Chapter 2: The Chemistry of Life – Study Guide

Chapter 2: The Chemistry of Life

Introduction

This chapter explores the fundamental chemical principles underlying biological systems. It covers the structure of atoms, the nature of chemical bonds, the properties of water, and the basics of acids, bases, and pH—all essential for understanding life at the molecular level.

Atoms and Subatomic Particles

Definition and Structure

• Atom: The smallest unit of matter that retains the properties of an element.

• Subatomic particles:

  • Proton: +1 charge, mass = 1 atomic mass unit (amu), located in the nucleus.

  • Neutron: No charge, mass = 1 amu, located in the nucleus.

  • Electron: -1 charge, negligible mass, orbits the nucleus in electron shells.

• Atomic number: Number of protons in an atom.

• Mass number: Number of protons + neutrons.

• Isotopes: Atoms of the same element with different numbers of neutrons.

  • Stable isotopes: Do not change over time.

  • Radioisotopes: Unstable; emit radiation.

Key facts:

• 92 elements occur naturally.

• Atoms = elements = matter.

• Atomic structure determines properties.

• Oxygen (65%), carbon (18.5%), and hydrogen (9.5%) make up 96.3% of body mass.

• Six essential elements: C, H, O, N, P, S.

The Periodic Table of Elements

Organization and Use

• Elements are arranged by increasing atomic number.

• Groups (columns) share similar chemical properties.

• Periods (rows) indicate increasing energy levels (shells).

Electron Shells and Reactivity

Electron Arrangement

• First shell: Maximum 2 electrons.

• Second and subsequent shells: Maximum 8 electrons each.

• Valence shell: The outermost electron shell; determines chemical reactivity.

• Atoms with full valence shells are stable (nonreactive); those with unfilled shells are reactive.

Molecules and Compounds

Definitions

• Element: Pure substance containing only one kind of atom.

• Molecule: Electrically neutral group of two or more atoms held together by chemical bonds.

• Compound: Molecule made up of two or more different elements (e.g., H2O, NaCl).

• Molecular weight: The sum of the atomic weights of all atoms in a molecule.

Chemical Bonds

Types of Bonds

• Covalent bonds: Atoms share electrons.

  • Nonpolar covalent: Equal sharing (e.g., O2, H2).

  • Polar covalent: Unequal sharing (e.g., H2O).

• Ionic bonds: Attraction between oppositely charged ions (cations and anions).

  • Cation: Positively charged ion (loses electrons).

  • Anion: Negatively charged ion (gains electrons).

• Hydrogen bonds: Weak attraction between a hydrogen atom (partially positive) and an electronegative atom (partially negative), often between water molecules or in DNA/proteins.

• Electronegativity: The tendency of an atom to attract electrons; increases up and to the right on the periodic table.

Chemical Reactions

Making and Breaking Bonds

• Reactants → Products

• Chemical reactions do not create or destroy matter (Law of Conservation of Matter).

Water and Life

Properties of Water

• Polarity: Water is a polar molecule, making it an excellent solvent for ionic and polar substances.

  • Hydrophilic: Polar; interacts with water.

  • Hydrophobic: Nonpolar; avoids water.

• Hydrogen bonding: Critical for the structure of DNA, proteins, and the unique properties of water.

• High specific heat: Water resists temperature change.

  • 1 calorie (cal) raises 1 gram of water by 1°C.

• High heat of vaporization: Water requires significant energy to change from liquid to gas, aiding in cooling (e.g., sweating).

• Ice floats: Solid water is less dense than liquid water, allowing ice to insulate aquatic environments.

• Cohesion: Water molecules stick together via hydrogen bonds (important for plant water transport).

• Surface tension: Water resists external force, allowing small objects to rest on its surface.

• Amphoteric property: Water can act as both a weak acid and a weak base, ionizing into H+ and OH-.

Acids, Bases, and pH

Definitions and Properties

• Acid: Substance that releases H+ ions in solution.

• Base: Substance that accepts H+ ions or releases OH- ions in solution.

• pH scale: Measures the concentration of H+ ions.

  • pH 7 = neutral (pure water)

  • pH < 7 = acidic

  • pH > 7 = basic (alkaline)

  • Most biological fluids: pH 6–8

• Buffer: A weak acid and its corresponding base that stabilizes pH by absorbing or releasing H+ or OH- ions.

Key Equations

• Mass number: $\text{Mass number} = \text{protons} + \text{neutrons}$

• Number of neutrons: $\text{Neutrons} = \text{mass number} - \text{atomic number}$

• Ionization of water: $\mathrm{H}_2\mathrm{O} \rightleftharpoons \mathrm{H}^+ + \mathrm{OH}^-$

• Acid dissociation (example): $\mathrm{HCl} \rightarrow \mathrm{H}^+ + \mathrm{Cl}^-$

• Base dissociation (example): $\mathrm{NaOH} \rightarrow \mathrm{Na}^+ + \mathrm{OH}^-$

Summary Table: Key Properties and Definitions

Term	Definition	Example/Application
Atom	Smallest unit of matter retaining element properties	Hydrogen atom (H)
Isotope	Atoms of same element, different neutrons	Carbon-12, Carbon-14
Covalent bond	Electron sharing between atoms	H2O, O2
Ionic bond	Attraction between oppositely charged ions	NaCl
Hydrogen bond	Weak attraction between polar molecules	Between water molecules, DNA strands
pH	Measure of H+ concentration	pH 7 = neutral
Buffer	Stabilizes pH by absorbing/releasing H+	Bicarbonate in blood

Additional info:

• Electronegativity differences determine bond polarity.

• Water’s high specific heat and heat of vaporization are crucial for temperature regulation in living organisms.

• Surface tension and cohesion are essential for processes like transpiration in plants.