Introduction to Chemical Principles in Biology

This section explores the foundational chemical concepts that underlie biological processes. Understanding atoms, elements, compounds, and chemical bonds is essential for studying life at the molecular level.

Concept 2.1: Matter, Elements, and Compounds

Definition of Matter

• Matter: 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 unique properties different from their constituent elements.

• Example: Water (H2O) is a compound made from hydrogen and oxygen.

Essential and Trace Elements

• Essential elements are required for life in large amounts. In humans, these include:

  • Carbon (C), Oxygen (O), Hydrogen (H), Nitrogen (N) (~96% of body mass)

  • Calcium (Ca), Phosphorus (P), Potassium (K), Sulfur (S), Sodium (Na), Chlorine (Cl), Magnesium (Mg) (~4%)

• Trace elements are required in minute quantities (<0.01%).

Concept 2.2: Atomic Structure and Properties

Atoms and Subatomic Particles

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

• Composed of subatomic particles:

  • Neutrons: No electrical charge; contribute to isotopes.

  • Protons: Positive charge; determine the element's identity.

  • Electrons: Negative charge; involved in bonding.

• Atoms are electrically neutral overall (number of protons = number of electrons).

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

• Atomic mass: Approximate total mass of an atom, measured in daltons.

• Electrons have negligible mass.

Isotopes and Radioactivity

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

• Radioactive isotopes: Unstable isotopes that decay spontaneously, emitting particles and energy.

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

• Applications: Used in medical diagnostics and radiometric dating.

Concept 2.3: Chemical Bonds and Molecular Interactions

Types of Chemical Bonds

• Covalent bonds: Atoms share pairs of valence electrons; can be single, double, or triple bonds.

• Ionic bonds: Atoms transfer electrons, resulting in oppositely charged ions (cations and anions) that attract each other.

• Hydrogen bonds: Weak attractions between a hydrogen atom covalently bonded to an electronegative atom (like O or N) and another electronegative atom.

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

Covalent Bonds: Polar and Nonpolar

• Nonpolar covalent bond: Electrons are shared equally between atoms.

• Polar covalent bond: Electrons are shared unequally, creating partial charges (e.g., in H2O).

• Electronegativity: An atom's ability to attract shared electrons in a bond.

Ionic Compounds (Salts)

• Formed by ionic bonds between cations and anions.

• Often found as crystals in nature.

• Not considered molecules; formula indicates the ratio of elements.

• Stable when dry, but dissociate easily in water.

Hydrogen Bonds and Van der Waals Forces

• Hydrogen bonds: Important in stabilizing the structures of proteins and nucleic acids.

• Van der Waals forces: Significant when many such interactions occur together (e.g., gecko feet adhesion).

Hybridization and Molecular Shape

• Hybridization: The mixing of atomic orbitals to form new hybrid orbitals, influencing molecular shape.

• Molecular shape determines function in biological systems.

Making and Breaking Bonds

• Chemical reactions: Processes that make and break chemical bonds, converting reactants to products.

• Reactions are reversible; products can become reactants in the reverse reaction.

• Chemical equilibrium: The point at which forward and reverse reaction rates are equal, and concentrations of reactants and products remain stable.

Half-Life Calculation Example

• Given: 60 grams of Np-240, half-life = 1 hour, time elapsed = 4 hours.

• Number of half-lives: 4

• Amount remaining: 60 g × (1/2)4 = 3.75 g

Summary Table: Types of Chemical Bonds

Additional info: Some explanations and examples have been expanded for clarity and completeness, including the half-life calculation and the summary table of bond types.

Chapter 3: Water and Life

Introduction

Water is a fundamental molecule for all known forms of life. Its unique chemical and physical properties make it essential for biological processes and the maintenance of life on Earth. This chapter explores the molecular structure of water, its emergent properties, and its role as a solvent, as well as the concepts of acids, bases, and pH.

Structure and Polarity of Water

• Water Molecule: Composed of two hydrogen atoms covalently bonded to one oxygen atom (H2O).

