Comprehensive Study Notes for General Chemistry: Foundations for Medical and Biological Sciences

Atoms and Atomic Theory

Atoms

Atoms are the fundamental units of matter, representing the smallest particles of an element that retain its chemical properties. Each atom consists of a dense nucleus containing protons (positively charged) and neutrons (neutral), surrounded by electrons (negatively charged) in specific regions called shells or orbitals. The arrangement of these subatomic particles determines the atom's chemical behavior.

• Protons: Positive charge, located in the nucleus.

• Neutrons: No charge, located in the nucleus.

• Electrons: Negative charge, orbit the nucleus in shells.

Atomic Theory

The atomic theory has evolved from the ancient Greek concept of indivisible particles (atomos) to Dalton's formalization in the 19th century. Key points include:

• All matter is composed of atoms.

• Atoms of a given element are identical in mass and properties.

• Atoms combine in fixed ratios to form compounds.

• Chemical reactions involve rearrangement of atoms, not their creation or destruction.

Later discoveries (Thomson, Rutherford, Bohr, and quantum mechanics) refined the model, introducing electrons, the nucleus, and probabilistic electron orbitals.

Elements and Atomic Number

An element is a pure substance consisting of only one type of atom, defined by its atomic number (number of protons). The atomic number determines the element's identity and position in the periodic table. In neutral atoms, the number of electrons equals the number of protons.

• Example: Hydrogen (atomic number 1), Oxygen (atomic number 8).

Isotopes and Atomic Weight

Isotopes are atoms of the same element with different numbers of neutrons, affecting their mass but not chemical properties. The atomic weight is the weighted average of the masses of an element's naturally occurring isotopes.

• Example: Carbon-12, Carbon-13, and Carbon-14 (used in radiocarbon dating).

The Periodic Table and Electronic Structure

The Periodic Table

The periodic table organizes elements by atomic number, electron configuration, and recurring chemical properties. It is divided into periods (rows) and groups (columns). Elements in the same group have similar chemical properties due to the same number of valence electrons.

• Group 1: Alkali metals (highly reactive, 1 valence electron).

• Group 2: Alkaline earth metals (2 valence electrons).

• Groups 3–12: Transition metals (variable oxidation states, biologically important).

• Group 17: Halogens (7 valence electrons, highly reactive nonmetals).

• Group 18: Noble gases (full valence shells, inert).

Electronic Structure and Electron Configurations

Electrons occupy shells (energy levels) around the nucleus, with each shell holding a maximum of $2n^2$ electrons (where $n$ is the shell number). Subshells (s, p, d, f) have specific capacities:

• s: 2 electrons

• p: 6 electrons

• d: 10 electrons

• f: 14 electrons

Electron configurations determine chemical behavior and periodic trends.

Electron-Dot (Lewis) Symbols

Lewis symbols represent valence electrons as dots around an element's symbol, useful for visualizing bonding and the octet rule.

• Example: Na (one dot), Cl (seven dots).

Ionic Compounds and Chemical Bonding

Ionic Compounds and the Octet Rule

Ionic compounds form when electrons are transferred from metals (which become cations) to nonmetals (which become anions), resulting in electrostatic attraction (ionic bonds). The octet rule states that atoms tend to gain, lose, or share electrons to achieve eight valence electrons.

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

• Anions: Negatively charged (e.g., Cl-).

Naming and Properties of Ionic Compounds

• Cation named first, then anion (with “-ide” suffix).

• Transition metal cations specify charge with Roman numerals.

• High melting/boiling points, crystalline structure, conduct electricity when molten or dissolved.

Acids, Bases, and Ions

• Acids: Increase H+ in solution.

• Bases: Increase OH- in solution.

• pH is determined by the balance of H+ and OH- ions.

Molecular Compounds and Covalent Bonding

Covalent Bonds and the Periodic Table

Covalent bonds involve the sharing of electrons between nonmetals. The number of bonds formed depends on the number of valence electrons.

