Chapter 2: The Chemical Context of Life – Study Notes

Chapter 2: The Chemical Context of Life

Introduction

This chapter explores the fundamental chemical principles that underlie biological processes. Understanding the structure of atoms, the nature of chemical bonds, and the properties of water is essential for studying life at the molecular level.

Atoms and Elements

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

• Element: A substance that cannot be broken down into other substances by chemical means. Each element is defined by its number of protons.

• Atomic Number: The number of protons in an atom's nucleus, which determines the element.

• Atomic Mass: The total mass of an atom, approximately equal to the sum of protons and neutrons.

• Isotopes: Atoms of the same element with different numbers of neutrons. Isotopes can have different physical properties and some may be radioactive.

• Example: Carbon-12 and Carbon-14 are isotopes of carbon.

Subatomic Particles

• Proton: Positively charged particle in the nucleus.

• Neutron: Neutral particle in the nucleus.

• Electron: Negatively charged particle orbiting the nucleus.

• Electrical Charge: Protons (+1), electrons (–1), neutrons (0).

• Contribution to Mass: Protons and neutrons contribute most of the atom’s mass; electrons have negligible mass.

Electron Arrangement and Behavior

• Electron Shells (Orbitals): Regions around the nucleus where electrons are likely to be found. Shells are filled in order of increasing energy.

• Valence Electrons: Electrons in the outermost shell, important for chemical bonding.

• Behavior: The arrangement of electrons determines how atoms interact and bond with each other.

• Energy Absorption and Release: Electrons can absorb energy and move to higher shells, or release energy and fall to lower shells.

Elements Essential to Life

• Major Elements: Carbon (C), Hydrogen (H), Oxygen (O), Nitrogen (N) – often abbreviated as CHON – are the most abundant in living organisms.

• Other Important Elements: Phosphorus (P), Sulfur (S), Calcium (Ca), Potassium (K), Magnesium (Mg), Iron (Fe), Sodium (Na), Chlorine (Cl).

Chemical Bonds

• Covalent Bond: A chemical bond formed when two atoms share one or more pairs of electrons.

• Polar Covalent Bond: Electrons are shared unequally, resulting in partial charges (δ+ and δ–) on the atoms.

• Nonpolar Covalent Bond: Electrons are shared equally between atoms.

• Ionic Bond: Formed when one atom donates an electron to another, resulting in oppositely charged ions that attract each other.

• Hydrogen Bond: A weak bond between a hydrogen atom (already covalently bonded to an electronegative atom) and another electronegative atom.

• Double Covalent Bond: Two pairs of electrons are shared between atoms (e.g., O2 molecule).

• Partial Charges: Occur in polar covalent bonds due to unequal sharing of electrons.

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

Ions and Molecules

• Ion: An atom or molecule with a net electric charge due to the loss or gain of electrons.

• Cation: Positively charged ion (loss of electrons).

• Anion: Negatively charged ion (gain of electrons).

• Molecule: Two or more atoms held together by covalent bonds.

• Molecular Formula: Shows the types and numbers of atoms in a molecule (e.g., H2O).

• Macromolecule: Large, complex molecules such as proteins, nucleic acids, carbohydrates, and lipids.

Properties of Water

• Cohesion: Attraction between molecules of the same substance (e.g., water molecules stick together).

• Adhesion: Attraction between molecules of different substances (e.g., water molecules stick to glass).

• Surface Tension: The measure of how difficult it is to stretch or break the surface of a liquid.

• High Heat of Vaporization: Water requires a lot of energy to change from liquid to gas.

• Density of Ice vs. Water: Ice is less dense than liquid water, so it floats.

• Solubility: Water is a universal solvent, dissolving many polar and ionic substances.

• Hydrophilic: Substances that dissolve in water (water-loving).

• Hydrophobic: Substances that do not dissolve in water (water-fearing).

• Example: Salt (NaCl) dissolves in water; oil does not.

Solutions and pH

• Solute: The substance that is dissolved.

• Solvent: The substance that does the dissolving (in biology, usually water).

• Solution: A homogeneous mixture of solute and solvent.

• Electrolyte: A substance that dissociates into ions in solution and conducts electricity.

• pH: A measure of hydrogen ion concentration; $\text{pH} = -\log_{10}[\text{H}^+]$

• Acid: A substance that increases the hydrogen ion concentration in a solution.

• Base: A substance that reduces the hydrogen ion concentration in a solution.

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

• Example: Pure water has a pH of 7; stomach acid has a pH of about 2.

Table: Comparison of Bond Types

Bond Type	How Formed	Relative Strength	Example
Covalent	Sharing of electron pairs	Strong	H2O, O2
Ionic	Transfer of electrons	Strong (in dry conditions)	NaCl
Hydrogen	Attraction between H (in polar bond) and electronegative atom	Weak	Between water molecules

Vocabulary

• Acid

• Adhesion

• Anion

• Atomic number

• Base

• Bond, Covalent

• Bond, Hydrogen

• Cation

• Charge, Negative

• Charge, Positive

• Cohesion

• Electron

• Electrolyte

• Element

• Heat of vaporization

• Hydrophilic

• Hydrophobic

• Ion

• Isotope

• Macromolecule

• Molecule

• Molecular formula

• Nucleus, Atomic

• Orbital shell

• Proton

• pH

• Polar (vs. non-polar) covalent bond

• Salt

• Solute

• Solution

• Solvent

• Valence electron

• Valence shell

Additional info:

• Some explanations and examples have been expanded for clarity and completeness.

• Table content inferred and organized for comparison of bond types.