Chapter 2: The Chemical Basis of Life – Essentials of Biology Study Notes

Study Guide - Smart Notes

Tailored notes based on your materials, expanded with key definitions, examples, and context.

2.1 Atoms and Atomic Bonds

Matter and Elements

Matter is anything that occupies space and has mass. It can exist in three states: solid, liquid, or gas. All matter is composed of elements, which are pure substances that cannot be broken down into other substances by ordinary chemical means.

Element: A substance that cannot be broken down into another substance by ordinary chemical means.
There are 92 naturally occurring elements.
Four elements—carbon (C), hydrogen (H), oxygen (O), and nitrogen (N)—make up about 96% of the body weight of most living organisms.

Origin of Elements

Elements are not produced by normal chemical reactions. Heavier elements, such as iron, are formed during supernova explosions of stars, which scatter these elements into space, eventually becoming part of planets and living organisms.

Example: The iron in human blood originated from the explosion of stars.

Atomic Structure

Atomic theory states that elements consist of atoms, which are the smallest units of matter retaining the properties of an element.

Atomic symbol: Abbreviation for an element (e.g., H for hydrogen, Na for sodium).
Subatomic particles:
- Neutrons: No electrical charge, found in the nucleus.
- Protons: Positive charge, found in the nucleus.
- Electrons: Negative charge, found outside the nucleus in orbitals.
Mass number: Sum of protons and neutrons; electrons have negligible mass.

Atomic Number and the Periodic Table

The atomic number is the number of protons in an atom and determines the element's identity. In a neutral atom, the number of electrons equals the number of protons.

The periodic table arranges elements by increasing atomic number in rows (periods) and columns (groups), reflecting recurring chemical and physical properties.

Isotopes

Isotopes are atoms of the same element with different numbers of neutrons, resulting in different mass numbers. Some isotopes are unstable and radioactive, emitting radiation as they decay.

Radioactive isotopes: Behave chemically like stable isotopes but can be used as tracers (e.g., PET scans), to sterilize medical equipment, or may cause cellular damage leading to cancer.

Arrangement of Electrons in an Atom

Electron Shells and Energy Levels

Electrons move constantly and occupy energy levels or shells around the nucleus. Each shell can hold a specific number of electrons:

First shell: 2 electrons
Second and subsequent shells: 8 electrons each (for atoms up to atomic number 20)

Octet Rule and Valence Shells

The octet rule states that atoms are most stable when their outermost (valence) shell contains eight electrons. Atoms will give up, accept, or share electrons to achieve a full valence shell, which determines their chemical properties.

Types of Chemical Bonds

Molecules and Compounds

A molecule is a group of atoms bonded together. A compound is a molecule containing atoms of more than one element.

Examples of molecules: O2, H2O, C6H12O6, N2
Examples of compounds: H2O, C6H12O6

Ionic Bonds

Ionic bonds form when atoms transfer electrons, resulting in charged atoms called ions. The attraction between oppositely charged ions holds them together.

Sodium (Na): Has one electron in its valence shell and usually gives it up, becoming Na+.
Chlorine (Cl): Has seven electrons in its valence shell and usually accepts one, becoming Cl-.
Ionic compounds: Often called salts (e.g., NaCl).

Covalent Bonds

Covalent bonds form when two atoms share electrons to fill their outer shells. These bonds can be single (sharing two electrons) or double (sharing four electrons).

Structural formula: Uses lines to show shared pairs (e.g., H–H).
Molecular formula: Shows the number of atoms involved (e.g., H2).
Double covalent bond: Two atoms share four electrons (e.g., O2).

Chemical Formulas and Reactions

Reactants and Products

Chemical reactions involve reactants (starting molecules) and products (molecules formed). In a chemical equation, reactants are shown to the left of the arrow, and products to the right.

Balanced Equations

A chemical equation is balanced when the same number of each type of atom appears on both sides.

Photosynthesis equation:

Molecular formula for glucose: C6H12O6

2.2 Water's Importance to Life

Role of Water in Life

Water is the most important molecule for life, making up 70–90% of all organisms. Its unique properties stem from its molecular structure.

Structure of Water

Polar covalent bond: Oxygen is more electronegative than hydrogen, so electrons are shared unequally, making oxygen slightly negative and hydrogens slightly positive.
Hydrogen bond: The slightly positive hydrogen of one water molecule is attracted to the slightly negative oxygen of another, creating a network of interactions.

Properties of Water

Solvency: Water dissolves many substances due to its polarity and hydrogen bonding.
Cohesion: Water molecules cling to each other (important for transport in plants).
Adhesion: Water molecules cling to other polar surfaces.
High surface tension: Water molecules at the surface stick together tightly.
High heat capacity: Water absorbs heat without a large change in temperature.
High heat of vaporization: It takes much energy to break hydrogen bonds for evaporation.
Varying density: Ice is less dense than liquid water, allowing it to float and insulate aquatic life.

Solvency

Hydrophilic: Molecules attracted to water.
Hydrophobic: Molecules not attracted to water.
Water causes ionic compounds like NaCl to dissociate into ions.

Cohesion and Adhesion

Cohesion: Water molecules stick to each other due to hydrogen bonding.
Adhesion: Water molecules stick to other polar surfaces.
These properties enable water transport in plants and animals.

High Surface Tension

Water molecules at the surface cling more tightly to each other than to the air above, mainly due to hydrogen bonding.
Example: Water striders can walk on water due to surface tension.

Heat Capacity and Heat of Vaporization

Water's hydrogen bonds allow it to absorb heat without a large temperature change.
Temperature of water rises and falls slowly.
High heat of vaporization means water requires much energy to evaporate, helping regulate temperature in organisms.

Varying Density

Ice is less dense than liquid water because water expands as it freezes.
Ice floats, insulating aquatic environments and making life possible beneath the surface.

2.3 Acids and Bases

Dissociation of Water

Water can dissociate into equal numbers of hydrogen ions (H+) and hydroxide ions (OH-):

Acidic Solutions

Acids release hydrogen ions (H+) or take up hydroxide ions (OH-).
Example: Hydrochloric acid dissociates as follows:

Basic Solutions

Bases take up hydrogen ions (H+) or release hydroxide ions (OH-).
Example: Sodium hydroxide dissociates as follows:

pH and the pH Scale

The pH scale measures the concentration of hydrogen ions in a solution, ranging from 0 (most acidic) to 14 (most basic).

pH below 7: Acidic (more H+ than OH-)
pH above 7: Basic (more OH- than H+)
pH of 7: Neutral (equal H+ and OH-)

Buffers and pH Regulation

Buffers are chemicals or combinations of chemicals that help maintain pH within normal limits by taking up excess H+ or OH-. The pH of human blood is tightly regulated between 7.35 and 7.45 by buffers. Failure to regulate pH can result in acidosis, which can be life-threatening.

Summary Table: Properties of Water

Property	Description	Biological Importance
Solvency	Dissolves many substances due to polarity	Facilitates chemical reactions and transport
Cohesion	Water molecules stick to each other	Enables water transport in plants
Adhesion	Water molecules stick to other surfaces	Assists movement of water in biological systems
High Surface Tension	Surface molecules cling tightly	Allows small organisms to move on water surface
High Heat Capacity	Absorbs heat with little temperature change	Stabilizes temperature in organisms and environments
High Heat of Vaporization	Requires much energy to evaporate	Helps cool organisms via evaporation
Varying Density	Ice is less dense than liquid water	Ice floats, insulating aquatic life