bonding

ionic bonding

• Between metals and non-metals 

• Electrons are transferred from metal → non-metal 

• Forming ions, metal forms positive, non-metal forms negative ions

• These ions are attracted to each other as they are oppositely charged 

• Electrostatic force of attraction

strength of ionic bonding (atomic radius)

Smaller ions means stronger ionic bonding. 

Ions are packed more closely together in a lattice 

shorter distance between oppositely charged ions meaning electrostatic forces are stronger

Ions bigger down group, more shells, more shielding down the group so weaker bonding 

Ions smaller across period as they have a higher nuclear charge

strength of ionic bonding (charge on ions)

Higher charge, stronger bonding 

Ions exert a greater force into each other 

covalent bonds 

• Between 2 non-metals 

• Atoms share pairs of outer electrons 

• A covalent bond is a shared pair of electrons 

• Molecules are neutral as no electrons are transferred 

• Atoms held by electrostatic forces between shared pair of electrons and nuclei 

simple covalent molecules

• Molecules have covalent bonds between atoms to hold them together. 

• Between molecules are intermolecular forces which are the bonds that need breaking to melt/boil the compound. 

DATIVE COVALENT BONDS

• One atom provides both the electrons in a covalent bond 

• Atom that accepts the electrons is electron–deficient 

• Atom donating electrons must have a lone pair 

• Charge of the receiving atom remains 

metallic bonding

• Positive metal ions are attracted to the negative delocalised electrons. 

• This is an electrostatic force of attraction. 

metal structure

• Metals exist as giant lattice structures. 

• The outermost electron on a metal atom is delocalised as is free to move. 

• This leves positive metal ions. 

• The positive metal ions are attracted to the delocalised negative electrons. 

• The positive ions are packed closely together in a sea of delocalised electrons.

structure of diamond

Tetrahedral shape as each carbon atom is covalently bonded to 4 other carbon atoms.

melting point of diamond

Very high due to the strong covalent bonds between the carbon atoms. 

conductivity of diamond

Cannot conduct electricity as there are no delocalised electrons. 

structure of graphite

Carbon atoms are arranged in sheets of hexagonal layers each carbon bonding to 3 others. The 4^th outer electron of each carbon is delocalised. The sheets of hexagons are connected with weak van der waals.

melting point of graphite

Very high due to the strong covalent bonds between the carbon atoms.

conductivity of graphite

Can conduct electricity due to having delocalised electrons that can move and carry charge through the structure.

repulsion

Lone pairs repel more than bonding pairs as they are closer to the nucleus of the central atom than the bonding pairs.

linear

electron pairs - 2

bonding pairs - 2

lone pairs - 0

bond angle - 180

linear diagram

X – Q – X

trigonal planar

 

 

electron pairs - 3

bonding pairs - 3

lone pairs - 0

angle - 120

v shape

electron pairs - 3

bonding pairs - 2

lone pairs - 1

angle - 118

trigonal bipyramidal

electron pairs - 5

bonding pairs - 5

lone pairs - 0

angle - 120,90

see - saw

electron pairs -5

bonding pairs - 4

lone pairs - 1

angle - 89,119 

t shape

electron pairs - 5

bonding pairs - 3

lone pairs - 2

angle - 89 

octraheral

electron pairs - 6

bonding pairs - 6

lone pairs - 0

angle - 90

square pyramid 

electron pairs - 6

bonding pairs - 5

lone pairs - 1

angle - 89

square planar

electron pairs - 6

bonding pairs - 4

lone pairs - 2

angle - 90

electronegativity

ELECTRONEGATIVITY is the power of an atom to attract the electron density/ electron pair towards itself in a covalent bond. 

factors affecting electronegativity

atomic radius

nulcear charge 

electronegativity - atomic radius

The smaller the atomic radius, the higher the electronegativity. 

There is a stronger attraction between the electrons and the nucleus

electronegativity - nuclear charge

The higher the nuclear charge, the higher the electronegativity. 

More protons in the nucleus create a stronger attraction with the electrons. 

A higher nuclear charge also decreases atomic radius.

electronegativity down the group

• Atomic radius increases, distance between electrons and nucleus increases. Shielding increases. 

Less attraction.  

electronegativity across the period

• Atomic radius increases, distance between electrons and nucleus increases. Shielding increases. 

• Less attraction.  

polarity

Polarity is about unequal sharing of electrons between atoms that are bonded together.

dipole

differences in charges caused by electron shift. 

when do polar molecules occur ?

DIFFERENT ATOMS 

• Different electronegativities so unequal sharing of electrons, dipole moment 

• Electrons more attracted to the more electronegative atom , delta +

• The less electronegative atom becomes delta - 

• The greater the difference in electronegativity, the more polar 

when do non -  polar molecules occur ?

SAME ATOMS 

• Electrons shared equally as equal attraction to both atoms 

• Same electronegativity 

• C-H is considered non-polar as electronegativities are almost equal.

polarity of molecules (polar)

Has polar bonds which do not cancel out 

H₂O

polarity of molecules (non - polar) no polar bonds

Contains no polar bonds, symmetrical molecule 

CH₄

polarity of molecules (non - polar) polar bonds

Contains polar bonds which cancel out, symmetrical molecule  

CO₂

VAN DER WAALS (INDUCED DIPOLE-DIPOLE ATTRACTION) 

• Occur because electrons are constantly moving around 

• This causes uneven distribution of electrons 

• Inducing a temporary dipole (some parts are charged due to electron density) 

• Temporary dipole induces a temporary dipole in the neighbouring molecule

• This causes an attraction between the delta + and - charges - induced dipole-dipole

PERMANENT DIPOLE-DIPOLE ATTRACTION

• In polar molecules

• Not in non-polar molecules even if have dipole bonds 

• Weak electrostatic forces of attraction between delta + and - of adjacent molecules 

• Delta + attracts delta - on next molecule/ bond

HYDROGEN BONDING

• When hydrogen is bonded to fluorine/oxygen/nitrogen 

• These elements are very electronegative so attract the electrons

IMPORTANCE OF H-BONDING IN ICE

• As a liquid, water forms hydrogen bonds which break and reform easily 

• When frozen, the molecules cannot move

• So the H-bonds hold the molecules in a fixed position 

• Ice is a 3D structure like a diamond 

• In order to fit, the molecules of water are packed less closely together 

• This means the ice is less dense than water and floats on top of it 

• This is why there is like beneath an icy surface which acts as an insulator

BOILING POINT OF HYDRIDES OF GROUP 4,5,6 AND 7

• Boiling points of H₂O, HF and NH₃ are the highest as they have H-bonding 

  • Hydrogen bonding requires more energy to overcome

• H₂O is the highest as it can form 4 H-bonds, HF and NH₃ only form 2

• Boiling points generally increase down a group due to more electrons and more van der waals 

• Group 4 hydrides have lowest b.p as they only have van der waals due to having almost the same electronegativity as hydrogen - nonpolar molecules 

• Group 5,6,7 have permanent dipole-dipole or hydrogen bonds which are stronger