Lecture 4

Chemical Bonds and Electronegativity

Types of Chemical Bonds

Chemical bonds are the forces that hold atoms together in molecules and compounds. The type of bond formed depends on the arrangement of electrons and the electronegativity of the atoms involved.

• Covalent Bonds: Atoms share pairs of electrons. Can be polar or nonpolar depending on electronegativity differences.

• Ionic Bonds: Electrons are transferred from one atom to another, resulting in oppositely charged ions that attract each other.

• Hydrogen Bonds: Weak attractions between a hydrogen atom covalently bonded to an electronegative atom (like oxygen or nitrogen) and another electronegative atom.

Example: Sodium chloride (NaCl) forms via ionic bonding between Na+ and Cl- ions.

Valence Electrons and Bond Formation

The number of bonds an atom can form is determined by its valence electrons—the electrons in its outermost shell.

• Valence Electrons: Directly involved in chemical reactions and bond formation.

• Bonding Capacity: Atoms tend to form enough bonds to fill their valence shell (octet rule).

Example: Oxygen has 6 valence electrons and typically forms 2 covalent bonds.

Electronegativity and Bond Polarity

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

• Polar Covalent Bonds: Formed when atoms have different electronegativities, resulting in unequal sharing of electrons and partial charges.

• Nonpolar Covalent Bonds: Formed when atoms have similar electronegativities, resulting in equal sharing of electrons.

Example: In water (H2O), oxygen is more electronegative than hydrogen, creating a polar covalent bond.

Comparison Table: Polar vs. Nonpolar Covalent Bonds

Bond Type	Electronegativity Difference	Electron Sharing	Example
Polar Covalent	High	Unequal	H2O
Nonpolar Covalent	Low/None	Equal	O2

Properties of Water

Structure and Polarity of Water Molecule

Water is a polar molecule due to the difference in electronegativity between oxygen and hydrogen. The molecule has a bent shape, with partial negative charge near oxygen and partial positive charges near hydrogens.

• Partial Charges: Oxygen (δ-) and Hydrogen (δ+)

• Hydrogen Bonding: Water molecules form hydrogen bonds with each other due to polarity.

Example: Water's polarity allows it to dissolve many substances, making it a universal solvent.

Hydrogen Bonding and Water's Unique Properties

Hydrogen bonds between water molecules give rise to several important properties essential for life.

• Cohesion: Water molecules stick to each other, contributing to surface tension.

• Adhesion: Water molecules stick to other surfaces.

• Surface Tension: The cohesive forces at the surface of water allow small objects to "float" or insects to walk on water.

• High Specific Heat: Water can absorb or release large amounts of heat with little temperature change.

• Ice Floats: Solid water (ice) is less dense than liquid water due to hydrogen bonding, allowing ice to float.

• Solvent Properties: Water dissolves polar and ionic substances effectively.

Example: Water's high specific heat helps moderate Earth's climate.

Comparison Table: Cohesion vs. Adhesion

Property	Description	Example
Cohesion	Attraction between water molecules	Surface tension
Adhesion	Attraction between water and other substances	Water climbing plant stems

Acids, Bases, and pH

pH Scale and Water Dissociation

The pH scale measures the concentration of hydrogen ions (H+) in a solution. Water can dissociate into H+ and OH- ions.

• pH Definition: $pH = -\log[H^+]$

• Neutral pH: Pure water has a pH of 7, with equal concentrations of H+ and OH-.

• Acid: Substance that increases H+ concentration (pH < 7).

• Base: Substance that decreases H+ concentration (pH > 7).

Example: Gastric juice is acidic (pH ~2), while bleach is basic (pH ~13).

Buffers and Biological Importance

Buffers are substances that minimize changes in pH by absorbing or releasing H+ ions. They are crucial for maintaining stable pH in biological systems.

• Function: Prevent harmful pH fluctuations in cells and organisms.

• Example: Blood contains bicarbonate buffer to maintain pH near 7.4.

pH Scale Table

pH Value	Type	Example
2	Acid	Gastric juice
7	Neutral	Pure water
13	Base	Bleach

Key Terms and Definitions

• Solute: Substance dissolved in a solution.

• Solvent: Substance that dissolves the solute (water is the universal solvent).

• Hydrophilic: "Water-loving"; substances that dissolve well in water.

• Hydrophobic: "Water-fearing"; substances that do not dissolve in water.

• Reactants: Starting materials in a chemical reaction.

• Products: Substances formed from a chemical reaction.

Summary

Understanding chemical bonds, water's unique properties, and the pH scale is fundamental to biology. These concepts explain how molecules interact, how life is sustained, and how organisms maintain homeostasis in changing environments.