1.3 Bonding 1.5 Structure CCEA

Metallic Bonding definition

The attraction between layers of cations and a sea of delocalised electrons

Structure of metallic bonding

A metallic lattice ( regular arrangement of ions )

Properties of metals

1. Electrical conductivity

2. High melting and boiling point

3. Malleable and ductile

4. Hardness

Metallic electrical conductivity

Electrons in metal delocalised they free to move and carry charge through lattice in certain direction when potential difference applied thus metals can conduct in solid and liquid state

High melting points

Large amount of energy is requires ro break the strong attractions between the positive ions and the negative electrons 

why does sodium have a higher melting/boiling point than potassium 

smaller ions with a high charge attract the electrons more strongly - higher melting points than larger ions with a lower charge .

Na has smaller cations than K -  so Na has higher Mp and Bp

why does magnesium and have a higher melting/boiling point than soduim

Mg cations have a higher charge than Na so Mg has higher Mp and Bp

Malleable and ductile

Layers of positive ions can slide over each other without disrupting the bonding

Hardness in metallic bonding

Strong attraction between positive ions and negative electrons, and strong regular structure

Ionic bonding 

The electrostatic force of attraction between ions of opposite charge in a reuglar ionic lattice.

Metal atoms to non-metal atoms

Structure of ionic bodning

an ionic lattice

ionic bonding

Electrostatic forces increase in strength as

the charge on the ions increases

the size of the ions decreases

1. Properties of ionic compounds

1. High melting and boiling points

2. Can’t conduct electricity in solid state, can molten or aqueous

3. Usually white, crystalline solids

High melting points of ionic compounds

Requires lots of energy to break strong electrostatic forces of attraction between ions of opposite charge

Conductivity in ionic compounds

Not in solid state but when molten or aqueous

Ions are not free to move about in the solid state but are free when the solid is melted or dissolved in water

Usually white, crystalline solid

Ionic compounds exist as regular arrangements of ions in a giant ionic lattice.

Regular arrangement leads to hard crystal structure being formed

Covalent bond

The electrostatic attraction between a shared pair of electrons and the nuclei of bonded atoms

Structure of covalent bonds

There are two types of covalent structure (simple covalent) and giant covalent.

What is a lone pair

A pair of unshared electrons

Why do covalent bonds happen

the electrons are more stable when attracted to two nuclei than when attracted to only one

What is a coordinate bond/ dative covalent

A shared pair of electrons between two atoms with both electrons shared by one atom

Why is a coordinate bond formed

Because nitrogen shares a pair of electrons

ammonia bond

ammonium ion diagram 

dirrection of arrow shows where the 2 electrons have come from

What is the octet rule

When reacting, an atom tends to gain, lose or share electrons to achieve eight in its outer shell.

contraction of the octet rule + examples

has less than 8 electrons in outer shell

BeCl2, BF3

Expanision of the octet rule

Has more than 8 electrons in the outer shell

PCl5, SF6

Propeties of moelcular covalent sturctures

1. Do not conduct elecitricty

2. Low melting and boiling points

3. Solubility

Molecular covalent compounds conductivity

They cant as they do not contain ions or delocalised electrons

Low melting and boiling points in molecular covalent compounds

Little energy required to break the weak Van der Waals forces between molecules. Attractive forces between molecules containing larger molecules - S₈, P₄, or heavier atoms such as I₂ , are much greater than the attrcitve forces between molecules such as N₂. Consequently sulfur and iodine have higher m.p and are more likly to be solids or liquids at r.t.p

Solubility in molecular covalent compounds

Molecular substances containing polar molecules - ammonia dissolve in polar solvents such as water. Molecular substances containing less polar molecules - sulfur or iodine dissolve in less polar solvents such as hexane.

Structure of diamond (allotrope of carbon)

Each carbon atoms is covalently bonded to four others in a tetrahedral arrangement, to form a rigid three-dimensional structure

Diamond diagram

Properties of diamond

1. High melting and boiling point

2. Poor conductivity

3. Hard, strong and brittle

High melting and boiling point meaning behind diamond

SInce strong covalent bonds must be broken before any atoms can be sepearted which requires a large amount of energy

Diamonds poor conductivity

There are no ions or delocalised electrons, so there is little electrical conductivity in either solid or liquid state.

Diamonds hardiness, strong and brittle

The covalent bonds are strong and directional, lots of energy needed to break the bonds. Hardest substance - used n drills and glass-cutting.

Graphites structure

Each carbon atom is covalently bonded to three others. The spare delocalised electron occupies space above and below each layer, All atoms in same layer held together by strong covalent bonds, and the different layers are held together by intermollecular forces.

Properties of graphite

1. High melting point

2. Good electrical conductivity

3. Hardness

Graphite high melting point reason

Lots of energy required to break the strong covalent bonds

Graphite good electrical conductivity reason

due to delocaliseded electrons in each plane, graphite is very good conducutor of electricity within each layer, even in the solid state (ususally for a non-metal)

Graphite hardness

Graphite much softer than diamond since different planes can slip over eachother fairly easily. Used in pencils and industrial lubricant.

What is electronegativity

The extent to which an atom attracts the bonding electrons in a covalent bond

What two factors affect electronegativity

-As the size of the nuclear charge increases, increased attraction between nuclear charge and pair of electrons in the covalent bond, so electrongeavity increases

-as Size of atom increases pair of electrons in covalent bond further away from the nucleus , there will be decreased attraction from nuclear charge so electronegativity decreases

Electronegativity across period

Increases , nuclear charge increases size of atoms decreases hence greater attraction between the nucleus and the pair of electrons in the covalent bond 

Electronegativity down a group

Decreases - pair of electrons in covalent bond further from nuclear so decreased attraction from nuclear charge. inner electrons also shield the pair of electrons in the bond from the nuclear charge 

electronegativity values of two atoms same/similar

pair of electrons shared equally and a covalent bond is nonpolar e.g I₂, O₂

Electronegativity of N, O , F

N- 3.0

O- 3.5

F- 4.0

ELECTRONEGATIVTY TABEL

What is a polar bond

A covalent bond in which there is an unequal sharing of bonding electrons.

What is partial charges

Delta negative δ-, and delta positive (less electronegative)

How is a bond dipole formed

Seeration of charge that develops between atoms in polar covalent bond

size increases as difference in electronegativity values increases

Why does a polar bond cancel

The shape of the molecule is symmetrical and the polarities of the bond cancel out

different electronegativities cancel out

Sodium, magneisum and aluminium meaning

Sodium magneisum and aluminium are metals (metalic structure). Strenght of metalic bonding depends upon the number of delocalsied electrons in metalic structure. Soidum has one, Mg 2, Al 3. Hence M.P and B.P increase from sodium to aluminium

Silicone meaning

Silicone exists as giant covalent strucutre. Great number of strong covalent bonds betweent the atoms have to be broken so high M.P and B.P

Phosphorus, sulfur, chlroine

are simple molecular species (molecular covalent). P consists of P₄ molecules, S - S₈ molecules, Cl - Cl₂ molecules. Forces between molecules are weak Van der Waals forces, so these elements have low melting and boiling points. Strength of VDW forces increases as size of the molecules increases more energy is needed to break these forces, so M.P and B.P increases from Cl₂ , P₄, S₈

Argon

Argon exists as isolated atoms (monotamic) with weak Van der Waals forces between the atoms. therefore Ar has lowest M.P and B.P in period.