Chemical Bonding

Atoms rarely exist alone in nature. They bond with other atoms to become more stable -- usually by achieving a full outer shell of electrons (like noble gases). The way atoms bond determines the properties of every substance around us.

Why Do Atoms Bond?

Atoms are most stable when their outermost shell is full. They achieve this by:

• Losing electrons (metals)
• Gaining electrons (non-metals)
• Sharing electrons (between non-metals)

This drive for stability leads to three main types of bonds.

Ionic Bonding

Ionic bonds form when a metal transfers electrons to a non-metal. The metal becomes a positive ion (cation), and the non-metal becomes a negative ion (anion). The opposite charges attract.

Example: Sodium Chloride (NaCl)

• Na (2, 8, 1) loses 1 electron to become Na⁺ (2, 8)
• Cl (2, 8, 7) gains 1 electron to become Cl⁻ (2, 8, 8)
• The Na⁺ and Cl⁻ attract each other = ionic bond

Properties of ionic compounds: High melting points, conduct electricity when dissolved or melted, form crystals, usually soluble in water.

Covalent Bonding

Covalent bonds form when non-metals share electrons.

Example: Water (H₂O)

• Oxygen needs 2 more electrons; each hydrogen needs 1 more
• Oxygen shares one electron with each hydrogen atom
• Result: 2 covalent bonds

Types of covalent bonds:

• Single bond: 1 shared pair (H-H)
• Double bond: 2 shared pairs (O=O)
• Triple bond: 3 shared pairs (N≡N)

Properties of covalent compounds: Lower melting points, usually do not conduct electricity, can be gases/liquids/soft solids.

Metallic Bonding

In metals, atoms release their valence electrons into a shared "sea of electrons." The positive metal ions are held together by this electron sea.

Properties: Good conductors of heat and electricity, malleable, ductile, shiny.

Lewis Structures (Dot Diagrams)

Lewis structures show valence electrons as dots around the element symbol.

Steps to draw Lewis structures:

1. Count total valence electrons of all atoms
2. Place the least electronegative atom in the center
3. Connect atoms with single bonds (each bond = 2 electrons)
4. Distribute remaining electrons to complete octets
5. Form double/triple bonds if needed

Electronegativity

Electronegativity is the ability of an atom to attract shared electrons in a bond.

• High electronegativity difference (>1.7): ionic bond
• Moderate difference (0.4-1.7): polar covalent bond
• Low difference (<0.4): non-polar covalent bond

Did You Know? The salt (NaCl) used in Nepali cooking is an ionic compound. The strong ionic bonds give it a high melting point (801°C), which is why salt does not melt in your kitchen!

Key Takeaways

• Ionic bonds: electron transfer between metals and non-metals
• Covalent bonds: electron sharing between non-metals
• Metallic bonds: sea of shared electrons among metal atoms
• Electronegativity difference determines bond type
• Lewis structures help visualize how atoms share or transfer electrons