Here’s what we’re covering for today:
- Ionic Bonding
- Covalent Bonding (intro)
- Covalent bonding
- Co-ordinate (dative)
- Metallic Bonding
What is Ionic Bonding?
A bond formed due to the attraction between positive & negative ions.
Also known as an ELECTROVALENT bond.
Ionic bonding occurs when:
- An atom LOSES electrons to form a CATION in order to gain a stable octet arrangement
- Energy is absorbed to lose the electron: this amount is the ionisation energy
- It donates the electrons to another atom
- The other electron RECEIVES electrons to form an ANION in order to gain a stable octet arrangement
- Energy is release to accept the electron: this amount is the electron affinity
- The oppositely-charged ions attract each other, & a ionic compound is formed
*Once you have learnt about ELECTRONEGATIVITY, you can revisit this section:
Q: If ionic bonds are between a metal & non-metal, why do some metals form covalent bonds (such as AlCl3)?
A: Ionic bonding can be explained in terms of atoms having a LARGE difference in electronegativity.
If there is a LARGE difference in electronegativity,
the valence electrons of the more electropositive atom (the metal) will tend to be attracted A LOT CLOSER to the more electronegative atom (the non-metal),
to the point where the electron can be considered to be DONATED.
If there is a SMALL difference in electronegativity,
the valence electrons of the more electropositive atom (the metal) will tend to be NOT AS CLOSE to the more electronegative atom (the non-metal),
to the point where the electron can be considered to be SHARED between the two.
In reality, there is no difference between ionic & covalent bonding.
Both lie on a spectrum, depending on electronegativity:
A visualisation of these 3 parts of the spectrum:
(it’s in the opposite order than the picture above. Apologies.)
What is Covalent Bonding?
A bond formed due to the sharing of valence electrons.
When two atoms share electrons, this means the electrons can exist in a space around BOTH atoms.
This will be discussed in-depth later, through the overlapping of orbitals.
Simple Covalent Bonding
Two or more atoms each contribute equal number of electrons to achieve stable octet arrangement.
Atoms share PAIRS of electrons.
Each pair of shared electrons is known as 1 bond:
- single-bond denotes 1 pair (e.g.: F2),
- double-bond denotes 2 pairs (e.g.: O2),
- triple-bond denotes 3 pairs (e.g.: N2), etc.
The electrons of each atom that are NOT involved in bonding are called
LONE ELECTRONS, & are usually in pairs called LONE PAIRS.
Co-ordinate Bonding (dative covalent bonds)
When ONE atom in a compound provides BOTH the electrons needed for a covalent bond.
In a simple covalent bond, both atoms provide electrons to be shared.
In a co-ordinate bond, ONLY 1 atom does so.
This occurs when:
- 1 atom has a LONE pair of electrons (2 electrons not bonded yet)
- The other atom has an UNFILLED orbital
In structural diagrams, you can show a co-ordinate bond by drawing an ARROW from the provider of the electrons.
Formation of Ammonium Ion, NH4+
|In ammonia (NH3),
nitrogen has a lone pair
|A H+ ion has an UNFILLED shell|
|Nitrogen provides the 2 lone electrons to be shared with the hydrogen ion|
|A positively charged ammonium ion is formed|
Formation of Aluminium Chloride Dimer, Al2Cl6
|At high temperatures:
AlCl3 exists as a MONOMER1 Al atom is covalently bonded with 3 Cl atoms.
|The Al atom only has 6 valence electrons,
it is ELECTRON-DEFICIENTHowever, this is stable at high temperatures.
|At low temperatures:
2 AlCl3 monomers will form coordinate bonds to become a Al2Cl6 DIMER.
|A Cl atom provides the lone pair for an
Al to complete an octet.The same occurs for the other Al.
What is Metallic Bonding?
Bond formed between atoms of metallic elements
In a metal,
- Atoms are packed closely in a lattice
- Atoms (which are electropositive) lose their valence electrons
- The electrons become DELOCALISED (free to move)
Each cation has a strong electrostatic bond with the sea of delocalised electrons.
The strength of a metallic bond is determined by:
- the number of electrons donated.
Higher number of donated electrons = higher attraction between cations & delocalised electrons
- the size of the atom.
Larger atomic radius = larger distance between protons & delocalised electrons
= weaker attraction
- the proton number of the metal.
Higher proton number = higher positive charge = stronger attraction.