Comprehensive Study Notes: Atomic Structure, Bonding, Water, and Organic Chemistry

Basic Atomic Structure

• Atoms are composed of sub-atomic particles:

  • Protons — positively charged

  • Neutrons — neutral

  • Electrons — negatively charged

• Protons and neutrons are located in the nucleus; electrons are found in orbitals surrounding the nucleus

The Elements

• Elements are pure substances that contain only one type of atom

• 92 naturally-occurring elements are known

• Only about 25 are found in living organisms

• Six elements compose 99% of the atoms of all living organisms:

  • Oxygen (O), Carbon (C), Hydrogen (H), Nitrogen (N), Calcium (Ca), Phosphorus (P)

Isotopes

• Isotopes are alternative forms of an element with the same number of protons but different numbers of neutrons

• Additional neutrons make some isotopes unstable

• To achieve a more stable state, many isotopes release tiny energetic particles from their nuclei; emissions are known as radiation

• Examples:

  • Carbon-12: ^{12}_{6} ext{C} ext{ has } 6p + 6n ext{ (stable/non-radioactive)}

  • Carbon-14: ^{14}_{6} ext{C} ext{ has } 6p + 8n ext{ (unstable/radioactive)}

Radioactivity

• Gradual disintegration of unstable atoms releases sub-atomic particles and energy as radioactivity

• Types of radiation:

  • Alpha radiation: slow-moving; stopped by paper or skin; hazardous when ingested

  • Beta radiation: faster-moving; stopped by foil or glass; hazardous when ingested

  • Gamma radiation: travels at the speed of light; stopped by concrete, lead, or steel; damages skin and internal organs

  • Neutron radiation: released by nuclear fission; travels through air; damages skin and internal organs

• Nuclear explosions release radiation:

  • Initial radiation: Gamma and neutron radiation in the first minute

  • Residual radiation: fission products with varying half-lives release beta particles and gamma radiation

• Source reference: AtomicArchive (link provided in original transcript)

Radiation exposure in the United States

• Diagnostic radiation

• Therapeutic radiation

• Radiometric dating

• Applications for the use of radioactive isotopes include: (details not provided in the transcript)

Chemical bonding

• The reactivity of an atom is determined by the number of electrons in its outermost (valence) electron shell

• Atoms with full valence shells are non-reactive (e.g., neon)

• Atoms with partly full valence shells can interact with other atoms through chemical bonding to form molecules (e.g., oxygen)

• Atoms can share one or more pairs of bonding electrons

Molecules and Compounds

• Molecule: two or more atoms chemically joined by covalent bonds (can be same or different elements) – examples: ext{H}2, ext{H}2 ext{O}, ext{NH}_3, ext{HCl}

• Compound: two or more atoms of different elements chemically joined by covalent or ionic bonds – example: ext{NaCl}

Covalent Bonds – sharing electrons

• Nonpolar covalent bond: electrons shared equally; bond is symmetrical (e.g., ext{H}_2)

• Polar covalent bond: electrons shared unequally because one atom is more electronegative; bond is asymmetrical (e.g., ext{H}_2 ext{O})

Electronegativity

• Electronegativity: measure of how tightly an atom in a molecule holds electrons

• Higher electronegativity → electrons held more tightly → partial negative charge (denoted as d^{-}) on the atom

• Lower electronegativity → electrons held less tightly → partial positive charge (denoted as d^{+}) on the atom

• Example: oxygen tends to attract bonding electrons more strongly than hydrogen

Ion formation

• Atoms have no net charge when the number of protons equals the number of electrons

• Ions form when valence electrons are donated or accepted to fill the outermost shell; electrons are not shared in this case (permanent transfer)

• Cation: loses an electron; becomes positively charged (e.g., ext{Na}^+, ext{K}^+, ext{Mg}^{2+}, ext{Ca}^{2+})

• Anion: gains an electron; becomes negatively charged (e.g., ext{Cl}^-, ext{OH}^-)

Ionic bonds – opposites attract

• Ionic bonds form between oppositely charged ions via electrostatic forces

• Ionic bonds are easily disrupted by polar molecules like water, which can pull ions apart and into solution; as water evaporates, ions can reform ionic bonds again

• Example illustration (from transcript): Sodium (yellow) and chloride (green) particles forming NaCl

The chemical bonding continuum

• The type of bond formed between two atoms is determined by the difference in electronegativity

• Larger electronegativity differences favor transfer of electrons (ionic bonds); smaller differences favor sharing of electrons (covalent bonds)

Electrolytes

• An electrolyte is any substance that is ionized in water and conducts electricity when dissolved

• Example: ext{NaCl} + ext{H}2 ext{O} ightarrow ext{Na}^+ + ext{Cl}^- ext{ (and } ext{H}2 ext{O})

• Most electrolytes are ions and are vital to the body’s proper functioning

Water: The solvent of life

• A solution is a homogeneous mixture; a solvent is the dissolving agent; the solute is what is dissolved

