Chemical Foundations of Life: Atomic Structure, Bonding & Properties of Water

Matter is defined as having mass and occupying space.
Atoms: The smallest unit of matter that retains the properties of an element.

Components of Atoms:
- Protons: Positively charged particles (+)
- Neutrons: Neutral particles (no charge)
- Electrons: Negatively charged particles (-)
Bohr Model: A visual representation of an atom, often depicted using an electron shell diagram.

Table of Subatomic Particles:
- Proton: Charge = +1, Mass = 1 amu, Location = Nucleus
- Neutron: Charge = 0, Mass = 1 amu, Location = Nucleus
- Electron: Charge = -1, Mass = negligible, Location = Orbitals surrounding nucleus
- Note: 1 amu = $1.67 imes 10^{-24}$ grams.

Atomic Number: The number of protons in an atom; unique to each element.
Atomic Mass: Sum of protons and neutrons.
Charge Neutrality: An atom is neutral when the number of protons equals the number of electrons.
Chemical Properties: Determined by electrons in the outermost shell.

Elements: Substances that cannot be broken down into simpler substances by ordinary chemical means. They possess specific chemical and physical properties.
Different atoms are identified by their number of protons, atomic number, and unique characteristics.

Mass Number: Equal to the number of protons plus neutrons in the nucleus.
Atomic Number: Equals the number of protons and identifies the element.
Periodic Table Information: Displays elements, atomic numbers, and mass numbers.

Electrons: Negatively charged particles found in orbitals located around the nucleus.
Neutral Atoms: Have equal numbers of protons and electrons; this balance results in no net charge.

Ions: Charged particles formed when an atom gains or loses electrons.
- Cations: Formed when there are more protons than electrons, resulting in a net positive charge.
- Anions: Formed when there are fewer protons than electrons, resulting in a net negative charge.

Isotopes: Variants of a single element that contain differing numbers of neutrons in the nucleus.
Radioactive Isotopes: Unstable isotopes that emit radiation as their nuclei disintegrate.
- Half-life: The time required for half of the radioactive nuclides in a sample to decay.

Common Isotopes: 12C, 13C, 14C
Carbon-14 Dating:
- Carbon-14 is absorbed by living organisms; its decay over time allows scientists to date organic materials.
- Half-lives for Carbon-14: 100% at death, 50% after 5,730 years, 25% after 11,460 years, 12.5% after 17,190 years.

Electron Shells:
- First shell: Holds a maximum of 2 electrons.
- Second shell: Holds a maximum of 8 electrons.
Arrangement of Electrons: Key to an atom's chemical behavior; electrons are found in various orbitals.
Covalent Bonds: Form when electrons are shared between atoms, creating stable associations.
Types of Covalent Bonds:
- Single Bonds: One pair of shared electrons.
- Double Bonds: Two pairs of shared electrons.
- Triple Bonds: Three pairs of shared electrons.
Examples of Covalent Compounds: Methane (CH₄), Ammonia (NH₃), Water (H₂O).

Definition: Formed by the electrostatic attraction between oppositely charged ions.
Cation vs. Anion:
- Cations are formed by the loss of electrons (e.g., Na⁺ from sodium atom).
- Anions are formed by the gain of electrons (e.g., Cl⁻ from chlorine atom).

Definition: An atom's tendency to attract electrons. Higher electronegativity corresponds to greater attraction for electrons.
Electronegativity Values: Ranges vary among different elements, with the most electronegative elements found in the upper right corner of the periodic table.

Nonpolar Covalent Bonds: Characterized by equal sharing of electrons (e.g., O₂, N₂) due to equal electronegativity.
Polar Covalent Bonds: Characterized by unequal sharing of electrons (e.g., H₂O), where one atom has a partial negative charge (δ−) and the other a partial positive charge (δ+).
Examples of partial charges: In water, oxygen has a δ− charge while hydrogen has a δ+ charge.

Cohesion: Water molecules stick to each other via hydrogen bonds.
Adhesion: Water molecules stick to other polar molecules due to hydrogen bonds.
Universal Solvent: Water can dissolve many substances, particularly polar molecules and ions but not nonpolar molecules.
Hydrogen Bonding: Occurs between partially charged regions of water molecules, contributing to water's unique properties.

High Specific Heat: Water can absorb significant heat without a considerable change in temperature.
High Heat of Vaporization: Water requires a lot of heat to evaporate, which helps in temperature regulation.
Density: Solid water (ice) is less dense than liquid water, allowing it to float.
Hydrophobic Exclusion: Water organizes nonpolar molecules through exclusion.
Ion Formation: Water can dissociate into hydrogen ions (H⁺) and hydroxide ions (OH⁻).

Definition: The pH scale measures the concentration of hydrogen ions in a solution, ranging from acidic (0) to basic (14).
pH Calculation:
- At 25°C, pure water has a concentration of $[H^+] = 10^{-7}$ moles per liter, corresponding to a neutral pH of 7.
pH Examples:
- Hydrochloric acid: pH 0 (very acidic)
- Pure water: pH 7 (neutral)
- Seawater: pH 8 (slightly basic)
- Sodium hydroxide: pH 14 (very basic)

Problem: What is the pH of a solution with an H⁺ concentration of $1.0 imes 10^{-3}$ M?
- Answer: pH = 3.
Problem: Comparing pH in plant cells, with vacuole pH at 3 and cytosol pH at 7, calculate the ratio of [H⁺] concentration:
- Answer: Correct answer is 10,000 times higher in the vacuole compared to cytosol.
Problem: pH of Great Salt Lake (~10) compared to streams (~7) indicates a H⁺ concentration ratio of:
- Answer: 1,000 times lower in the Great Salt Lake than in the streams.