Chapter 2: The Chemical Context of Life

Introduction

This chapter explores the fundamental chemical principles that underlie biological processes. Understanding the nature of atoms, elements, compounds, and chemical bonds is essential for studying life at the molecular level.

Concept 2.1: Matter Consists of Chemical Elements in Pure Form and in Combinations Called Compounds

Definition of Matter

• Matter is anything that takes up space and has mass.

• All living organisms are composed of matter.

Elements and Compounds

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

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

• Example: Formic acid (CH2O2) is a compound made of carbon (C), hydrogen (H), and oxygen (O).

Elements in the Human Body

The human body is primarily composed of a few key elements. The following table summarizes their relative abundance:

Additional info: Trace elements such as boron, chromium, cobalt, copper, fluorine, iodine, manganese, molybdenum, selenium, silicon, tin, vanadium, and zinc are present in amounts less than 0.01% of body mass.

Concept 2.2: An Element’s Properties Depend on the Structure of Its Atoms

Atomic Structure

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

• Atoms are composed of subatomic particles:

  • Neutrons: No electrical charge

  • Protons: Positive charge

  • Electrons: Negative charge

• Protons and neutrons form the atomic nucleus; electrons form a "cloud" around the nucleus.

• Proton mass and neutron mass are nearly identical and are measured in daltons.

Atomic Number and Atomic Mass

• Atomic number: Number of protons in the nucleus; determines the element’s identity.

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

• Atomic mass: Total mass of the atom, approximately equal to the mass number.

• Example: Oxygen has 8 protons (atomic number 8).

Isotopes

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

• Radioactive isotopes: Decay spontaneously, giving off particles and energy.

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

Concept 2.3: The Formation and Function of Molecules Depend on Chemical Bonding Between Atoms

Chemical Bonds

• Atoms with incomplete valence shells can share or transfer valence electrons with other atoms, resulting in chemical bonds.

• Chemical bond: Attraction that holds atoms together in molecules or compounds.

Covalent Bonds

• Covalent bond: Sharing of a pair of valence electrons by two atoms.

• Single covalent bond: Sharing of one pair of electrons.

• Double covalent bond: Sharing of two pairs of electrons.

• Molecule: Two or more atoms held together by covalent bonds.

• Structural formula: Notation representing atoms and bonding (e.g., H—H).

• Molecular formula: Abbreviated notation (e.g., H2).

Electronegativity and Bond Polarity

• Electronegativity: An atom’s attraction for electrons in a covalent bond.

• Nonpolar covalent bond: Atoms share electrons equally.

• Polar covalent bond: One atom is more electronegative, causing unequal sharing of electrons and partial charges.

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

Ionic Bonds

• Sometimes atoms strip electrons from their bonding partners, resulting in ions.

• Ion: A charged atom or molecule.

• Cation: Positively charged ion.

• Anion: Negatively charged ion.

• Ionic bond: Attraction between an anion and a cation.

• Example: Sodium (Na) transfers an electron to chlorine (Cl), forming Na+ and Cl-, which combine to form sodium chloride (NaCl).

Hydrogen Bonds

• Hydrogen bond: Weak attraction between a hydrogen atom covalently bonded to a more electronegative atom (such as oxygen or nitrogen) and another electronegative atom.

• Example: Hydrogen bonds form between water molecules, contributing to water’s unique properties.

Additional info:

• Electron shells and valence electrons play a key role in determining chemical reactivity. Elements with full valence shells are chemically inert.

• The periodic table organizes elements by atomic number and electron configuration, which influences their chemical behavior.