Chapter 2: The Chemical Context of Life – Study Notes

Chapter 2: The Chemical Context of Life

Introduction

This chapter explores the fundamental chemical principles that underlie biological processes. Understanding the nature of matter, elements, compounds, and the unique properties of water 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 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 emergent properties that are different from those of their constituent elements.

• Example: Sodium (Na) and chlorine (Cl) are both dangerous in pure form, but together they form sodium chloride (NaCl), or table salt, which is safe to eat.

Elements of Life

• About 20–25% of the 90+ natural elements are essential for life.

• Essential elements are required for an organism to live a healthy life and reproduce.

• Trace elements are required in minute quantities (e.g., iodine for thyroid function in vertebrates).

• Deficiency Example: Lack of iodine can cause goiter in humans.

Table: Elements in the Human Body

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%
Sodium	Na	0.2%
Chlorine	Cl	0.2%
Magnesium	Mg	0.1%
Trace elements	-	<0.01%

Evolution of Tolerance to Toxic Elements

• Some elements are toxic to organisms (e.g., arsenic).

• Some species have adapted to environments with high concentrations of toxic elements (e.g., sunflowers can absorb heavy metals and are used in phytoremediation).

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 charge), protons (positive charge), and electrons (negative charge).

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

• Proton and neutron mass ≈ 1 dalton each.

Atomic Number and Atomic Mass

• 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 in the nucleus.

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

• Example (Sodium): Atomic number = 11, Mass number = 23, Number of neutrons = 12.

Isotopes

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

• Radioactive isotopes: Decay spontaneously, releasing particles and energy.

• Applications: Dating fossils, tracing atoms in metabolism, diagnosing medical disorders.

• Radiation from decaying isotopes can damage cellular molecules.

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, forming chemical bonds.

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

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

• Single bond: Sharing of one pair of electrons (e.g., H–H).

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

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

• Nonpolar covalent bond: Electrons are shared equally.

• Polar covalent bond: Electrons are shared unequally, leading to partial charges.

• Ionic bond: Attraction between oppositely charged ions (cation and anion) formed by electron transfer.

• Ionic compounds (salts): Compounds formed by ionic bonds, often found as crystals.

Weak Chemical Interactions

• Weak bonds (hydrogen bonds, van der Waals interactions) are important for the structure and function of large biological molecules.

• Hydrogen bond: Forms when a hydrogen atom covalently bonded to one electronegative atom is attracted to another electronegative atom (usually oxygen or nitrogen).

• Van der Waals interactions: Weak attractions due to transient local partial charges.

Molecular Shape and Function

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

• Molecular shape is crucial for the recognition and interaction of biological molecules (e.g., enzyme-substrate, hormone-receptor).

• Molecules with similar shapes can have similar biological effects (e.g., morphine and endorphins).

Concept 2.4: Chemical Reactions Make and Break Chemical Bonds

Chemical Reactions

• Chemical reactions involve the making and breaking of chemical bonds.

• Reactants: Starting materials in a chemical reaction.

• Products: Final materials in a chemical reaction.

• Example: Photosynthesis: $6CO_2 + 6H_2O \xrightarrow{sunlight} C_6H_{12}O_6 + 6O_2$

• All chemical reactions are theoretically reversible.

• Chemical equilibrium: Point at which forward and reverse reaction rates are equal.

Concept 2.5: Hydrogen Bonding Gives Water Properties That Help Make Life Possible on Earth

Structure and Polarity of Water

• Water is a polar molecule: oxygen has a partial negative charge, hydrogens have partial positive charges.

• Hydrogen bonds form between water molecules, leading to unique properties.

Emergent Properties of Water

• Cohesive behavior: Water molecules stick together due to hydrogen bonding (cohesion), aiding transport in plants.

• Adhesion: Water molecules cling to other substances, such as cell walls.

• Surface tension: Measure of how difficult it is to break the surface of a liquid; water has high surface tension.

• Ability to moderate temperature: Water absorbs and releases heat with little temperature change due to high specific heat.

• Expansion upon freezing: Ice is less dense than liquid water, so it floats, insulating aquatic life below.

• Versatility as a solvent: Water dissolves many substances due to its polarity.

Temperature and Heat

• Kinetic energy: Energy of motion.

• Thermal energy: Total kinetic energy due to molecular motion.

• Temperature: Average kinetic energy of molecules.

• Heat: Thermal energy in transfer.

• Calorie (cal): Amount of heat to raise 1g of water by 1°C.

• Specific heat of water: 1 cal/(g·°C).

• Water’s high specific heat is due to hydrogen bonding.

• Evaporative cooling: As water evaporates, its surface cools, stabilizing temperatures in organisms and bodies of water.

Floating of Ice on Liquid Water

• Hydrogen bonds in ice are more ordered, making ice less dense than liquid water.

• Ice insulates water below, allowing life to persist under frozen surfaces.

• Melting of polar ice due to climate change threatens habitats and global ecosystems.

Water: The Solvent of Life

• Solution: Homogeneous mixture of substances.

• Solvent: Dissolving agent (water in aqueous solutions).

• Solute: Substance dissolved.

• Water forms hydration shells around ions and can dissolve polar molecules.

• Hydrophilic substances: Affinity for water.

• Hydrophobic substances: Do not have an affinity for water (often nonpolar).

Solute Concentration in Aqueous Solutions

• Molecular mass: Sum of masses of all atoms in a molecule.

• Mole (mol): 6.02 × 1023 molecules (Avogadro’s number).

• Molarity (M): Number of moles of solute per liter of solution.

Acids, Bases, and pH

Acids and Bases

• Water can dissociate into hydrogen ions (H+) and hydroxide ions (OH−).

• Acid: Increases H+ concentration in solution.

• Base: Reduces H+ concentration (either by accepting H+ or producing OH−).

• Strong acids/bases: Dissociate completely in water (e.g., HCl, NaOH).

• Weak acids/bases: Reversibly release or accept H+ (e.g., carbonic acid).

pH Scale

• pH is defined as $pH = -\log[H^+]$

• In pure water at 25°C, $[H^+][OH^-] = 10^{-14}$

• Acidic solutions: pH < 7; Basic solutions: pH > 7; Neutral: pH = 7

• Most biological fluids have pH between 6 and 8.

Buffers

• Buffers: Substances that minimize changes in concentrations of H+ and OH− in a solution.

• Most buffers consist of a weak acid and its corresponding base.

• Example: Carbonic acid (H2CO3) in blood helps maintain pH stability.

Ocean Acidification

• CO2 from fossil fuel combustion dissolves in oceans, forming carbonic acid and lowering ocean pH.

• Increased H+ reduces carbonate ion concentration, threatening organisms that build shells and coral reefs.

• Predicted 40% decline in carbonate ion concentration by 2100.