• Polar Covalent Bonds: The electrons in the O-H bonds spend more time near the oxygen atom, making it partially negative and the hydrogens partially positive.

• Polarity: Water is a polar molecule, meaning it has an uneven distribution of charge.

• Hydrogen Bonding: The polarity allows water molecules to form hydrogen bonds with each other, which are weak individually but strong collectively.

Definition: Electronegativity is the tendency of an atom to attract electrons in a covalent bond.

Four Emergent Properties of Water

1. Cohesive Behavior

   • Cohesion: Hydrogen bonds hold water molecules together, resulting in high surface tension (the difficulty of breaking the surface of a liquid).

   • Adhesion: Water molecules can also be attracted to other substances, such as plant cell walls, helping counteract gravity in plants.

   • Example: Water moving up plant stems (capillary action).

2. Ability to Moderate Temperature

   • High Specific Heat: Water can absorb or release a large amount of heat with only a slight change in its own temperature.

   • Specific Heat: The amount of heat required to raise 1 gram of a substance by 1°C. For water, it is 1 cal/(g·°C).

   • Heat Absorption and Release: Heat is absorbed when hydrogen bonds break and released when they form.

   • Definition: Kinetic energy is the energy of motion; thermal energy is the kinetic energy associated with the random movement of atoms or molecules.

3. Expansion Upon Freezing

   • Ice Floats: Water is less dense as a solid than as a liquid because hydrogen bonds form a crystalline lattice that spaces molecules apart.

   • Biological Importance: If ice sank, bodies of water would freeze solid, making life impossible in aquatic environments.

4. Versatility as a Solvent

   • Solution: A homogeneous mixture of substances.

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

   • Solute: The substance that is dissolved.

   • Hydration Shell: When ionic compounds dissolve, each ion is surrounded by water molecules.

   • Water dissolves: Ionic and polar (hydrophilic) substances, and large molecules with ionic/polar regions.

   • Water does not dissolve: Non-polar (hydrophobic) substances, such as oils.

Hydrophilic vs. Hydrophobic Substances

• Hydrophilic: Substances that have an affinity for water (usually polar or charged).

• Hydrophobic: Substances that repel water (usually non-polar), such as oils and many cell membrane components.

Solute Concentration and Molarity

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

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

Formula:

• $\text{Molarity (M)} = \frac{\text{Number of moles of solute}}{\text{Number of liters of solution}}$

• Example: To make a 1 M NaCl solution, dissolve 58 g of NaCl (molecular mass = 58 g/mol) in 1 L of water.

Acids, Bases, and pH

Acid-Base Chemistry in Water

• Acid: Increases the H+ concentration in a solution (pH < 7).

• Base: Reduces the H+ concentration in a solution (pH > 7).

• pH Scale: Measures the concentration of H+ ions; ranges from 0 (most acidic) to 14 (most basic).

• Neutral Solution: [H+] = [OH-], pH = 7.

• Strong acids/bases: Dissociate completely in water.

• Weak acids/bases: Reversibly release and accept H+ ions.

Formulas:

• $\text{pH} = -\log[\text{H}^+]$

• Example: A solution with [H+] = 0.0235 M has pH = $-\log(0.0235) = 1.63$.

Buffers

• Definition: Buffers are substances that minimize changes in concentrations of H+ and OH- in a solution.

• Composition: Usually consist of a weak acid and its corresponding base.

• Function: Maintain stable pH in biological systems, crucial for cellular processes.

• Example: The bicarbonate buffer system in blood.

Ocean Acidification

• Definition: The process by which excess atmospheric CO2 dissolves in seawater, forming carbonic acid and lowering ocean pH.

• Biological Impact: Reduces carbonate available for marine organisms to build shells and skeletons.

• Environmental Concern: Threatens marine biodiversity and ecosystem stability.

Summary Table: Properties of Water

Key Terms

• Polar molecule

• Hydrogen bond

• Cohesion

• Adhesion

• Specific heat

• Solvent

• Solute

• Hydrophilic

• Hydrophobic

• Molarity