Chemical Bonding and Electronegativity

Overview

Understanding chemical bonding and electronegativity is essential in chemistry. These concepts describe how atoms connect to form molecules, affecting the properties of these compounds.

• Chemical Bonding: The process through which atoms combine to create compounds and elements.
• Electronegativity: A measure of an atom's ability to attract shared electrons, vital for understanding chemical reactivity and bond types.

Types of Chemical Bonds

• Ionic Bonds: Involves the transfer of electrons between metals and non-metals.
• Covalent Bonds: Entails the sharing of electrons between non-metal atoms.
• Metallic Bonds: Characterised by a "sea" of free electrons among metal atoms.

Properties and Comparisons

Bond Type	Electrical Conductivity	Strength	Bond Formation
Ionic	High (when molten)	Strong	Electrostatic force
Covalent	Low	Variable	Electron sharing
Metallic	High	Strong	Electron sea

Examples

Ionic Bonds: Sodium chloride (NaCl) exemplifies classic ionic behaviour, exhibiting high melting points and conductivity when molten.

Covalent Bonds: Water (H₂O) serves as an example, with lower melting points due to the sharing of electrons.

Metallic Bonds: Aluminium (Al) is a good conductor and is malleable, attributed to its free electrons.

Electronegativity

Definition and Trends

• Measurement: Employs the Pauling scale, an effective method for comparing electronegativity.
• Trends:
  • Increases across a period (e.g., fluorine possesses high electronegativity).
  • Decreases down a group (due to a larger atomic radius).

Bond Polarity

• Determined by the electronegativity difference between atoms.
• Leads to dipoles or partial positive/negative charges.

Example: In polar HCl, the hydrogen region is positive, while the chlorine region is negative.

Determining Bond Type

• Calculate electronegativity differences:
  • 0.0 to 0.4: Non-polar covalent
  • 0.5 to 1.7: Polar covalent
  • Greater than 1.7: Ionic

Practical Examples

• HCl: Polar covalent (E.N. difference approximately 0.96)
• NaCl: Ionic (E.N. difference approximately 2.23)

Interactive Practice Problems

Problem Set 1

Element Pair	Electronegativity 1	Electronegativity 2	Difference	Bond Type
H - Cl	2.1	3.0	0.9	Polar covalent
Na - Cl	0.9	3.0	2.1	Ionic

Problem Set 2

Exercise 1: Draw Lewis structures for H₂O, CH₄, NH₃

Solution:

• H₂O: Oxygen atom has 6 valence electrons, and each hydrogen contributes 1. The oxygen forms single bonds with each hydrogen and has two lone pairs.
• CH₄: Carbon has 4 valence electrons and each hydrogen contributes 1. Carbon forms four single bonds with hydrogen atoms.
• NH₃: Nitrogen has 5 valence electrons and each hydrogen contributes 1. Nitrogen forms three single bonds with hydrogen atoms and has one lone pair.

Exercise 2: Predict properties based on bond types.

Solution:

• Ionic compounds (like NaCl) typically have high melting points, are brittle, and dissolve well in water.
• Polar covalent compounds (like H₂O) often have moderate melting points and good water solubility.
• Non-polar covalent compounds (like CH₄) typically have low melting points and poor water solubility.

Conclusion

Electronegativity and chemical bonding are crucial for understanding molecule formation and properties. Accurate calculations of electronegativity help predict bond types and molecular behaviour, essential for success in chemistry.