The Chemical Context of Life: Atoms, Elements, and Chemical Bonds

Chapter 2: The Chemical Context of Life

Overview: A Chemical Connection to Biology

Biology is deeply intertwined with the principles of chemistry and physics. Living organisms are composed of matter, which is organized into a hierarchy from atoms to molecules to cells. At each level, new properties emerge that are not present at the previous level, illustrating the concept of emergent properties.

Concept 2.1: Matter, Elements, and Compounds

Definition of Matter, Elements, and Compounds

• Matter: Anything that takes up space and has mass.

• Element: A substance that cannot be broken down into other substances by chemical reactions.

• Compound: A substance consisting of two or more elements in a fixed ratio, with properties different from its constituent elements (emergent properties).

Example: Table salt (NaCl) is a compound formed from sodium (a metal) and chlorine (a gas), resulting in an edible substance with new properties.

The Elements of Life

• About 20–25% of the 92 natural elements are essential for life.

• Major elements: Carbon, hydrogen, oxygen, and nitrogen make up 96% of living matter.

• Other important elements: Calcium, phosphorus, potassium, and sulfur account for most of the remaining 4%.

• Trace elements: Required in minute quantities (e.g., iron, iodine).

Element	Symbol	Percentage of Body Mass (including water)
Oxygen	O	65.0%
Carbon	C	18.5%
Hydrogen	H	9.5%
Nitrogen	N	3.3%
Calcium	Ca	1.5%
Phosphorus	P	1.0%
Potassium	K	0.4%
Sulfur	S	0.3%
Sodium	Na	0.2%
Chlorine	Cl	0.2%
Magnesium	Mg	0.1%

Concept 2.2: Atomic Structure and Properties

Atoms and Subatomic Particles

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

• Protons: Positively charged, found in the nucleus.

• Neutrons: No charge, found in the nucleus.

• Electrons: Negatively charged, form a cloud around the nucleus.

Atomic Number, Mass Number, and Isotopes

• Atomic number: Number of protons in the nucleus (also equals the number of electrons in a neutral atom).

• Mass number: Sum of protons and neutrons.

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

• Radioactive isotopes: Unstable isotopes that decay, emitting particles and energy.

Formulas:

• Atomic number: $\text{Atomic number} = \text{Number of protons}$

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

• Number of neutrons: $\text{Number of neutrons} = \text{Mass number} - \text{Atomic number}$

Energy Levels and Electron Shells

Electrons have potential energy based on their distance from the nucleus. The farther an electron is from the nucleus, the higher its potential energy. Electrons exist in discrete energy levels called electron shells. Electrons can move between shells by absorbing or releasing energy.

Electron Distribution and the Periodic Table

The chemical behavior of an atom is determined by the distribution of electrons in its shells, especially the outermost shell (valence shell). The periodic table arranges elements by increasing atomic number and shows their electron configurations.

Electron Orbitals

An orbital is a three-dimensional space where an electron is likely to be found. Each shell contains a specific number of orbitals, and each orbital can hold up to two electrons. The arrangement of orbitals influences how atoms interact and bond.

Concept 2.3: Chemical Bonds and Molecular Structure

Covalent Bonds

Covalent bonds involve the sharing of pairs of valence electrons between atoms. These bonds can be single, double, or triple, depending on the number of shared electron pairs. Covalent bonds can form between atoms of the same or different elements.

• Nonpolar covalent bond: Electrons are shared equally (e.g., H2).

• Polar covalent bond: Electrons are shared unequally, resulting in partial charges (e.g., H2O).

Ionic Bonds

Ionic bonds form when one atom transfers an electron to another, resulting in oppositely charged ions (cations and anions) that attract each other. Ionic compounds, such as salts, are formed by these bonds.

Weak Chemical Interactions

• Hydrogen bonds: Weak attractions between a hydrogen atom covalently bonded to an electronegative atom (like O or N) and another electronegative atom. These are crucial for stabilizing DNA, protein structures, and water's properties.

• Van der Waals interactions: Weak attractions due to temporary partial charges from uneven electron distribution. Individually weak, but collectively significant in stabilizing large molecules and enabling adhesion (e.g., gecko feet).

Molecular Shape and Function

The shape of a molecule is determined by the positions of its atoms' orbitals and is critical for its function. Molecular shape affects how molecules interact and recognize each other, such as in enzyme-substrate binding or hormone-receptor interactions.

Summary Table: Types of Chemical Bonds

Bond Type	Description	Relative Strength	Example
Covalent	Sharing of electron pairs	Strong	H2O, O2
Ionic	Transfer of electrons, attraction between ions	Strong (in dry conditions)	NaCl
Hydrogen	Attraction between H and electronegative atom	Weak (individually)	Between water molecules
Van der Waals	Temporary partial charges	Very weak (individually)	Gecko adhesion

Key Terms

• Element

• Compound

• Atom

• Isotope

• Covalent bond

• Ionic bond

• Hydrogen bond

• Van der Waals interaction

• Electronegativity

• Valence electrons

Additional info: This guide expands on the provided notes with definitions, examples, and context to ensure a comprehensive understanding of the chemical basis of life, suitable for exam preparation in a General Biology course.