Biology & Chemistry: Foundations for Life

Biology & Chemistry: Foundations for Life

Introduction

This chapter explores the intimate connection between chemistry and the evolution of life. Understanding the chemical basis of biological processes is essential for grasping how living systems function, adapt, and evolve.

Why Does Biology Matter?

Importance of Biology in Society

• We are living things: Biology helps us understand ourselves and other organisms.

• Medicine and health: Biological knowledge underpins medical advances and healthcare.

• Agriculture: Biology informs crop improvement, pest control, and sustainable food production.

• Ecosystem services: Biological systems provide essential services such as pollination, water purification, and nutrient cycling.

• Climate change: Biological processes influence and are influenced by global climate patterns.

Why Chemistry Matters in Biology

Role of Chemistry in Life and the Environment

• Biochemistry within an organism: All living things rely on chemical reactions for metabolism, growth, and reproduction.

• Biochemistry in the environment: Chemical cycles, such as the carbon cycle, regulate the flow of elements and energy through ecosystems.

Example: The global carbon cycle illustrates how carbon moves between the atmosphere, organisms, oceans, and rocks, affecting climate and life.

Basic Atomic Structure

Parts of the Atom

• Protons: Positively charged particles (+1 charge) located in the nucleus.

• Neutrons: Neutral particles (no charge) also found in the nucleus.

• Electrons: Negatively charged particles (-1 charge) that orbit the nucleus.

Key Point: Most of an atom's volume is empty space, with the nucleus occupying a tiny central region.

The Periodic Table

Elements and Their Properties

• Elements: Pure substances consisting of only one type of atom.

• The Periodic Table organizes elements by atomic number and reveals patterns in their chemical properties.

• Types of chemical bonds include polar covalent, nonpolar covalent, ionic, and hydrogen bonds.

Atomic Number and Mass Number

• Atomic number (Z): The number of protons in the nucleus; defines the element.

• Mass number (A): The total number of protons and neutrons in the nucleus.

Formula: $\text{Mass number} = \text{Number of protons} + \text{Number of neutrons}$ $\text{Atomic number} = \text{Number of protons}$

Example: Carbon has an atomic number of 6 (6 protons). Its most common isotope, carbon-12, has 6 neutrons (mass number 12).

Isotopes

Definition and Importance

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

• Isotopes have different mass numbers but the same chemical properties.

• The atomic weight of an element is the weighted average of all naturally occurring isotopes.

Example: Carbon-12, Carbon-13, and Carbon-14 are isotopes of carbon with 6, 7, and 8 neutrons, respectively.

Electron Arrangement

Orbitals and Electron Shells

• Electrons occupy orbitals, specific regions around the nucleus.

• Each orbital can hold up to two electrons.

• Orbitals are grouped into electron shells, numbered 1, 2, 3, etc., with lower numbers closer to the nucleus.

• The valence shell is the outermost electron shell; the number of unpaired electrons in this shell determines an atom's chemical reactivity.

Key Point: Atoms are most stable when their valence shells are full.

Chemical Bonds

Covalent Bonds

• Covalent bonds: Formed when two atoms share pairs of valence electrons, resulting in a full outer shell for each atom.

• Nonpolar covalent bonds: Electrons are shared equally between atoms (e.g., H2 molecule).

• Polar covalent bonds: Electrons are shared unequally, leading to partial charges (e.g., H2O molecule).

Example: In water (H2O), oxygen is more electronegative than hydrogen, so electrons are pulled closer to oxygen, creating a partial negative charge on O and partial positive charges on H.

Electronegativity

• Electronegativity: The ability of an atom to attract electrons in a covalent bond.

• Increases up and to the right on the periodic table (O > N > S ≈ C ≈ H ≈ P).

Ionic Bonds

• Ionic bonds: Formed when electrons are transferred from one atom to another, resulting in oppositely charged ions that attract each other.

• Cation: Positively charged ion (loses electrons).

• Anion: Negatively charged ion (gains electrons).

Example: Sodium (Na) donates an electron to chlorine (Cl), forming Na+ and Cl-, which combine to make NaCl (table salt).

Continuum of Bond Types

• The degree of electron sharing in chemical bonds forms a continuum from equal sharing (nonpolar covalent) to complete transfer (ionic).

Bond Type	Electron Sharing	Example	Charge Distribution
Nonpolar Covalent	Equal	H2, O2	No charge
Polar Covalent	Unequal	H2O, NH3	Partial charges (δ+, δ-)
Ionic	Transferred	NaCl	Full charges (+, -)

Additional info: Hydrogen bonds, though not covered in detail here, are weak attractions between partially charged regions of polar molecules, such as between water molecules.