the lattice
• compounds can be hard or soft
  graphite is soft as the intermolecular forces between the carbon layers are weak
  silicon oxide and diamond are hard as it is difficult the break their 3D network of strong covalent bonds
• most compounds are insoluble in water
• most covalent substances do not conduct electricity
  e.g. diamond and silicon (IV) oxide do not conduct electricity as all four outer electrons on every carbon atom is involved in a covalent bond, so there are no free electrons available
• there are some covalent substances that are exceptions because they do conduct electricity
  graphite has delocalised electrons between the carbon layers, which can move along the layers when a voltage is applied
  graphene is an excellent conductor of electricity due to the delocalised electrons