• Group 14: 4 bonds (e.g., carbon)

• Group 15: 3 bonds (e.g., nitrogen)

• Group 16: 2 bonds (e.g., oxygen)

• Group 17: 1 bond (e.g., fluorine)

Characteristics of Molecular Compounds

• Low melting/boiling points

• Poor electrical conductivity

• Diverse structures (from simple to complex)

Lewis Structures and Molecular Formulas

• Show arrangement of atoms and valence electrons

• Follow the octet rule (except for hydrogen: duet rule)

Polarity and Electronegativity

• Polar covalent bonds: Unequal sharing due to electronegativity differences (e.g., H2O).

• Nonpolar molecules: Equal sharing or symmetric arrangement (e.g., CH4).

Naming Binary Molecular Compounds

• First element: full name; second element: root + “-ide”

• Prefixes indicate number (mono-, di-, tri-, etc.)

• Examples: CO (carbon monoxide), CO2 (carbon dioxide)

Chemical Reactions and Equations

Classification of Chemical Reactions

• Combination (Synthesis): A + B → AB

• Decomposition: AB → A + B

• Single Replacement: A + BC → AC + B

• Double Replacement: AB + CD → AD + CB

• Combustion: Hydrocarbon + O2 → CO2 + H2O

• Neutralization: Acid + Base → Salt + Water

• Redox: Electron transfer (oxidation and reduction)

Balancing Chemical Equations

• Write unbalanced equation

• Count atoms of each element

• Add coefficients to balance

• Check conservation of mass

Redox Reactions

• Oxidation: Loss of electrons

• Reduction: Gain of electrons

• Example: Cellular respiration (glucose oxidation)

Mole Concept and Chemical Calculations

The Mole and Avogadro’s Number

• 1 mole = $6.022 \times 10^{23}$ entities (Avogadro’s number)

• Links atomic/molecular scale to measurable quantities

Gram–Mole Conversions

• Molar mass: Mass of 1 mole (g/mol)

• Formula: $\text{Number of moles} = \frac{\text{Mass (g)}}{\text{Molar mass (g/mol)}}$

• Example: 36.03 g H2O / 18.015 g/mol = 2.0 moles

Stoichiometry and Solution Calculations

• Use balanced equations for mole ratios

• Molarity (M): $\text{M} = \frac{\text{moles of solute}}{\text{volume of solution (L)}}$

• Dilution: $M_1V_1 = M_2V_2$

Physical Quantities and Measurement

Metric System and Units

• Length: meter (m), centimeter (cm), millimeter (mm), kilometer (km)

• Mass: kilogram (kg), gram (g), milligram (mg), microgram (μg)

• Volume: liter (L), milliliter (mL), cubic centimeter (cm3)

Significant Figures

• All nonzero digits are significant

• Zeros between nonzero digits are significant

• Leading zeros are not significant

• Trailing zeros are significant if a decimal is present

Fundamental Chemical Laws

• Law of Conservation of Mass: Mass is neither created nor destroyed in chemical reactions.

• Law of Definite Proportions: A compound always contains the same elements in the same ratio by mass.

• Law of Multiple Proportions: When two elements form more than one compound, the masses of one element that combine with a fixed mass of the other are in small whole number ratios.

Solutions and Concentration

Mixtures and Solutions

• Homogeneous: Uniform composition (e.g., saline)

• Heterogeneous: Non-uniform (e.g., oil and water)

Units of Concentration

• Molarity (M): moles/L

• Mass/volume percent (% m/v): g/100 mL

• Parts per million (ppm)

• Molality (m): moles/kg solvent

Electrolytes

• Strong electrolytes: Completely dissociate (e.g., NaCl)

• Weak electrolytes: Partially dissociate (e.g., acetic acid)

• Non-electrolytes: Do not dissociate (e.g., glucose)

