A2/A-level Chemistry - Bonding

Ionic bonding

The electrostatic attraction between positive and negative ions in all directions. It holds together cation (positive ions) and anions (negative ions) in ionic compounds with a regular arrangement.  

Common cations: 

• metal ions 
• e.g.Na⁺, Mg²⁺, Al³⁺ 
• ammonium ions 
• e.g.NH₄⁺ 

Common anions:

• non-metalions 
• e.g.Cl^-, O^2- 
• polyatomicions 
• e.g.NO₃^-, SO₄^2- 

Ionic Compouds: Dot and Cross Diagram/ Lewis structure

• The simplest ionic compounds contain metal ions and non-metal ions 
• The outer shell electron(s) of a metal atom are transferred to the outer shell of a non-metal atom 
• Positive (loses electron(s)- metal) 
• and negative ions (gains electron(s)- non-metal) are formed 
• The ions formed often have outer shells with the same electron configuration as the nearest noble gas 

(Octet rule)

example 1: Sodium Chloride, NaCl 

• The outer electron of Na (sodium) is transferred to the outer shell of a chlorine atom, forming a sodium ion, Na⁺, and chloride ion, Cl^- 
• The electron structures of Na⁺ and Cl^- are now the same as the nearest noble gas, fullfill Octet rule.
• Square brackets are used to show that the charge is spread over each ion and the ions are separate entities  
• When it comes to AS/ A-level chemistry exam, only the outer shell electrons are shown as the inner shells are full and not involved in bonding 

example 2: Magnesium Chloride, MgCl₂

• The 2 electrons in the outer shell of the magnesium atom are transfered to the outer shells of 2 chloride atoms to form a magnesium ion, Mg2+, and 2 chloride ions, Cl-.
• Electron structure of each of the ions are therefore the same as the nearest noble gas, fullfilled octet rule. 

Structure of Ionic Compounds 

• Each ion attracts 
• oppositely charged ions 
• in all directions.
• Giant ionic lattice structure 

contains billions of billions of ions, the actual number of is only determined by the size of the crystal.

example: Sodium chloride, NaCl

• Each Na+ ion is surrounded by 6 Cl^- ions
• Each Cl- ion is surrounded by 6 Na⁺ ions
• Regular, cubic arrangement of Na⁺ and Cl^- ions gives the cubic shape of the crystal
• Each ion is surrounded by oppositely charged ions, forming a giant ionic lattice

Propterties of Ionic Compounds: 

Melting and Boiling points

• Most ionic compounds have 
• high melting and boiling points
• Melting points are higher for lattices containing ions with greater ionic charges as there is a stronger attraction between the ions
• Ionic attraction also depends of the size of the ions, the smaller the ions, the stronger the attraction because there is less shielding from the nucleus 

Common question in AS/ A-level paper: Why does ionic compounds have high melting and boiling point? 👨‍🏫

Answer: It is because a lot of energy is required to overcome the strong electrostatic attraction between the oppositely charged ions in the giant ionic lattice

Solubility 

Many ionic compounds dissolve in polar solvents (eg. water) 

• Polar water molecules break down the lattice and surround each ion in solution 
• In a compound make of ions with large charges, the ionic attraction may be too strong for water to be able to break down the lattice structure. The compound will not then be very soluble 

Solubility requires ✌️ processes:

1. The ionic lattice must be broken down
2. Polar molecules must attract and surround ions

• The solubilty of an ionic compound in water depends on the relative strengths of the attractions within the giant ionic lattice and the attractions between the ions and water molecules.
• Solubility decreases as ionic charge increases, but predictions of solubilty should be treated with caution

Electrical Conductivity

In solid states, ionic compounds do NOT conduct electrcity. This is because:

• the ions are in a fixed position in the giant ionic lattice
• there are no mobile charge carriers

In molten states (melted/ dissolved in water), ionic compound do conduct electrcity🔌, it is because:

• the solid ionic lattice breaks down
• the ions are now free to move as mobile charges carriers

Ionic compounds are conductrors of electricity in liquid and aquesous states.

Summary of Ionic compounds

Most ionic compounds:

• have high melting and boiling points.
• tends to dissolve in polar solvents such as water.
• conduct electricity only in the liquid state or in aquous solution. 

