Chapter 2 | BIO 110

Nature of Atoms

• Structure of Atoms:

  • Composed of protons, neutrons, and electrons.

  • Protons: Positive charge (located in nucleus).

  • Neutrons: No charge (located in nucleus).

  • Electrons: Negative charge (located in orbitals).

• Atomic Number:

  • Defined as the number of protons in an atom.

  • An element is characterized by its atomic number.

  • In neutral atoms, the number of protons equals the number of electrons, making it electrically neutral.

• Atomic Mass:

  • The sum of the number of protons and neutrons.

  • Each proton and neutron has a mass of approximately 1 Dalton.

• Ions:

  • Formed when atoms gain or lose electrons.

  • Cation: More protons than electrons (net positive charge).

  • Anion: Fewer protons than electrons (net negative charge).

• Isotopes:

  • Atoms with the same number of protons but different numbers of neutrons.

  • Radioactive isotopes: Unstable; emit radiation as the nucleus decays.

  • Half-life: Time taken for half of the nuclei in a sample to decay.

• Energy Levels:

  • Electrons farther from the nucleus have higher potential energy.

  • During chemical reactions, electrons can be transferred.

  • Oxidation: Loss of an electron.

  • Reduction: Gain of an electron.

• The Periodic Table:

  • Lists elements based on electron distribution.

  • Periods: Rows that indicate the number of electron shells.

  • Groups: Columns that indicate elements with identical valence electron configurations.

Elements Found in Living Systems

• There are 90 naturally occurring elements, but only 12 are abundant in living organisms.

• Key Elements:

  • Four elements (C, H, O, N) constitute 96.3% of human body weight.

  • Trace Elements: I.e., Iodine (important for thyroid function).

• Deficiencies:

  • Can lead to diseases (e.g., nitrogen deficiency in plants, iodine deficiency in humans leading to goiter).

Nature of Chemical Bonds

• Molecules: Groups of atoms held together by chemical bonds.

• Ionic Bonds:

  • Formed by the attraction between oppositely charged ions.

  • Example: Sodium chloride (NaCl) is formed by the transfer of an electron from sodium to chlorine.

• Covalent Bonds:

  • Atoms share electrons to achieve filled outer shells.

  • Single, Double, and Triple Bonds based on number of shared electrons.

• Electronegativity:

  • The affinity of an atom for electrons.

  • Affects how electrons are shared in covalent bonds:

  • Nonpolar Covalent Bonds: Equal sharing of electrons.

  • Polar Covalent Bonds: Unequal sharing of electrons, leading to partial charges.

• Chemical Reactions:

  • Involve making or breaking chemical bonds.

  • Reactants: Initial molecules; Products: Final molecules formed.

  • Influenced by temperature, concentration, and catalysts.

Water: A Vital Compound

• General Properties of Water:

  • Highly polar molecule, allowing for hydrogen bonding.

  • Cohesion: Atoms of water stick together, creating surface tension.

  • Adhesion: Water molecules stick to other polar substances.

• Thermal Properties:

  • High specific heat: Requires substantial energy to change temperature.

  • High heat of vaporization: Evaporation leads to cooling.

  • Ice is less dense than liquid water, allowing it to float and not freeze solid.

• Solvent Properties:

  • Water can dissolve many polar molecules and ions, making it crucial in biochemical reactions.

Acids and Bases

• Acids:

  • Substances that dissociate in water to increase [H+] and lower pH.

  • Stronger acids produce more hydrogen ions.

• Bases:

  • Substances that decrease [H+] in a solution by combining with hydrogen ions.

• pH Scale:

  • Ranges from 0 (acidic) to 14 (basic); 7 is neutral.

• Buffers:

  • Substances that resist changes in pH by either releasing or absorbing hydrogen ions.

  • Important for maintaining stable pH levels in biological systems.