Chapter 2: The Chemical Context of Life

Introduction

Understanding biology requires a solid foundation in chemistry, as all biological processes are governed by chemical principles. This chapter explores the fundamental chemical concepts essential for studying life, including the nature of matter, atomic structure, elements, and the types of chemical bonds that form the basis of biological molecules.

Matter, Elements, and Compounds

Definitions and Properties

• Matter: Anything that takes up space and has mass. All living and non-living things are composed of matter.

• Element: A substance that cannot be broken down into other substances by chemical reactions. Each element is defined by its number of protons.

• Compound: A substance consisting of two or more elements combined in a fixed ratio. Compounds have characteristics different from those of their constituent elements.

Example: Sodium (Na) and chlorine (Cl) are elements; when combined, they form sodium chloride (NaCl), a compound with properties distinct from either element.

The Elements of Life

Major and Minor Elements

• Of the 92 naturally occurring elements, only a small number are essential for life.

• Major elements: Carbon (C), Hydrogen (H), Oxygen (O), and Nitrogen (N) make up about 96% of living matter.

• Other essential elements: Calcium (Ca), Phosphorus (P), Potassium (K), Sulfur (S), Sodium (Na), and a few others are required in smaller amounts.

Additional info: Trace elements are required by organisms in minute quantities, such as iron (Fe) and iodine (I).

Atoms and Subatomic Particles

Atomic Structure

• Atom: The smallest unit of matter that retains the properties of an element.

• Composed of subatomic particles:

  • Protons (p+): Positively charged, located in the nucleus.

  • Neutrons (n): No charge, located in the nucleus.

  • Electrons (e-): Negatively charged, orbit the nucleus in electron shells.

• Protons and neutrons have nearly identical mass, measured in atomic mass units (amu).

Atomic Number and Mass Number

• Atomic number (Z): Number of protons in the nucleus; defines the element.

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

• Atomic mass: The atom’s total mass, often approximated by the mass number.

Example: Carbon-12 has 6 protons and 6 neutrons; its mass number is 12.

Isotopes and Radioactivity

Isotopes

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

• All isotopes of an element have the same number of protons but differ in mass number.

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

Radioactive Isotopes

• Radioactive isotopes: Unstable isotopes that decay spontaneously, emitting radiation.

• Used as diagnostic tools in medicine (e.g., radioactive tracers in imaging).

• Can be used to date ancient biological materials (radiometric dating).

Example: Fludeoxyglucose (radioactive glucose) is used to detect cancerous tissues in PET scans.

Electron Arrangement and Chemical Properties

Energy Levels and Electron Shells

• Potential energy: Energy that matter possesses due to its location or structure.

• Electrons occupy energy levels called electron shells around the nucleus.

• The chemical behavior of an atom is determined by the distribution of electrons, especially those in the outermost shell (valence shell).

• Atoms are most stable when their valence shell is full (2 electrons for the first shell, 8 for the second and third).

Electron Orbitals

• Orbital: A three-dimensional space where an electron is found 90% of the time.

• Each electron shell consists of a specific number of orbitals (e.g., 1s, 2s, 2p).

The Periodic Table and Chemical Reactivity

Organization and Trends

• Elements are arranged by increasing atomic number.

• Rows (periods) correspond to the number of electron shells.

• Columns (groups) contain elements with similar valence electron configurations and chemical properties.

Additional info: Elements in the same group often form similar types of bonds and have comparable reactivity.

Chemical Bonds

Covalent Bonds

• Covalent bond: A strong chemical bond formed when two atoms share one or more pairs of valence electrons.

• Single bond: Sharing of one pair of electrons.

• Double bond: Sharing of two pairs of electrons.

• Electronegativity: The attraction of an atom for the electrons in a covalent bond.

• Nonpolar covalent bond: Electrons are shared equally between atoms.

• Polar covalent bond: Electrons are shared unequally, resulting in partial charges.

Example: In water (H2O), oxygen is more electronegative than hydrogen, resulting in polar covalent bonds.

Ionic Bonds

• Ionic bond: Formed when one atom transfers an electron to another, resulting in oppositely charged ions that attract each other.

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

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

• Compounds formed by ionic bonds are called ionic compounds or salts.

Example: Sodium chloride (NaCl) forms from the transfer of an electron from sodium to chlorine.

Weak Chemical Interactions

• Hydrogen bonds: Weak bonds that form when a hydrogen atom covalently bonded to one electronegative atom is attracted to another electronegative atom (commonly oxygen or nitrogen).

• Van der Waals interactions: Weak attractions between molecules or parts of molecules that result from transient local partial charges.

Example: Hydrogen bonding is critical for the properties of water and the structure of DNA.

Molecular Shape and Function

Three-Dimensional Structure

• The shape of a molecule is determined by the positions of its atoms’ orbitals.

• Molecular shape is crucial for biological function, such as enzyme-substrate specificity and hormone-receptor binding.

• Molecules with similar shapes can mimic each other and bind to the same biological receptors.

Example: Morphine and endorphins have similar shapes and can bind to the same brain receptors.

Chemical Reactions

Making and Breaking Bonds

• Chemical reaction: The making and breaking of chemical bonds, leading to changes in the composition of matter.

• Reactants: Starting materials in a chemical reaction.

• Products: Resulting materials after the reaction.

• Chemical reactions are reversible; products can become reactants in the reverse reaction.

• Chemical equilibrium: The point at which the forward and reverse reactions occur at the same rate, and the concentrations of reactants and products remain constant.

Example: Photosynthesis:

Redox Reactions

• Oxidation-reduction (redox) reactions: Chemical reactions that involve the transfer of electrons between substances.

• Oxidation: Loss of electrons.

• Reduction: Gain of electrons.

Additional info: Redox reactions are fundamental to cellular respiration and photosynthesis.

Table: Comparison of Bond Types