Covalent bonding

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

Common covalent compounds:

• non-metalic elements eg. H₂, O₂
• compounds of non-metallic elements eg. H₂O, CO₂
• polyatomic ions eg. NH₄⁺

The Covalent bond- Orbital Overlap

• A covalent bond is the overlap of atomic orbitals, each containing 1 electron, to give a shared pair of electrons
• The shared pair of elections is attracted to the nuclei of both the bonding atoms 
• The bonded atoms often have outer shells with the same electron structure as the nearest noble gas
• The attractoon in a covalent bond is localised, acting solely between the shared pair of electrons and the nuclei of the 2 bonded atoms. The result can be a small unit- a molecule, consisiting of 2 or more atoms. eg. H₂ and H₂O
• A molecule is the smallest part of a covalent compound that can exist whilst retaining the chemical properties of the coimpound. 

Single Covalent Bonds- Dot and Cross Diagrams 

• Used to account for electrons in covalent bonding 

            -in covalent bonding, electrons are shared

            -in ionic bonding, electrons are transferred 

• The use of dots and crosses allows the origin of each electron to be shown clearly 
• Each bonding structure will NOT have the electron structure of the nearest noble gas 
• Displayed formula shows the relative positioning of atoms and the bonds between them as lines 
• Lone pairs are paired electrons that are not shared. These can also be added to the displayed formulae as in water, H₂O, and ammonia, NH₃ 

Covalent Bonding in Boron 

• Period 2, Electron configuration 1s2,2s2,2p1 – so only 
• 3 outer shell electrons can be paired
• Boron forms covalent compounds e.g. borontrifluoride, BF₃, in which its 3 outer shell electrons are paired 
•  A molecule of BF₃ therefore only has 6 electrons around the boron atom 
•  BF₃ shows that predictions for bonding cannot be solely based on the noble gas electron structure  

Covalent Bodning in Phosphorus, Sulphur and Chlorine

•  Phosphorus trifluoride, PF₃, sulphurdioxide, SF₂, and chlorine monofluoride, ClF, follow the expected pattern of formulae, with the bonded atoms having a noble gas electron structure  
• For elements in period 2, the n=2 outershell can only hold 8 electrons 
• But for phosphorus, sulphur and fluorine,the n=3 outer shell can hold 18 electrons, so more electrons are available for bonding  

• Different numbers of unpaired electrons lead to different possibilities for covalent compounds of sulphur  
• Different arrangements for the 6 outer shell electrons of sulphur and the different number of bonds possible to bond with fluorine 

Multiple Covalent bonds: Double Covalent Bonds

• Multiple covalent bonds 
• exist when 2 atoms share more than 1 pair of electrons In a double covalent bond, the electrostatic attraction is between 2 shared pairs of electrons and the nuclei of the bonding atoms. e.g. molecule of O₂, O=O, CO₂, O=C=O 
• All atoms have 8 electrons in their outershell and the electron structure of the nearest noble gas C=C and C=O bonds are very important in organic chemistry.

Multiple Covalent bonds: Triple Covalent Bonds 

• In a triple covelent bond the electrostatic attraction is between 3 shared pairs of electrons and the nuclei of the bonded atoms. eg. N2, HCN
• All molecules have the electron structure of the nearest noble gas 

Dative Covalent Bonds 

A covlanet bond which the shared pair of electrons has been supplied by one of the bonding atoms only, where the shared pair was originally a lone pair of electrons on one of the bonded atoms.

example: Ammonium ion, NH₄⁺, a very common compound in AS/ A-level Chemistry

• From the reaction of ammonia, NH₃ and a hydrogen H+, the ammonia molecule donates its lone pair of electrons to the H⁺ ion, forming a dative covalent bond in NH₄⁺. 
• The dative colvenlend bond is shown by 
• an arrow pointing from N ➔ H as to show that the nitrogen atom provides both electrons to the covalent bond. In the NH₄⁺ion, all 4 bonds are equivalent and you can’t tell which is a dative covalent bond. The arrowhead ➔ for the dative covalent bond just helps with the accounting for all electrons.

Average Bond Enthalpy: a measurement of covalent bond strength

• Average bond enthapy serves as a measurement of covalent bond strength
• The larger the value of the average bond enthalpy, the stronger the covalent bond

This is the end of the topic!

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