Chapter 2: The Chemical Context of Life – Study Notes

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Chapter 2: The Chemical Context of Life

Introduction to Matter and Elements

All living organisms are composed of matter, which exists in pure forms as elements or in combinations as compounds. Understanding the chemical basis of life is essential for studying biology.

Matter: Anything that takes up space and has mass.
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, with properties different from its constituent elements.
Molecule: Two or more atoms held together by covalent bonds.
Example: Sodium (Na) and chlorine (Cl) are elements; sodium chloride (NaCl) is a compound with different properties from its elements.

The Elements of Life

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

About 20–25% of elements are essential to life.
Major elements: Carbon (C), hydrogen (H), oxygen (O), and nitrogen (N) make up about 96% of living matter.
Other important elements: Calcium (Ca), phosphorus (P), potassium (K), and sulfur (S) make up most of the remaining 4%.
Trace elements: Required in minute quantities (e.g., iron, iodine).

Element	Symbol	Percentage of Body Mass
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%

Atomic Structure and Subatomic Particles

The properties of elements depend on the structure of their atoms, which are composed of subatomic particles.

Atom: The smallest unit of matter that retains the properties of an element.
Subatomic particles:
- Neutron: No electrical charge.
- Proton: Positive charge; determines the atom's identity.
- Electron: Negative charge.
Neutrons and protons form the atomic nucleus; electrons form a cloud around the nucleus.
Neutron and proton mass are nearly identical.

Atomic Number, Mass Number, and Isotopes

Atoms of different elements differ in their number of subatomic particles.

Atomic number: Number of protons in the nucleus.
Mass number: Sum of protons and neutrons.
Atomic mass: The atom's total mass, approximately equal to the mass number.
Isotopes: Atoms of the same element with different numbers of neutrons.
Radioactive isotopes: Unstable isotopes that decay spontaneously, emitting particles and energy.
Applications: Dating fossils, tracing metabolic processes, diagnosing medical disorders.

Isotope	Protons	Neutrons
Protium (¹H)	1	0
Deuterium (²H)	1	1
Tritium (³H)	1	2

Energy Levels of Electrons

Electrons in an atom have different amounts of potential energy, depending on their distance from the nucleus.

Energy: The capacity to cause change.
Potential energy: Energy due to location or structure.
Kinetic energy: Energy in motion.
Electrons exist in energy levels called electron shells.
Electrons in outer shells have more energy than those in inner shells.
First shell: up to 2 electrons; second shell: up to 8 electrons.

Reactivity of Atoms and Valence Electrons

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

Valence electrons: Electrons in the outermost shell.
Atoms with full valence shells are chemically inert (e.g., noble gases: helium, neon, argon).
Atoms with incomplete valence shells are reactive and tend to form chemical bonds.

Chemical Bonds and Molecules

Atoms with incomplete valence shells can share or transfer electrons, resulting in chemical bonds that hold atoms together in molecules.

Chemical bond: Attraction that holds atoms together.
Covalent bond: Sharing of a pair of valence electrons by two atoms.
Single bond: Sharing of one pair of electrons; double bond: sharing of two pairs.
Structural formula: Representation of atoms and bonds (e.g., H—H, O=O).
Molecular formula: Abbreviated representation (e.g., H2, O2).
Compound: Combination of two or more different elements.

Name & Formula	Electron Distribution	Lewis Dot/Structural Formula	Space-Filling Model
Hydrogen (H2)	2 electrons shared	H—H	[model]
Oxygen (O2)	4 electrons shared	O=O	[model]
Water (H2O)	2 single bonds	H—O—H	[model]
Methane (CH4)	4 single bonds	H—C—H	[model]

Additional info: Space-filling models visually represent the 3D structure of molecules.

Electronegativity and Polar Covalent Bonds

Atoms in a molecule attract electrons to varying degrees, a property called electronegativity.

Electronegativity: An atom's attraction for electrons in a covalent bond.
Nonpolar covalent bond: Electrons shared equally (e.g., O2, CO2).
Polar covalent bond: Electrons shared unequally, resulting in partial charges (e.g., H2O).
Unequal sharing causes partial positive (δ+) and negative (δ−) charges.

Ions and Ionic Bonds

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

Ion: Charged atom or molecule.
Cation: Positively charged ion.
Anion: Negatively charged ion.
Ionic bond: Attraction between a cation and an anion.
Ionic compounds (salts): Compounds formed by ionic bonds, often found as crystals (e.g., NaCl).

Weak Chemical Bonds

In addition to strong covalent bonds, weak chemical bonds play important roles in biological systems.

Hydrogen bond: Forms when a hydrogen atom covalently bonded to one electronegative atom is attracted to another electronegative atom (usually O or N).
Van der Waals interactions: Weak attractions between molecules due to transient local partial charges.
Weak bonds help reinforce the shapes of large molecules and enable molecules to adhere to each other.

Chemical Reactions

Chemical reactions involve the making and breaking of chemical bonds, transforming reactants into products.

Reactants: Starting molecules in a chemical reaction.
Products: Resulting molecules from a chemical reaction.
Example: Photosynthesis

In this reaction, carbon dioxide and water (reactants) are converted into glucose and oxygen (products) using sunlight.