Essential Chemistry for Biology: Atoms, Elements, and Water

Chemistry Foundations for Biology

Introduction

Understanding basic chemistry is essential for studying biology, as all living things are composed of chemicals. This section introduces the chemical principles that underpin biological processes, focusing on the elements, atoms, and molecules that make up living organisms.

The Elements of Life

The Periodic Table and Biological Elements

The periodic table organizes all known chemical elements by their atomic number, symbol, and atomic mass. Elements are substances that cannot be broken down into simpler substances by chemical means.

• Matter: Anything that occupies space and has mass.

• Atomic number: The number of protons in an atom, unique to each element.

• Atomic mass: The average mass of an atom of an element, accounting for all isotopes.

• There are over 100 known elements, but only a subset are essential for life.

Elements Essential for Life

Only a few elements are abundant in living organisms. These elements are crucial for the structure and function of biological molecules.

• Major elements in biology (with symbols):

  • C = carbon

  • H = hydrogen

  • O = oxygen

  • N = nitrogen

  • P = phosphorus

  • S = sulfur

  • K = potassium

  • Ca = calcium

  • Fe = iron

  • Na = sodium

  • Cl = chlorine

  • I = iodine

• These elements make up the majority of the mass of living organisms.

Elemental Composition of Living Things

The chemical composition of living organisms is distinct from that of the Earth's crust. Biomolecules are primarily composed of a few key elements.

• Oxygen (O): 65% of body mass

• Carbon (C): 18.5%

• Hydrogen (H): 9.5%

• Nitrogen (N): 3.3%

• Calcium (Ca): 1.5%

• Phosphorus (P): 1.0%

• Other elements (K, S, Na, Cl, Mg) and trace elements (B, Cr, Co, Cu, F, I, Fe, Mn, Mo, Se, Si, Sn, V, Zn) are present in smaller amounts but are still essential.

Atoms and Subatomic Particles

Structure of the Atom

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

• Protons: Positively charged particles found in the nucleus.

• Neutrons: Neutral particles also located in the nucleus.

• Electrons: Negatively charged particles that orbit the nucleus in electron shells.

The number of protons defines the element, while the number of neutrons can vary, resulting in different isotopes. Electrons determine the atom's chemical behavior.

Counting Subatomic Particles

The atomic number and mass number are used to determine the number of subatomic particles in an atom.

• Atomic number = number of protons

• Mass number = number of protons + number of neutrons

• Number of neutrons = mass number - atomic number

• In a neutral atom, number of electrons = number of protons

Example: For carbon-12 (atomic number 6, mass number 12):

• Protons: 6

• Neutrons: 6

• Electrons: 6

Isotopes and Radioactivity

Isotopes

Isotopes are atoms of the same element that have different numbers of neutrons, and thus different mass numbers. Some isotopes are stable, while others are radioactive (radioisotopes), meaning their nuclei decay over time, releasing radiation.

• Example: Carbon-12, Carbon-13, and Carbon-14 are isotopes of carbon.

• Radioisotopes have important applications in medicine, research, and dating of ancient materials.

Chemical Bonds and Molecules

Types of Chemical Bonds

Atoms combine to form molecules through chemical bonds. The main types of chemical bonds are:

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

• Covalent bonds: Formed when two atoms share one or more pairs of electrons. Covalent bonds are the strongest type of chemical bond.

• Hydrogen bonds: Weak attractions between a hydrogen atom with a partial positive charge and another atom with a partial negative charge (often oxygen or nitrogen). Hydrogen bonds are important in the structure of water and biological molecules.

Ionic Bonding

Ionic bonds occur when atoms transfer electrons to achieve full outer electron shells, resulting in the formation of ions.

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

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

• Example: Sodium chloride (NaCl) is formed by the ionic bond between Na+ and Cl-.

Covalent Bonding

Covalent bonds involve the sharing of electron pairs between atoms. Molecules such as H2, O2, and H2O are held together by covalent bonds.

• Single covalent bond: Sharing of one pair of electrons (e.g., H-H in H2).

• Double covalent bond: Sharing of two pairs of electrons (e.g., O=O in O2).

• Structural formulas and space-filling models are used to represent molecules.

Polar and Nonpolar Molecules

When electrons are shared unequally in a covalent bond, the molecule becomes polar, with partial positive and negative charges at different ends. Water (H2O) is a classic example of a polar molecule.

• Polar molecules dissolve well in water (hydrophilic).

• Nonpolar molecules do not dissolve well in water (hydrophobic).

Water: Structure and Properties

Water as a Polar Molecule

Water's unique properties arise from its polar structure and ability to form hydrogen bonds.

• Oxygen atom is more electronegative, pulling electrons closer and creating a partial negative charge.

• Hydrogen atoms have a partial positive charge.

• This polarity allows water molecules to form hydrogen bonds with each other and with other substances.

Properties of Water

• Cohesion: Water molecules stick together due to hydrogen bonding, aiding in the transport of water in plants.

• Surface tension: Water has a high surface tension, making it difficult to break the surface.

• High specific heat: Water can absorb or release large amounts of heat with little temperature change, helping to stabilize temperatures in organisms and environments.

• Ice floats: Solid water (ice) is less dense than liquid water due to the arrangement of hydrogen bonds, allowing ice to float and insulate aquatic life.

• Solvent properties: Water dissolves many substances, especially polar and ionic compounds, making it the universal solvent in biological systems.

Solutions, Acids, and Bases

A solution is a homogeneous mixture of two or more substances. The solvent is the dissolving agent (often water), and the solute is the substance being dissolved.

• Hydrophilic substances: Dissolve easily in water (e.g., salts, sugars).

• Hydrophobic substances: Do not dissolve in water (e.g., oils, fats).

pH, Acids, and Bases

The pH scale measures the concentration of hydrogen ions (H+) in a solution, ranging from 0 (most acidic) to 14 (most basic), with 7 being neutral.

• Acids: Substances that increase H+ concentration (pH < 7).

• Bases: Substances that decrease H+ concentration (pH > 7).

• Salts: Compounds formed from the neutralization of an acid and a base.

Buffer Systems

Buffers are substances that help maintain a stable pH in biological systems by neutralizing excess acids or bases. The bicarbonate buffer system is important in maintaining blood pH.

• Example equation for the bicarbonate buffer system:

$\mathrm{CO_2 + H_2O \leftrightarrow H_2CO_3 \leftrightarrow H^+ + HCO_3^-}$

Summary Table: Major Elements in Biology

Element	Symbol	Approximate % of Body Mass	Biological Role
Oxygen	O	65%	Component of water and organic molecules; involved in cellular respiration
Carbon	C	18.5%	Backbone of all organic molecules
Hydrogen	H	9.5%	Component of water and most organic molecules
Nitrogen	N	3.3%	Component of proteins and nucleic acids
Calcium	Ca	1.5%	Bone and teeth structure, muscle function
Phosphorus	P	1.0%	Component of nucleic acids, ATP, and cell membranes
Potassium	K	0.4%	Nerve function, fluid balance
Sulfur	S	0.3%	Component of some amino acids and proteins
Sodium	Na	0.2%	Nerve function, fluid balance
Chlorine	Cl	0.2%	Fluid balance, stomach acid
Magnesium	Mg	0.1%	Enzyme function

Additional info: Some explanations and examples have been expanded for clarity and completeness, including the summary table and buffer system equation.