AS-Level Chemistry Chapter 4 Chemical Bonding Types of Bonds

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Ionic Bonding

Formation of ionic bond

1) Ionic bond is the electrostatic force of attraction between oppositely-charged ions formed by the complete transfer of electrons from an atom to another atom.

2) Ionic bond is also called electrovalent bond.

3) i. An atom(usually a metal) that loses electron(s) will form a positive ion(cation).

ii. The electron is then transferred to another atom.

iii. The atom that gains the electron(usually a non-metal) will form a negative ion(anion).

iv. The cations and anions are then attracted by strong electrostatic force of attraction. The force of attraction constitutes the ionic bond.

4) The force of attraction between cation and anion is very strong, therefore ionic bond is a very strong bond.

5) Ionic bonds are non-directional, each cation will attract any neighboring anion and vice versa to form a huge ionic lattice.

6) The compound formed as a result of ionic bond is called ionic compound. An example is sodium chloride, NaCl.

Dot-and-cross diagram

1) A dot-and-cross diagram shows:

i. the outer electron shell only.

ii. that the charge of the ion is spread evenly using a square bracket.

iii. the charge of each ion.

2) It is also called the Lewis diagram.

Strength of ionic bonds

1) The strength of ionic bond is a measure of the electrostatic force of attraction between the ions.

2) The force of attraction between the oppositely-charged ions is proportional to the charge on the ions and inversely proportional to the square of distance between the ions.

3) The strength of ionic bond is manifested in the melting point of the ionic compound.

4) i. For instance, the melting point of NaCl is higher than NaBr.

ii. This is because the size of Br⁻ ion is larger than Cl⁻ ion. Therefore the distance between Br⁻ and Na⁺ is larger than that of between Cl⁻ and Na⁺.

iii. As a result, the electrostatic force of attraction between Na⁺ and Cl⁻ is stronger than that of between Na⁺ and Br⁻ ion.

5) i. The melting point of NaCl is lower than MgCl2.

ii. This is because Mg²⁺ ion has a higher charge than Na⁺ ion. Besides that, the size of Mg²⁺ ion is smaller than Na⁺ ion.

iii. The above two factors causes the electrostatic force of attraction between Mg²⁺ and Cl⁻ to be stronger than that of between Na⁺ and Cl⁻.

Covalent Bonding

Formation of covalent bond

1) Covalent bond is the electrostatic force of attraction that two neighboring nuclei have for a localized pair of electrons shared between them.

2) Covalent bond is formed without transferring electrons, instead, the atoms share their valence electron(s) to achieve duplet/octet electronic configuration.

3) The shared pair of electrons constitutes the covalent bond.

Single bond

1) Single bond is formed when one pair of electrons is shared between two atoms.

2) Examples of compounds with single bonds:

Double bond

1) Double bond is formed when two pairs of electrons are shared between two atoms.

2) Examples of compounds with double bonds:

Triple bond

1) Triple bond is formed when three pairs of electrons are shared between two atoms.

2) Examples of compounds with triple bonds:

Lone pair and bond pair of electrons

1) The pair of electrons used in covalent bonding is called the bond pair while the pair of electrons not used in covalent bonding is called the lone pair.

Octet-deficient and expanded octet species

1) In general, atoms tend to share their electrons to a achieve a duplet/octet electronic configuration - the octet rule.

2) i. In octet-deficient species, the central atom has less than eight electrons.

ii. Some examples are boron trifluoride, BF3 and nitrogen monoxide, NO.

3) i. In expanded octet species, the central atom has more than eight electrons.

ii. An example is phosphorus(V) chloride, PCl5.

iii. This is possible only for Period 3 elements and beyond, this is because starting from Period 3, the atoms have empty d orbitals in the third energy level to accommodate more than eight electrons.

Co-ordinate bond (dative covalent bond)

1) A co-ordinate bond is formed when one atom provides both the electrons needed for a covalent bond.

2) Conditions of forming a co-ordinate bond:

1) one atom has a lone pair of electrons.

2) another atom has an unfilled orbital to accept the lone pair, in other words, an electron-deficient species.

3) Once the bond is formed, it is identical to the other covalent bonds. It does not matter where the electrons come from.

4) In a displayed formula, a co-ordinate bond is represented by an arrow, the head of the arrow points away from the lone pair which forms the bond.

5) An example is the reaction between ammonia and hydrogen chloride. In this reaction, ammonium ion is formed by the transfer of hydrogen ion(an octet deficient species) from hydrogen chloride to the lone pair of electrons in the ammonia molecule.

6) Another example is aluminium chloride. At high temperature, it exists as AlCl3. At low temperature(around 180-190°C), it exists as Al2Cl6, a dimer(two molecules joined together). This is possible because lone pairs of electrons from the chlorine atom form co-ordinate bonds with the aluminium atom.

Identify the central atom and terminal atom(s).

For example, in ammonia, the nitrogen is the central atom while the hydrogens are the terminal atoms.

1) During the sharing of electrons, the terminal atoms must attain octet configuration(or duplet for hydrogen) but not necessarily for the central atom.

2) i. If the central atom is from Period 2 of the Periodic Table, the total number of electrons surrounding it cannot exceed eight(but can less than eight).

ii. If the central atom is from Period 3 and beyond, the total number of electrons surrounding it can exceed eight

3) i. For polyatomic anions, the negative charge will be distributed among the most electronegative atom(s). This is to decrease the charge density on a particular atom and to stabilize the ion.

ii. For polyatomic cation, the positive charge will be distributed among the less

electronegative atom(s). The reason is same as above.

4) If the terminal atom already has octet configuration(for example, Cl⁻), it will contribute two electrons to the central atom to form a co-ordinate bond.

Effect of lone pair on bond angle

1) For methane, ammonia and water, the electron pair geometries are tetrahedral. However, the molecular geometries are different.

2) In methane, all the bonds are identical, repulsion between the bonds is the same. Thus, methane has a perfect tetrahedral structure with bond angle 109.5°.

3) In ammonia, the repulsion between the lone pair and the bond pairs is stronger than in methane. This forces the bond angle to decrease slightly to 107°.

4) In water, there are two lone pairs and thus the repulsion is the greatest, the two bond pairs are pushed closer to one another and the bond angle is reduced to 104.5°.

Effect of electronegativity on bond angle

1) Water and hydrogen sulfide have the same general shape with the same number of bond pairs and lone pairs. However, their bond angles are different.

2) This is because oxygen has a higher electronegativity than sulphur. The bond pairs of electrons are closer to the oxygen atom compared to the sulfur atom.

3) This results in greater repulsion in the O-H bonds than in the S-H bonds. Therefore, the bond angle increases from 92.5° to 104.5°.

Sigma(σ) bond and pi(π) bond

1) A sigma bond is formed by orbitals from two atoms overlapping end-to-end.

2) In a sigma bond, the electron density is concentrated between the two nuclei.

3) A pi bond is formed by the p orbitals from two atoms overlapping sideways.

4) In a pi bond, there are two regions of high electron density alongside the nuclei.

5) A pi bond is weaker than a sigma bond because the overlapping of charge clouds is less than in a sigma bond.

6) In covalent molecules, single bonds are sigma bonds(σ), a double bond consists of one sigma bond and one pi bond(1σ, 1π), and a triple bond consists of one sigma bond and two pi bonds(1σ, 2π).