General Biology: Chemical Context of Life and Properties of Water

Chapter 2: Chemical Context of Life

Living Things in an Aqueous Environment

All living organisms exist in environments where water is present, and their chemistry is fundamentally based on the properties of water and the elements that compose them.

• Matter: Anything that has mass and occupies space.

• The universe is composed of matter, energy, and space.

Elements and Compounds

• Elements: Pure substances that cannot be broken down by chemical means. There are 92 naturally occurring elements.

• Major elements in living organisms:

  • H - Hydrogen

  • O - Oxygen

  • C - Carbon

  • N - Nitrogen

  • Together, these make up about 96% of essential elements in living things.

• Compounds: Substances consisting of two or more elements in a fixed ratio (e.g., CO2, H2O, NaCl, MgCl2).

Elements of Life

• Essential elements: 25-25% of naturally occurring elements are essential for life.

• Bulk elements (most of living things):

  • S - Sulfur

  • Ca - Calcium

  • K - Potassium

  • P - Phosphorus

  • Other elements in trace amounts (e.g., Fe - Iron, I - Iodine)

Atoms and Subatomic Particles

An atom is the smallest unit of matter that retains the properties of an element.

• Atoms are made of subatomic particles:

  • Protons: Positive charge (+1), located in the nucleus

  • Neutrons: No charge (0), located in the nucleus

  • Electrons: Negative charge (-1), orbit the nucleus

• Atoms are electrically neutral when the number of protons equals the number of electrons.

• Atoms are highly reactive and unstable when not balanced.

Atomic Nucleus and Atomic Number

• The atomic nucleus contains protons and neutrons.

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

• Mass number: Sum of protons and neutrons.

• Dalton: Unit of atomic mass, defined as 1/12th the mass of a carbon-12 atom.

Isotopes and Radioactivity

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

• Stable isotopes: Do not change over time.

• Radioactive isotopes: Unstable, decay over time, emitting radiation.

• Half-life: Time required for half the atoms in a radioactive sample to decay.

• Radioactive dating: Using isotopes (e.g., C-14) to estimate the age of biological materials.

Energy Levels and Electron Shells

• Energy: The capacity to cause change.

  • Potential energy: Stored energy due to position or structure.

  • Kinetic energy: Energy of motion.

  • Heat (thermal) energy: Random movement of atoms/molecules.

  • Photons of light: Main source of energy for life, used in photosynthesis.

• Law of Conservation of Energy: Energy cannot be created or destroyed, only transformed.

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

  • First shell: 2 electrons

  • Second and third shells: 8 electrons each

The Periodic Table

• Organizes elements by atomic number and electron configuration.

• 8 columns (groups), 3 rows (periods) in the simplified table.

• Elements in the same column have the same number of electrons in their outer shell (valence electrons).

• Valence shell: Outermost electron shell; determines chemical reactivity.

• Atoms with full valence shells are inert (non-reactive).

Atomic Orbitals and Molecular Shapes

• Electron orbitals define the regions where electrons are likely to be found.

• First shell: Spherical (s orbital)

• Second shell: s and p orbitals (p orbitals have a dumbbell shape)

• Shapes of orbitals influence molecular geometry (e.g., tetrahedral shape in methane).

Chemical Bonds

• Strong Bonds:

  • Covalent bonds: Atoms share pairs of electrons; very strong; form molecules.

  • Ionic bonds: Electrons are transferred from one atom to another, creating charged ions that attract each other.

• Weak Bonds:

  • Hydrogen bonds: Weak attractions between a hydrogen atom and an electronegative atom (e.g., O or N).

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

Covalent Bonds and Molecules

• Single, double, and triple covalent bonds:

  • Single bond: One pair of shared electrons (e.g., H2)

  • Double bond: Two pairs of shared electrons (e.g., O2)

  • Triple bond: Three pairs of shared electrons (e.g., N2)

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

• Compounds: Molecules containing at least two different elements.

Valence and Electronegativity

• Valence: Number of covalent bonds an atom can form (e.g., C = 4, H = 1, O = 2, N = 3).

• Electronegativity: The tendency of an atom to attract electrons in a covalent bond.

Polar and Nonpolar Covalent Bonds

• Polar covalent bond: Electrons are shared unequally, resulting in partial charges (δ+ and δ−) on atoms (e.g., H2O).

• Nonpolar covalent bond: Electrons are shared equally; no partial charges (e.g., CH4).

Ionic Bonds and Ionization

• Formed by transfer of electrons, creating cations (+) and anions (−).

• Example: NaCl (table salt)

  • Na (atomic number 11) donates an electron to Cl (atomic number 17).

  • Na+ (cation), Cl− (anion)

• Dissociation: Ionic compounds like NaCl dissolve in water, separating into ions.

Weak Bonds in Biology

• Hydrogen bonds: Important in water, DNA, and protein structure.

• Van der Waals interactions: Help geckos climb walls, stabilize protein structure.

Chemical Reactions and Equilibrium

• Chemical reactions: Atoms are rearranged; matter is conserved.

• Chemical equilibrium: Forward and reverse reaction rates are equal; concentrations of reactants and products remain constant.

• Dynamic equilibrium: Reactions continue, but no net change in concentrations.

Chapter 3: Water and Life

Phases of Water

• Solid: Ice

• Liquid: Water

• Gas: Water vapor/steam

Properties of Water

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

• Adhesion: Water molecules stick to other substances.

• Surface tension: Water has a high surface tension due to hydrogen bonding; difficult to stretch or break the surface.

• Moderation of temperature: Water absorbs and releases heat slowly, helping to stabilize temperatures in organisms and environments.

• Floating of ice: Ice is less dense than liquid water, so it floats.

Biological Importance of Water's Properties

• Water's cohesion and adhesion are essential for transport in plants.

• High surface tension allows small organisms to move on water's surface.

• Moderation of temperature helps maintain stable conditions for life.

• Ice floating insulates bodies of water, protecting aquatic life in winter.

Medical Application: Infant Respiratory Distress Syndrome (IRDS)

• Premature babies may lack surfactant, a substance that reduces surface tension in the lungs.

• Without surfactant, alveoli collapse, making breathing difficult.

• Surfactant therapy is used to treat IRDS in premature infants.

Summary Table: Types of Chemical Bonds

Bond Type	Description	Strength	Example
Covalent	Atoms share electrons	Strong	H2O, CH4
Ionic	Transfer of electrons; attraction between ions	Strong (in dry conditions)	NaCl
Hydrogen	Attraction between H and electronegative atom	Weak	Between water molecules
Van der Waals	Transient local charges	Very weak	Gecko feet adhesion

Key Equations

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

• Law of Conservation of Mass: $\text{Total mass of reactants} = \text{Total mass of products}$

• pH and pOH relationship: $\text{pH} + \text{pOH} = 14$

Additional info:

• Some diagrams and tables were inferred and expanded for clarity.

• Examples and applications were added to illustrate biological relevance.