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 structure of atoms, elements, and compounds is essential for grasping how life functions 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.

• 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.

Example: Water (H2O) is a compound made of hydrogen and oxygen in a 2:1 ratio.

Essential and Trace Elements in Life

• Essential Elements (make up ~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 (<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 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.

Example: Carbon-12 has 6 protons, 6 neutrons, and 6 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.

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

Formula: $\text{Atomic Mass} = \text{Number of Protons} + \text{Number of Neutrons}$

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.

Example: Carbon-14 is a radioactive isotope of carbon.

Formula for Remaining Mass: $\text{Remaining Mass} = \text{Initial Mass} \times \left(\frac{1}{2}\right)^{\text{Number of Half-Lives}}$

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

Chemical Bonds

• Covalent Bonds: Atoms share pairs of valence electrons.

  • Single Covalent Bond: Sharing one pair of electrons (e.g., H—H).

  • Double Covalent Bond: Sharing two pairs of electrons (e.g., O=O).

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

  • Nonpolar Covalent Bond: Electrons shared equally.

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

• Ionic Bonds: Atoms transfer electrons, forming ions (cations and anions) that attract each other.

• Hydrogen Bonds: Weak bonds formed when a hydrogen atom covalently bonded to one electronegative atom is attracted to another electronegative atom (commonly oxygen or nitrogen).

• Van der Waals Interactions: Weak attractions due to transient local partial charges.

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.

Property	Description
Crystal Formation	Salts form lattice structures.
Stability	Stable when dry, dissociate in water.
Ratio	Formula shows element ratio, not discrete molecules.

Hybridization

• Hybridization: The process by which atomic orbitals mix to form new, hybrid orbitals suitable for bonding.

• Determines molecular shape and function.

Example: Carbon’s sp3 hybridization in methane (CH4).

Concept 2.4: Chemical Reactions Make and Break Chemical Bonds

Chemical Reactions

• Reactants: Starting molecules in a chemical reaction.

• Products: Resulting molecules after the reaction.

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

• Chemical reactions are reversible, indicated by double arrows ($\rightleftharpoons$).

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

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

Additional info:

• Understanding electron configuration and orbitals is crucial for predicting chemical behavior and bonding.

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