Chapter 2: The Chemical Context of Life – Study Notes

Chemistry of Life: Creating Compounds

Introduction to Matter, Elements, and Compounds

Organisms are composed of matter, which is anything that takes up space and has mass. The basic building blocks of matter are elements and compounds.

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

• Atom: The smallest particle of an element that retains its chemical properties.

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

• Example: Sodium (Na) and chlorine (Cl) combine to form sodium chloride (NaCl), a compound with properties distinct from either element.

Elements of Life

Essential and Trace Elements

Of the 92 naturally occurring elements, only 20-25% are essential for life. These are required for an organism to survive, grow, and reproduce.

• Essential elements: Oxygen, carbon, hydrogen, and nitrogen make up about 96% of living matter.

• Trace elements: Required in minute quantities, such as iodine, which is necessary for thyroid function. Deficiency can cause disorders like goiter.

• 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%

  

Evolution of Tolerance to Toxic Elements

Adaptation and Phytoremediation

Some elements are toxic, but certain species have evolved to tolerate or even utilize them.

• Phytoremediation: The use of plants, such as sunflowers, to absorb and detoxify heavy metals from contaminated soils.

• Example: Sunflowers were used to clean up soils contaminated with lead and zinc after environmental disasters.

Element Properties Depend on Atomic Structure

Subatomic Particles and Atomic Nucleus

The properties of elements are determined by their atomic structure.

• Subatomic particles:

  • Protons: Positive charge

  • Neutrons: No charge

  • Electrons: Negative charge

• Atomic nucleus: Contains protons and neutrons; electrons form a cloud around the nucleus.

Atomic Number and Atomic Mass

Definitions and Calculations

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: Approximated by mass number.

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

• Example: Sodium has 11 protons and 12 neutrons, so its mass number is 23.

Isotopes

Stable and Radioactive Isotopes

Isotopes are atomic forms of an element with different numbers of neutrons.

• Stable isotopes: Do not change over time.

• Radioactive isotopes: Nucleus decays spontaneously, emitting particles and energy.

• Applications: Used in dating fossils, tracing metabolic processes, and medical diagnostics (e.g., PET scans).

• Example: Carbon-12, Carbon-13, and Carbon-14 are isotopes of carbon with different numbers of neutrons.

Energy and Electrons

Electron Shells and Potential Energy

Electrons have potential energy due to their position relative to the nucleus.

• Electron shells: Electrons occupy specific energy levels (shells) around the nucleus.

• Energy absorption/release: Electrons move to higher shells by absorbing energy and to lower shells by releasing energy.

Electrons & Chemical Bonds

Valence Electrons and Reactivity

The chemical behavior of an atom is determined by the number of electrons in its outermost shell (valence shell).

• Inert atoms: Atoms with full valence shells are unreactive.

• Reactive atoms: Atoms with incomplete valence shells can share or transfer electrons, forming chemical bonds.

Covalent Bonds

Bond Formation and Molecules

Covalent bonds involve the sharing of valence electrons between atoms.

• Single bond: Sharing one pair of electrons.

• Double bond: Sharing two pairs of electrons.

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

Electronegativity & Covalent Bonds

Polar and Nonpolar Covalent Bonds

Electronegativity is the ability of an atom to attract electrons in a covalent bond.

• Polar covalent bond: Unequal sharing of electrons, resulting in partial charges.

• Electronegativity trend: Increases from bottom left to top right of the periodic table.

• Example: Water (H2O) is a polar molecule due to the higher electronegativity of oxygen.

Ionic Bonds

Formation of Ions and Ionic Compounds

Ionic bonds result from the transfer of electrons between atoms, creating charged ions.

• Cation: Positively charged ion (lost electron).

• Anion: Negatively charged ion (gained electron).

• Ionic bond: Attraction between cation and anion.

• Ionic compounds: Also called salts, such as sodium chloride (NaCl).

Weak Chemical Interactions

Hydrogen Bonds and Van der Waals Interactions

Many biological molecules are stabilized by weak interactions.

• Hydrogen bond: Attraction between a hydrogen atom covalently bonded to an electronegative atom and another electronegative atom nearby.

• Van der Waals interactions: Weak attractions due to transient asymmetrical electron distributions; important in gecko adhesion.

Chemical Reactions

Making and Breaking Bonds

Chemical reactions involve the making and breaking of covalent bonds.

• Reactants: Starting molecules.

• Products: Final molecules.

• Example: Photosynthesis: $6\ \text{CO}_2 + 6\ \text{H}_2\text{O} + \text{sunlight} \rightarrow \text{C}_6\text{H}_{12}\text{O}_6 + 6\ \text{O}_2$

Hydrogen Bonding & Water

Properties of Water

Water is a polar molecule, and hydrogen bonding between water molecules gives rise to several important properties.

• Cohesion: Water molecules stick together due to hydrogen bonding.

• Adhesion: Water molecules cling to other substances.

• Surface tension: The measure of how hard it is to break the surface of a liquid; water has high surface tension due to hydrogen bonding.

Moderation of Temperature by Water

Heat Absorption and Release

Water can absorb or release large amounts of heat with only slight changes in temperature, helping to moderate Earth's climate.

• Example: Coastal areas have milder climates due to water's heat capacity.

Floating of Ice on Liquid Water

Density and Insulation

Ice floats because hydrogen bonds in ice are more ordered, making it less dense than liquid water.

• Importance: Floating ice insulates water below, allowing aquatic life to survive in cold climates.

Water: The Solvent of Life

Solutions, Solvents, and Solutes

Water is an excellent solvent due to its polarity, dissolving many ionic and polar substances.

• Solution: Homogeneous mixture of substances.

• Solvent: The dissolving agent (water in aqueous solutions).

• Solute: The substance dissolved.

• Aqueous solution: Solution in which water is the solvent.

• Hydrophilic: Substances with affinity for water.

• Hydrophobic: Substances that repel water (usually nonpolar).

Acids, Bases, & pH

pH Scale and Buffers

The pH scale measures the concentration of hydrogen ions (H+) in a solution.

• Acids: Increase H+ concentration (pH < 7).

• Bases: Decrease H+ concentration (pH > 7).

• Buffers: Substances that minimize changes in H+ or OH- concentrations, maintaining stable pH in biological systems.

• Formula: $\text{pH} = -\log_{10}[\text{H}^+]$

• Example: Most biological fluids have pH between 6 and 8; internal pH of cells is close to 7.