Chapter 2: The Chemical Context of Life

Introduction to Biology and Chemistry

Understanding the chemical basis of life is essential for studying biology. Biological systems are governed by chemical principles, including the structure and behavior of atoms, molecules, and the interactions between them.

• Order: Living organisms exhibit complex organization, from molecules to cells.

• Energy: Life requires energy transformations, such as metabolism and ATP production.

• Growth & Reproduction: Cells grow and reproduce through processes like cellular respiration, photosynthesis, and the cell cycle.

• Evolution: Genetic changes drive evolution and diversity.

• Response to Environment & Regulation: Organisms respond to stimuli and regulate internal conditions.

Learning Outcomes

• Recognize that matter is composed of elements and atoms.

• Describe an atom’s structure based on its atomic number and atomic mass.

• Understand valence electrons and their role in chemical bonding.

• Distinguish between covalent, ionic, hydrogen bonds, and van der Waals forces.

• Explain electronegativity and its influence on chemical behavior.

Atoms

Definition and Structure

An atom is the smallest unit of matter that retains the properties of an element. Atoms are composed of subatomic particles: protons, neutrons, and electrons.

• Element: A substance that cannot be broken down by chemical means.

• Proton: Positively charged particle located in the nucleus.

• Neutron: Neutral particle located in the nucleus.

• Electron: Negatively charged particle orbiting the nucleus.

Example: A helium atom has 2 protons, 2 neutrons, and 2 electrons.

Periodic Table

Organization and Biological Relevance

The periodic table organizes elements by atomic number and properties. Certain elements are especially important for life, such as carbon, hydrogen, oxygen, and nitrogen.

• Major elements in the human body: Oxygen (65%), Carbon (18.5%), Hydrogen (9.5%), Nitrogen (3.3%)

• Trace elements: Required in small amounts (e.g., iron, iodine)

Example: Oxygen is the most abundant element in the human body.

Atomic Number and Atomic Mass

Definitions and Calculations

Atomic number is the number of protons in an atom, which defines the element. Atomic mass (mass number) is the sum of protons and neutrons.

• Atomic number (Z): Number of protons

• Atomic mass (A): Number of protons + number of neutrons

Equation:

Example: Helium (He): Atomic number = 2, Atomic mass = 4

Electron Configuration and Valence Electrons

Electron Shells and Stability

Electrons are arranged in shells around the nucleus. Valence electrons are those in the outermost shell and determine chemical reactivity.

• First shell holds up to 2 electrons.

• Second shell holds up to 8 electrons.

• Atoms are most stable when their outer shell is full.

Example: Carbon has 4 valence electrons; oxygen has 6.

Isotopes and Radioactivity

Isotopes

Isotopes are atoms of the same element with different numbers of neutrons, resulting in different atomic masses.

• Stable isotopes: Do not change over time.

• Radioactive isotopes: Decay spontaneously, emitting radiation.

Example: Carbon-12 (6 protons, 6 neutrons) vs. Carbon-14 (6 protons, 8 neutrons)

Application: Radioactive isotopes are used in dating fossils and medical imaging.

Chemical Bonds

Covalent Bonds

Covalent bonds involve the sharing of valence electrons between atoms, resulting in stable molecules.

• Single covalent bond: Sharing one pair of electrons (e.g., H-H)

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

• Triple covalent bond: Sharing three pairs of electrons (e.g., N≡N)

Example: Carbon dioxide (CO2):

Polarity and Electronegativity

Electronegativity is an atom’s ability to attract electrons in a bond. Differences in electronegativity lead to polar or nonpolar covalent bonds.

• Nonpolar covalent bond: Equal sharing of electrons (e.g., H2)

• Polar covalent bond: Unequal sharing, resulting in partial charges (e.g., H2O)

Equation:

Example: Water (H2O) is polar because oxygen is more electronegative than hydrogen.

Ionic Bonds

Ionic bonds result from the transfer of electrons from one atom to another, creating charged ions that attract each other.

• Cation: Positively charged ion (e.g., Na+)

• Anion: Negatively charged ion (e.g., Cl-)

• Example: Sodium chloride (NaCl) forms from Na+ and Cl-

Hydrogen Bonds

Hydrogen bonds are weak attractions between a hydrogen atom (with partial positive charge) and an electronegative atom (such as oxygen or nitrogen).

• Important in water, DNA, and protein structure.

• Individually weak, but collectively strong in biological systems.

Example: Hydrogen bonding between water molecules gives water its unique properties.

Van der Waals Interactions

Van der Waals forces are weak attractions due to temporary shifts in electron distribution, allowing molecules to stick together.

• Significant in large molecules and biological macromolecules.

Example: Geckos use van der Waals forces to adhere to surfaces.

Summary Table: Types of Chemical Bonds

Key Terms

• Atom

• Element

• Isotope

• Ion

• Covalent bond

• Ionic bond

• Hydrogen bond

• Van der Waals force

• Electronegativity

• Valence electron

Additional info: These notes expand on fragmented points and images from the original materials, providing full academic context and definitions for General Biology students.