Chapter 2: The Chemical Context of Life – General Biology Study Notes

Chapter 2: The Chemical Context of Life

Introduction

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

Concept 2.1: Matter Consists of Chemical Elements in Pure Form and in Combinations Called Compounds

Definition of Matter, Elements, and Compounds

• Matter: Anything that takes up space and has mass. All organisms are composed of matter.

• Element: A substance that cannot be broken down to other substances by chemical reactions. Each element is made of unique atoms.

• Compound: A substance consisting of two or more elements in a fixed ratio. Compounds have properties different from their constituent elements.

Example: Water (H2O) is a compound made from hydrogen and oxygen.

Essential and Trace Elements in Life

• Essential Elements (about 96% of living matter): Carbon (C), Oxygen (O), Hydrogen (H), Nitrogen (N)

• Other Essential Elements (~4%): Calcium (Ca), Phosphorus (P), Potassium (K), Sulfur (S), Sodium (Na), Chlorine (Cl), Magnesium (Mg)

• Trace Elements: Required in minute quantities (less than 0.01%), e.g., iron (Fe), iodine (I)

Element	Symbol	Percentage of Body Mass
Oxygen	O	65.0%
Carbon	C	18.5%
Hydrogen	H	9.5%
Nitrogen	N	3.3%
Calcium	Ca	1.5%
Phosphorus	P	1.0%
Potassium	K	0.4%
Sulfur	S	0.3%
Sodium	Na	0.2%
Chlorine	Cl	0.2%
Magnesium	Mg	0.1%

Concept 2.2: An Element’s Properties Depend on the Structure of Its Atoms

Atomic Structure

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

• Subatomic Particles:

  • Neutrons: No electrical charge; contribute to atomic mass and isotopes.

  • Protons: Positive charge; determine the element’s identity.

  • Electrons: Negative charge; involved in chemical bonding.

• Atoms are electrically neutral overall because they have equal numbers of protons and electrons.

Atomic Number and Atomic Mass

• Atomic Number: Number of protons in the nucleus; unique to each element.

• Atomic Mass (Mass Number): Sum of protons and neutrons in the nucleus; measured in daltons.

• Electrons are much lighter and do not contribute significantly to atomic mass.

Example: Carbon-12 has 6 protons and 6 neutrons ($6+6=12$).

Isotopes and Radioactivity

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

• Radioactive Isotopes: Unstable isotopes that decay spontaneously, emitting particles and energy.

• Half-life: The time required for half of the atoms in a radioactive sample to decay.

Applications: Radioactive isotopes are used in medical diagnostics and radiometric dating.

Formula for remaining mass after n half-lives:

$\text{Remaining mass} = \text{Initial mass} \times \left(\frac{1}{2}\right)^n$

Concept 2.3: The Formation and Function of Molecules and Ionic Compounds Depend on Chemical Bonding Between Atoms

Chemical Bonds

• Atoms with incomplete valence shells can share or transfer electrons, forming chemical bonds.

• Main Types of Chemical Bonds:

  • Covalent Bonds (strong)

  • Ionic Bonds (strong)

  • Hydrogen Bonds (weak)

  • Van der Waals Interactions (weak)

Covalent Bonds

• Involve sharing of valence electrons between atoms.

• Single Covalent Bond: Sharing one pair of electrons.

• Double Covalent Bond: Sharing two pairs of electrons.

• Electronegativity: An atom’s attraction for electrons in a covalent bond.

• Nonpolar Covalent Bond: Electrons shared equally.

• Polar Covalent Bond: Electrons shared unequally, leading to partial charges.

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

Ionic Bonds

• Formed when atoms transfer electrons, creating ions.

• Cation: Positively charged ion.

• Anion: Negatively charged ion.

• Oppositely charged ions attract, forming ionic bonds.

Example: Sodium chloride (NaCl) is formed from Na+ and Cl-.

Properties of Ionic Compounds (Salts)

• Often form crystals in nature.

• Not considered molecules; formula indicates element ratio in the crystal.

• Stable when dry, but dissociate easily in water.

Hydrogen Bonds

• Form when a hydrogen atom covalently bonded to one electronegative atom is attracted to another electronegative atom.

• Common in water and biological molecules (partners usually oxygen or nitrogen).

Van der Waals Interactions

• Occur when electrons are distributed unevenly, creating temporary charges.

• Weak individually, but can be collectively strong (e.g., gecko’s toe hairs adhering to surfaces).

Hybridization

• Occurs when atomic orbitals mix to form new, hybrid orbitals during bonding.

• Determines molecular shape and function.

Example: Carbon’s s and p orbitals hybridize to form tetrahedral structures in methane (CH4).

Concept 2.4: Chemical Reactions Make and Break Chemical Bonds

Chemical Reactions

• Involve the making and breaking of chemical bonds.

• Reactants: Starting molecules.

• Products: Resulting molecules.

• Reactions are reversible; products can become reactants in the reverse reaction.

• Chemical Equilibrium: Point at which forward and reverse reactions occur at the same rate; concentrations of reactants and products remain constant.

Example: $2H_2 + O_2 \rightarrow 2H_2O$

Equilibrium Example: $3H_2 + N_2 \rightleftharpoons 2NH_3$

Additional info:

• Atomic orbitals are three-dimensional spaces where electrons are found 90% of the time.

• Electron configuration determines chemical behavior and bonding capacity.

• Practice problems may involve calculating isotopes, half-lives, and electron configurations.