1.3 Structures 💎

Structure

how atoms are arranged once they are bonded together

Giant structure

3D network of atoms or ions

Lattice

Regular arrangement of particles

Type of structure

depends on its bonding

Metallic structure

regular lattice of positive metal ions, surrounded by sea of delocalised electrons

Examples of metallic

magnesium or sodium

Delocalised electrons

electron able to move freely and carry charge

Properties of metallic structures

• high melting/ boiling points as high temperatures required to break strong electrostatic forces

• good conductivity as delocalised electrons can move, carry charge and transfer heat energy quickly

• layers of ions can slide over each other, but still held together by delocalised electrons so is both malleable and ductile

Malleable

can be hammered into shape

Ductile

can be drawn out into wires

Alloy

mixture of two or more elements, one of which a metal, and resulting mixture has metallic properties

Reason alloys are harder than pure metals

different sized atoms distort layers of ions in the structure, preventing layers from sliding

Gold used in jewellery

usually an alloy with silver, copper and zinc

Proportion of gold in alloys

measured in carats, 24ct being pure and 18ct being 75%

Molecular covalent structures (most common)

simple molecules held together by weak forces of attraction

Examples of molecular covalent

iodine or carbon dioxide

Weak forces of attraction

Van der Waals’ forces, are stronger the larger the molecule

Properties of molecular covalent substances

• low melting and boiling points as small amount of energy is need to break weak forces between molecules

• don’t conduct electricity as molecules are neutral so no ions/ delocalised electrons to carry charge

• most are insoluble in water

Ammonia

has molecular covalent structure but as a whole is charged, so will form ionic bonds to non metals

Giant covalent structures

3D structure of atoms joined together by lots of strong covalent bonds

Examples of giant covalent structures

diamond, graphite or graphene

Allotropes

different forms of same element in the same physical state e.g diamond, graphite and graphene

Carbon atoms have

up to 4 covalent bonds (shares all outer electrons)

Diamond

each carbon atom is covalently bonded to 4 other in 3D tetrahedral structure

Properties of diamond

• high melting and boiling points as large amount of energy needed to break many strong covalent bonds

• doesn’t conduct electricity as no ions or delocalised electrons can carry charge

• very hard due to strong covalent bonds in rigid tetrahedral structure

Uses of diamond

cutting tools for hard materials and jewellery

Graphite

carbon atom forms 3 covalent bonds (+1 delocalised electron) to form layers of hexagonal rings, held by weak forces of attraction

Properties of graphite

• high melting and boiling points as large amount of energy needed to break many strong covalent bonds

• conducts electricity as each atom has 1 delocalised electron that is free to move through layers and carry charge

• soft due to weak forces between layers, allowing them to slide easily

Uses of graphite

lubricant and pencils as carbon slides onto page

Graphene

single-atom thick layer of graphite with strong covalent bonds between each atom, arranged in hexagons

Properties of graphene

• high melting and boiling points as large amount of energy needed to break many strong covalent bonds

• excellent conductor of electricity and heat as each atom has 1 delocalised electron that is free to move through layers and carry charge

• very strong (100x steel) due to many covalent bonds

• low density and transparent as is only single atom thick

• flexible and stretchy as layer of atoms can bend without breaking bonds

Uses of graphene

batteries and solar cells due to good electrical conductivity

Giant ionic lattice

huge 3D regular structure held together by the electrostatic forces of attraction between positive and negative ions

Examples of ionic compounds

sodium chloride or magnesium oxide

Properties of ionic compounds

• high melting and boiling points as large amount of energy needed to break strong ionic bonds

• do not conduct electricity when solid as ions are held in fixed positions, cannot move and carry charge

• conduct electricity when molten or dissolved in water as ions are free to move and carry charge

• most are soluble in water

• brittle as when struck layers of ions move beside one another, opposite charges repel, so shatters crystal structure

Ions when molten/ dissolved in solution

break free so can move and carry charge

Electrostatic force of attraction

attractive force between oppositely charged particles

Reason some molecules are soluble

react with water, usually only giant ionic structures but also acids

Classifying structures

Structure	Melting point	Solubility in water	Electrical conductivity
Giant ionic lattice	High	Soluble	Conductive when molten or dissolved
Molecular covalent	Low	Insoluble	Not conductive
Giant covalent	Very high	Insoluble	Not conductive (except graphite/ graphene)
Metallic	Usually high	Insoluble	Conductive

Classifying example

• A is a metallic structure

• B is a giant ionic lattice

• C is a giant covalent structure