Ap bio unit 1 topic 1

Chemistry Review

Matter

  • Definition: Matter is anything that takes up space and has mass.

  • Examples: Rocks, metals, oils, gases, organisms, etc. are all forms of matter.

Elements and Compounds

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

    • Natural Elements: There are 92 elements that occur in nature, organized in the periodic table.

  • Compound: A substance consisting of two or more different elements combined in a fixed ratio.

    • Examples:

    • Water (H₂O)

    • Sodium Chloride (NaCl)

Essential and Trace Elements

  • Essential Elements: Of the 92 naturally occurring elements, 20-25% are essential for survival and reproduction. The elements CHOPN (Carbon, Hydrogen, Oxygen, Phosphorus, Nitrogen) make up 96% of living matter.

  • Trace Elements: These are required in very small quantities; examples include:

    • Iron: Important for red blood cells to carry oxygen.

    • Zinc: Crucial for the immune system.

    • Iodine: Regulates metabolism through thyroid hormones.

    • Copper: Necessary for making blood cells and maintaining healthy nerves.

    • Fluoride: Strengthens teeth and prevents cavities.

Atomic Structure

  • Atomic Number: The number of protons in an atom.

  • Atomic Mass: The number of protons plus neutrons averaged over all isotopes.

Periodic Table

  • Elements are organized with:

    • Periods (horizontal rows): Elements in the same row have the same total number of electron shells.

    • Groups (vertical columns): Elements in the same column have the same number of valence electrons.

Types of Bonds

Stability through Chemical Bonds

  • Elements form chemical bonds to achieve stability through various methods, guided by the Octet Rule: Elements gain, lose, or share electrons to complete their valence shell (similar to noble gases).

  • Valence Shell: Outermost layer of electrons in an atom.

Types of Bonds

  1. Covalent Bonds: Formed when two or more atoms share electrons, typically between nonmetals.

    • Types of Covalent Bonds:

      • Single Bond: 1 pair of shared electrons.

      • Double Bond: 2 pairs of shared electrons.

      • Triple Bond: 3 pairs of shared electrons.

    • Nonpolar Covalent Bond: Electrons are shared equally.

      • Example: Oxygen gas (O₂).

    • Polar Covalent Bond: Electrons are shared unequally, resulting in partial charges.

      • Example: Water (H₂O).

  2. Ionic Bonds: Formed by the attraction between oppositely charged ions, typically between a metal and a nonmetal, resulting in the transfer of electrons.

    • Cation: Positively charged ion (e.g., Na⁺).

    • Anion: Negatively charged ion (e.g., Cl⁻).

    • Examples: Sodium Chloride (NaCl), Lithium Fluoride (LiF).

  3. Hydrogen Bonds: An attraction between a partially positive hydrogen atom in one polar covalent molecule and an electronegative atom in another.

    • Forms intermolecular bonds, holding water molecules together.

    • Hydrogen bonds are responsible for the unique properties of water.

Properties of Water

  1. Polarity: Created by unequal sharing of electrons between oxygen and hydrogen in water molecules.

  2. Cohesion: Attraction between molecules of the same kind, enabling water transport against gravity in plants and contributing to surface tension.

    • Example: Hydrogen bonds hold water molecules (H₂O) together.

  3. Adhesion: Attraction to other polar or charged molecules, allowing water to cling to cell walls.

    • Facilitates water movement in plants (e.g., xylem).

  4. Capillary Action: Upward movement of water due to cohesive and adhesive forces. Adhesion is greater than cohesion.

  5. Temperature Control:

    • High Specific Heat: Water resists changes in temperature, stabilizing environments (benefits marine life).

    • Requires heat to break hydrogen bonds and releases heat when they form.

    • High Heat of Vaporization: Water requires significant energy to evaporate, allowing for evaporative cooling processes in organisms (e.g., sweating in humans).

  6. Density: Ice expands and is less dense than liquid water, allowing aquatic life to survive under ice sheets.

    • Hydrogen bonds cause water molecules to form a crystalline structure in ice.

  7. Solvent Properties: Water is known as the universal solvent due to its ability to dissolve many polar compounds and ionic substances, adhering to the principle "like dissolves like".

    • Solvent: The dissolving agent

    • Solute: The substance that gets dissolved.

  8. pH and Buffers:

    • pH: A measure of the acidity or basicity of a solution.

      • Acid: Releases hydrogen ions (H⁺) in water.

      • Base: Accepts H⁺ or releases hydroxide ions (OH⁻).

    • Buffers: Solutions that resist pH changes, maintaining stability in biological systems.

Concept Check Questions

  1. Potassium Chloride in Water: Mixing would result in dissolution due to ionic nature and lower electronegativity.

  2. Water's Property for Marine Life: Water holds oxygen and absorbs carbon, facilitating respiration for aquatic organisms.

  3. Impact of More Dense Ice: If ice were denser than water, it would sink, disrupting plant growth and marine food chains.