BIO 110- Chapter 2

General Biology - Bio 110

Chapter 2: Chemistry of Life

Contents Overview

  • Part I: Life on Earth: An Overview

  • Part II: Chemistry of Life

    • a. Basic Chemistry (Chapter 2)

    • b. Chemistry of Organic Molecules (Chapter 3)

  • Part III: The Cell

    • a. Cell Structure and Function (Chapter 4)

    • b. Membrane Structure and Function (Chapter 5)

Part II: Chemistry of Life

2.1 Chemical Elements

a. Elements

  • Matter exists in three states: solid, liquid, gas.

  • Composed of elements: substances that cannot be broken down by ordinary chemical means.

  • Total of 92 elements, with six (CHNOPS) essential to life, making up 95% of body weight:

    • Elements: Carbon (C), Hydrogen (H), Nitrogen (N), Oxygen (O), Phosphorus (P), Sulfur (S).

  • Additional important elements: Iron (Fe), Potassium (K), Magnesium (Mg), Calcium (Ca).

b. Atoms

  • Atomic theory: Atoms are the smallest part of an element exhibiting its properties.

  • Atomic symbols: Represented by one or two letters (e.g., H for hydrogen, Na for sodium).

  • Subatomic particles:

    • Protons (positively charged)

    • Neutrons (uncharged)

    • Electrons (negatively charged)

  • Protons and neutrons are in the nucleus; electrons orbit the nucleus.

c. Atomic Number and Mass Number

  • Atomic number: Number of protons in the nucleus of the atom.

  • Mass number: Sum of protons and neutrons in the nucleus.

  • Mass is used instead of weight, as mass remains constant regardless of gravity.

  • Notation: Atomic number is a subscript; mass number is a superscript (e.g., 4He).

d. The Periodic Table

  • Constructed based on chemical and physical properties.

  • Displays electrically neutral atoms, with atomic number indicating protons and electrons.

  • To find neutrons: Mass number - Atomic number.

  • Atoms are arranged into periods (horizontal) and groups (vertical columns) with similar properties.

  • Noble gases (Group VIII) are inert and non-reactive.

e. Isotopes

  • Isotopes: Variants of an element with the same number of protons but different neutrons.

  • Example: Carbon isotopes - Carbon-12 (6 neutrons), Carbon-13 (7 neutrons), Carbon-14 (8 neutrons).

  • Carbon-14 is radioactive and decays into Nitrogen-14.

2.2 Compounds and Molecules

  • Compounds: Formed when two or more elements bond together.

  • Molecules: Smallest part of a compound, e.g., Water (H2O).

a. Ionic Bonding

  • Ionic bond: Electrovalent bond formed between oppositely charged ions.

  • Example: Table salt (NaCl).

    • Sodium donates an electron; chlorine accepts one to achieve stable outer shells.

b. Covalent Bonding

  • Covalent bond: Atoms share electrons to achieve stability.

  • Hydrogen sharing example: H—H (structural) or H2 (molecular).

  • Double covalent bond: Two pairs of electrons shared (e.g., O=O).

c. Polar and Non-polar Covalent Bonds

  • Non-polar covalent bonds: Electrons shared equally.

  • Polar covalent bonds: Electrons shared unequally due to differing electronegativities (e.g., water).

2.3 Chemistry of Water

Water Structure and Properties

  • Water molecule: Comprises one oxygen atom and two hydrogen atoms, forming polar covalent bonds.

  • Hydrogen bonding arises from attraction between hydrogen and oxygen atoms of adjacent molecules.

Properties of Water

  1. High heat capacity: Moderates temperature changes and supports life.

  2. Universal solvent: Dissolves many substances due to its polarity.

  • Hydrophilic (water-attracting) vs. hydrophobic (water-repelling) properties.

  1. Cohesive and adhesive properties:

  • Cohesion: Water molecules cling to each other, enabling liquid state under normal conditions.

  • Adhesion: Water molecules cling to polar surfaces; important for nutrient and waste transport in organisms and plants.

  1. Capillary action: Movement of water in narrow spaces against gravity; essential for plant water transport.

2.4 Acids and Bases

a. Acidic Solutions

  • Acids release hydrogen ions (H+) in water, e.g., lemon juice, vinegar.

  • Strong acids (HCl) dissociate nearly completely in water.

b. Basic Solutions

  • Bases accept H+ or release hydroxide ions (OH-), e.g., ammonia.

  • Strong bases (NaOH) also dissociate in water.

c. pH Scale

  • Ranges from 0 (acidic) to 14 (basic), with 7 being neutral.

  • Each change of 1 pH unit represents a tenfold change in concentration of H+ or OH-.