• An aqueous solution is one in which water is the solvent

Hydrogen bonds

• Water is a small, asymmetric, polar molecule

• Water molecules form hydrogen bonds with other polar molecules

• A hydrogen atom from water (δ−) is attracted to a more electronegative atom (δ+) in another molecule

• Opposite partial charges attract to form weak bonds that, collectively, are strong

• No electron sharing or transferring is involved

Hydrophilic vs Hydrophobic

• Hydrophilic substances are ions or polar molecules that stay in solution due to interactions with water’s partial charges (e.g., sugar)

• Hydrophobic substances are mainly nonpolar and do not interact with water well (e.g., oils)

Properties of water

1) Adhesion – water forms hydrogen bonds with other substances (e.g., plastic, cells)
2) Cohesion – water molecules form hydrogen bonds with each other
3) High surface tension – at the water–air interface, water molecules bind more tightly to each other than to air

Significance of water’s high surface tension

• Cohesive forces between molecules create surface tension

• Example context: leaf surfaces and water interactions in plants

Capillary action

• Water moves through narrow spaces due to adhesion to polar surfaces and cohesion among water molecules

• Rate of capillary action increases as diameter decreases

• Capillary action depends on adhesion, cohesion, surface tension, and gravity

• Water moves from roots to tree tops via capillary action

Practical aspects of water

• Capillary action has practical applications such as wicking (e.g., watering plants via a string or towel)

Water distribution in different states

• Water is denser as a liquid than as a solid; ice is less dense and floats on liquid water

• As water cools, hydrogen bonds form a lattice that spaces molecules apart more in ice than in liquid water

• Ice insulates the liquid water below, helping organisms survive freezing temperatures

Water’s high specific heat

• A large amount of heat is required to raise the temperature of water by 1 degree

• Large bodies of water stabilize temperatures; coastal climates are moderated by oceanic heat capacity

Water’s high heat of vaporization

• A large amount of energy is required to break hydrogen bonds and convert liquid water to vapor

• Evaporation (e.g., sweating) cools organisms as water molecules escape from the surface

The dissociation of water

• Water can undergo spontaneous reversible dissociation: ext{H}_2 ext{O}
  ightleftharpoons ext{H}^+ + ext{OH}^-

• Protons do not exist freely in solution; they associate with water to form hydronium: ext{H}2 ext{O} + ext{H}2 ext{O}
  ightarrow ext{H}_3 ext{O}^+ + ext{OH}^-

• In pure water, [H⁺] = [OH⁻]

• Acidic solutions have more H⁺ than OH⁻ (e.g., lemon juice, stomach acids)

• Basic solutions have more OH⁻ than H⁺ (e.g., ammonia, baking soda)

pH Scale

• The pH scale indicates the concentration of hydrogen ions [H⁺] in a solution

• pH is a logarithmic scale: ext{pH} = -
  abla ext{log} [ ext{H}^+]

• Acidic solutions have pH < 7; basic solutions have pH > 7; neutral solution has pH = 7

• Examples range from battery acid (very low pH) to pure water (pH ≈ 7) to household bleach (basic)

Organisms regulate the pH of cell and body fluids

• [H⁺] must be tightly regulated for proper function

• Homeostatic mechanisms maintain pH in body fluids:

  • Chemical buffers (e.g., bicarbonate) absorb or release H⁺ in solution

  • Respiratory mechanisms regulate CO₂ and bicarbonate levels

  • Renal (kidney) mechanisms absorb or secrete H⁺

The importance of carbon

• Carbon is the most versatile atom on Earth

• Four valence electrons

• Can form an almost limitless array of molecular shapes with other atoms

• Different combinations of single and double bonds create diverse structures

• Carbon–carbon bonds were crucial in chemical evolution

• Organic compounds contain mostly carbon (e.g., methane, CH₄); inorganic compounds contain little or no carbon (e.g., water, H₂O)

Functional groups: determinants of chemical behavior

• Carbon skeleton provides overall shape; functional groups attached to carbon influence chemical behavior and interactions

• Functional groups can be polar or nonpolar

Polar functional groups

• Amino groups (-NH₂ or -NH₃⁺ when ionized): accept a proton in solution

• Carboxyl group (-COOH or –COO⁻ when ionized): releases a proton in solution

• Carbonyl groups (C=O): sites that link molecules; aldehydes at molecule ends; ketones in molecule interior

• Hydroxyl groups (-OH): act as weak acids and can donate H⁺ in solution

• Phosphate groups (-PO₄^{2-}): have two negative charges when ionized; function in energy transfer or as part of nucleic acids; also good buffers (accept or donate H⁺)

• Sulfhydryl groups (-SH): can link together via covalent disulfide bonds (S–S); common in proteins

Nonpolar functional groups

• Methyl groups (-CH₃): highly stable, non-reactive; do not form hydrogen bonds and do not interact with water

• Hydrophobic molecules can have more than one functional group

Lighthearted note

• The transcript includes chemistry jokes and memes illustrating chemical ideas (e.g., Neutron joke and others) which can be used as mnemonic tools during study, though not essential to the core content