The Chemical Basis of Life: Atoms, Molecules, and Water

Chapter 2: The Chemical Basis of Life, I: Atoms, Molecules, and Water

Key Concepts:

  • Atoms
  • Chemical Bonds and Molecules
  • Properties of Water
  • pH and Buffers

Atoms

  • Definition: Atoms are the smallest functional units of matter that form all chemical substances.
  • Characteristics: An atom is the smallest unit of an element that retains the chemical properties of that element.
  • Chemical Elements: Each specific type of atom is a chemical element.
    • Examples: Nitrogen, Oxygen, Helium.

Composition of Atoms

  • Atoms consist of three subatomic particles:
    1. Protons:
    • Charge: Positive (+)
    • Location: Found in the nucleus
    1. Neutrons:
    • Charge: Neutral (no charge)
    • Location: Found in the nucleus
    1. Electrons:
    • Charge: Negative (−)
    • Location: Found in orbitals surrounding the nucleus

Atomic Structure

Distinguishing Elements

  • Atomic Number:
    • Definition: The number of protons in an atom.
    • Characteristics:
    • Atomic number also equals the number of electrons in a neutral atom (net charge of zero, except in ions).
  • Atomic Mass:
    • Definition: Approximately equal to the sum of protons and neutrons in the atom.

Ions and Isotopes

  • Ions: Charged atoms or molecules.
    • Cations: Atoms that have lost electrons (net positive charge, e.g., Ca²⁺).
    • Anions: Atoms that have gained electrons (net negative charge).
  • Isotopes: Different forms of an element with varying numbers of neutrons.
    • Examples: Carbon-12 (6 protons, 6 neutrons), Carbon-14 (6 protons, 8 neutrons).

Electrons and Orbitals

  • Electron Orbitals: Regions surrounding the nucleus where electrons are likely to be found.
    • Visualization: Atoms have been historically visualized as miniature solar systems, an oversimplification, but useful.
  • Electron Cloud Model: Electrons travel within orbitals that can be represented as a cloud around the nucleus.

Rutherford's Experiment

  • Alpha Particles: Composed of two protons and two neutrons.
  • Experiment Overview: Rutherford passed alpha particles through gold foil to study atomic structure.
  • Detection: The descent pattern of alpha particles was detected using zinc sulfide screens surrounding the foil.
  • Conclusion: 98% of alpha particles passed through undeflected, demonstrating that atoms consist mainly of empty space.

Orbital Structure

Types of Orbitals

  • s Orbitals: Spherical shape, with each holding 2 electrons.
  • p Orbitals: Dumbbell-shaped (propeller), with three sub-orbitals that can each hold 2 electrons (total of 6 in 2p).

Electron Shells

  • Electron Shells:
    • 1st shell: 1s orbital (holds 2 electrons).
    • 2nd shell: 1s (2 electrons) and 3 p orbitals (up to 6 electrons total).
  • Valence Electrons: Electrons in the outermost shell available for bonding with other atoms.

The Periodic Table

Organization

  • Rows (Periods): Indicate the number of electron shells.
  • Columns (Groups): Indicate the number of valence electrons.
    • Example: Column 1 has 1 valence electron, Column 2 has 2, etc.

Atomic Mass

Mass Information

  • Definition: Atomic mass is the weighted average of the masses of the isotopes of an element.
  • Protons and Neutrons: Nearly equal in mass; both are approximately 1,800 times that of an electron.

Isotopes

  • Definition: Multiple forms of an element differing by neutron count.
  • Example: Carbon has an atomic mass average of 12.011 due to its isotopes.

Mass vs. Weight

  • Mass: The amount of matter in an object.
  • Weight: The gravitational pull on a given mass, which varies by location (e.g., weight on the moon vs. Earth).

Measurement Units

  • Dalton: Unit for atomic mass (1 Da = 1/12 the mass of a carbon atom).
  • Mole: 1 mole of any element = Avogadro's number of atoms.

Essential Elements in Living Organisms

  • Primary Elements: Hydrogen, Oxygen, Carbon, Nitrogen make up 95% of living matter.
  • Trace Elements: Less than 0.01% yet crucial for function (e.g., essential minerals).

Chemical Bonds and Molecules

  • Definition of Molecule: Two or more atoms bonded together; Compound: Molecule with two or more different elements (e.g., C6H12O6).
  • Molecular Formula: Represents composition using chemical symbols and subscripts (e.g., H2O).

Types of Chemical Bonds

  1. Covalent Bonds:
    • Electrons are shared to satisfy valence shells.
    • Types of covalent bonds:
      • Polar Covalent: Unequal sharing of electrons.
      • Nonpolar Covalent: Equal sharing of electrons.
  2. Ionic Bonds:
    • Electrons are transferred, creating charged ions (cations and anions).
  3. Hydrogen Bonds:
    • Attraction between a hydrogen atom in a polar molecule and an electronegative atom in another.

Covalent Bond Example

  • Single Bond: 1 pair of shared electrons (e.g., H-F).
  • Double Bond: 2 pairs of shared electrons (e.g., O=O).
  • Triple Bond: 3 pairs of shared electrons (e.g., N≡N).

Octet Rule

  • Atoms are stable with full outer shells; most require 8 electrons (with Hydrogen and Helium as exceptions).

Polar and Nonpolar Bonds

  • Polar Bonds: Electrons shared unequally, leading to partial positive and negative charges.
  • Nonpolar Bonds: Electrons shared equally with no charge difference.

Water Molecule Properties

  • Water's polar covalent nature leads to its behavior in biological systems (e.g., solvent capabilities, hydrogen bonding).

Hydrogen Bonding

  • Hydrogen bonds are weak bonds that form between polar molecules; they can collectively create strong interactions, as seen in DNA structure.

Ionic Bonds

  • Formation of cations and anions by electron loss/gain; ionic compounds (salts) result from the attraction of oppositely charged ions.

Chemical Reactions

  • Defined as the transformation of substances into different substances (reactants to products).
  • Properties: require energy, often need catalysts (enzymes), tend toward equilibrium, and predominantly occur in aqueous solutions.

Properties of Water

  • Solutions: Composed of solutes dissolved in solvents; water as the universal solvent.
  • Hydrophilic vs. Hydrophobic:
    • Hydrophilic: Substances that dissolve in water.
    • Hydrophobic: Substances that do not dissolve in water.

Amphipathic Molecules

  • Molecules that contain both polar (hydrophilic) and nonpolar (hydrophobic) portions.
  • Example: Detergents can form micelles in water due to their amphipathic nature.

Solution Concentration

  • Concentration: Solute amount per unit volume.
  • Molarity: Number of moles of solute dissolved in 1 liter of solution (1 M NaCl = 58.4 g NaCl per liter).

Water's Physical Properties

  • Exhibits high specific heat, high heat of vaporization, and low density in solid form (ice floats).

Effects of Solutes on Water Properties

  • Adding solutes can lower the freezing point and raise the boiling point of water.

Water's Biological Functions

  • Functions in lubrication, chemical reactions, waste removal, evaporative cooling, and providing structural support.

Acids and Bases

  • Acids: Release H+ ions; strong acids increase H+ more than weak.
  • Bases: Lower H+ concentration by releasing OH- or binding H+.

pH Scale

  • Measures H+ concentration on a scale from 0-14; 7 is neutral,
  • Common substance pH examples are human stomach fluid (1.3), orange juice (3.5), etc.

Buffer Systems

  • Maintain stable pH in biological systems, shifting to absorb or release H+ as needed.
  • Important for organisms to regulate pH within narrow ranges (e.g., human blood pH between 7.35-7.45).