Chapter 2: Basic Chemistry – Foundations for Biology

2.1 Chemical Elements

Introduction to Matter and Elements

Understanding the chemical basis of life begins with the study of matter and elements. Matter is anything that has mass and occupies space, and it exists in three states: solid, liquid, and gas. All matter, whether living or non-living, is composed of basic substances called elements.

• Matter: Anything with mass and volume.

• Element: A pure substance that cannot be broken down into substances with different properties; composed of one type of atom.

• There are 92 naturally occurring elements.

• Six elements (acronym CHNOPS) make up 95% of the body weight of organisms: Carbon, Hydrogen, Nitrogen, Oxygen, Phosphorus, Sulfur.

Atoms and Subatomic Particles

An atom is the smallest unit of an element that retains the properties of that element. Atoms are composed of subatomic particles:

• Protons: Positively charged, 1 atomic mass unit (amu), located in the nucleus.

• Neutrons: No charge, 1 amu, located in the nucleus.

• Electrons: Negatively charged, negligible mass, located in electron shells around the nucleus.

Particle	Electric Charge	Atomic Mass Unit (AMU)	Location
Proton	+1	1	Nucleus
Neutron	0	1	Nucleus
Electron	-1	0	Electron shell

Atomic Number, Mass Number, and Atomic Mass

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

• Mass Number: Sum of protons and neutrons in the nucleus.

• Atomic Mass: Approximately equal to the mass number.

• Elements are represented by unique atomic symbols (e.g., H for hydrogen, Na for sodium).

The Periodic Table

The periodic table organizes elements by increasing atomic number (periods) and groups elements with similar chemical properties (groups).

• Rows are called periods; columns are called groups.

• Group VIII (noble gases) are inert.

• Atoms in the table are electrically neutral (number of protons = number of electrons).

Isotopes

Isotopes are atoms of the same element with different numbers of neutrons, resulting in different atomic masses. Some isotopes are radioactive and decay spontaneously, emitting energy.

• Example: Carbon isotopes (12C, 13C, 14C).

• Radioactive isotopes are used in biological research and medicine.

2.2 Molecules and Compounds

Definitions and Examples

A molecule consists of two or more atoms bonded together. A compound is a molecule containing at least two different elements. The formula of a compound indicates the number and type of atoms present.

• Molecule: Smallest part of a compound retaining its chemical properties (e.g., H2, NaCl).

• Compound: Molecule with different elements (e.g., CO2, H2O, C6H12O6).

• Formula: Shows the number of each atom (e.g., C6H12O6 for glucose).

Atom Type	Number in Glucose (C6H12O6)
Carbon	6
Hydrogen	12
Oxygen	6

Chemical Bonding

Bonds between atoms are formed by interactions of electrons in the outermost shells. The process of bond formation is called a chemical reaction.

Types of Chemical Bonds

• Ionic Bonds: Formed when electrons are transferred from one atom to another, resulting in oppositely charged ions that attract each other (e.g., NaCl formation).

• Covalent Bonds: Formed when atoms share electrons to achieve stability.

• Nonpolar Covalent Bonds: Electrons are shared equally (e.g., H2, O2, CH4).

• Polar Covalent Bonds: Electrons are shared unequally, creating partial charges (e.g., H2O).

Bond Type	Electron Sharing	Example
Ionic	Transferred	NaCl
Nonpolar Covalent	Equally Shared	O2, H2
Polar Covalent	Unequally Shared	H2O

2.3 Chemistry of Water

Structure and Polarity

Water is a polar molecule due to the unequal sharing of electrons between oxygen and hydrogen atoms. This polarity allows water molecules to form hydrogen bonds with each other.

• Hydrogen Bond: Weak attraction between a slightly positive hydrogen atom and a slightly negative atom (often oxygen or nitrogen).

• Hydrogen bonds are crucial for the structure and function of proteins and DNA.

Properties of Water

• High Heat Capacity: Water absorbs large amounts of heat with little temperature change, helping organisms maintain stable internal temperatures.

• High Heat of Vaporization: Water requires significant energy to evaporate, aiding in cooling mechanisms (e.g., sweating).

• Solvent Properties: Water dissolves many substances due to its polarity. Hydrophilic molecules dissolve in water, while hydrophobic molecules do not.

• Cohesion: Water molecules stick to each other, contributing to surface tension.

• Adhesion: Water molecules stick to other polar surfaces, enabling capillary action.

• Density of Ice: Ice is less dense than liquid water, allowing it to float and insulate aquatic environments.

2.4 Acids and Bases

pH and Ionization

The pH scale measures the concentration of hydrogen ions (H+) in a solution, ranging from 0 (acidic) to 14 (basic), with 7 being neutral. Water ionizes to release equal numbers of H+ and OH- ions.

• Acids: Substances that release H+ ions in water.

• Bases: Substances that take up H+ ions or release OH- ions.

• The pH scale is logarithmic; each unit change represents a tenfold change in H+ concentration.

Buffers

Buffers are chemicals that help maintain pH within narrow limits, essential for biological systems. Human blood, for example, is maintained at a pH of 7.4.

• If blood pH drops below 7.0, acidosis occurs; above 7.8, alkalosis occurs—both are life-threatening.

• Buffers help prevent drastic changes in pH.

Additional info:

• Bohr models and diagrams are used to visualize electron distribution in atoms, which is essential for understanding chemical bonding.

• Structure-function relationships, such as water's polarity and hydrogen bonding, are central themes in biology.