Acids, Bases, and Buffers

Acids and Bases

• Arrhenius: Acids produce H+, bases produce OH- in water

• Brønsted–Lowry: Acids are proton donors, bases are proton acceptors

Acid Dissociation Constant (Ka)

• Quantifies acid strength; strong acids fully dissociate, weak acids partially dissociate

pH and Buffer Solutions

• pH: $\text{pH} = -\log [\text{H}^+]$

• Buffer: Solution of weak acid and its conjugate base (or vice versa) that resists pH changes

• Henderson–Hasselbalch equation: $\text{pH} = \text{p}K_a + \log \left( \frac{[\text{A}^-]}{[\text{HA}]} \right)$

Organic Chemistry: Hydrocarbons and Functional Groups

Alkanes

• Saturated hydrocarbons (single bonds), general formula: $C_nH_{2n+2}$

• Isomerism: Same formula, different structures (e.g., n-butane vs. isobutane)

• Naming: Longest chain, number substituents, use prefixes (meth-, eth-, prop-, etc.)

• Reactions: Combustion, halogenation

Alkenes and Alkynes

• Alkenes: Double bonds, $C_nH_{2n}$; Alkynes: Triple bonds, $C_nH_{2n-2}$

• Cis–trans isomerism in alkenes

• Reactions: Addition (hydrogenation, halogenation, hydration)

Aromatic Compounds

• Benzene: Delocalized π-electrons, resonance, aromaticity (Hückel’s rule: $4n+2$ π-electrons)

• Naming: Substituent positions (ortho, meta, para)

• Reactions: Electrophilic aromatic substitution (halogenation, nitration, sulfonation)

Alcohols, Phenols, Ethers, Thiols, and Halides

• Alcohols: Contain –OH group; classified as primary, secondary, tertiary

• Phenols: –OH group on aromatic ring; more acidic than alcohols

• Ethers: R–O–R', relatively inert, used as solvents

• Thiols: R–SH, form disulfide bonds (important in proteins)

• Halides: R–X (X = F, Cl, Br, I), undergo substitution/elimination reactions

Amines, Aldehydes, Ketones, Carboxylic Acids, and Derivatives

Amines

• Derived from ammonia; classified as primary, secondary, tertiary

• Basicity due to lone pair on nitrogen

• Form amine salts with acids (more water-soluble)

Aldehydes and Ketones

• Contain carbonyl group (C=O); aldehydes at end, ketones within chain

• Naming: Replace “-e” with “-al” (aldehyde) or “-one” (ketone)

• Reactions: Oxidation (aldehydes → acids), reduction (to alcohols)

Carboxylic Acids and Derivatives

• Carboxyl group (–COOH), weak acids

• Derivatives: Esters (–COOR), amides (–CONH2), anhydrides

• Reactions: Esterification, amide formation, hydrolysis

Biomolecules: Amino Acids, Proteins, Enzymes, Carbohydrates, Lipids, and Nucleic Acids

Amino Acids and Proteins

• General structure: H2N–CHR–COOH

• Zwitterions at physiological pH; isoelectric point (pI)

• Proteins: Primary (sequence), secondary (α-helix, β-sheet), tertiary, quaternary structure

• Enzymes: Biological catalysts, specificity, affected by temperature, pH, inhibitors

Carbohydrates

• Monosaccharides (glucose), disaccharides (sucrose, lactose), polysaccharides (starch, glycogen, cellulose)

• D- and L- forms, Fischer and Haworth projections

• Energy source, structural roles, cell recognition

Lipids

• Fats (triglycerides), phospholipids, steroids

• Fatty acids: Saturated (no double bonds), unsaturated (one or more double bonds)

• Energy storage, membrane structure, signaling

Nucleic Acids

• DNA and RNA: Polymers of nucleotides (phosphate, sugar, base)

• Watson–Crick model: Double helix, base pairing (A–T, G–C)

• Genetic information storage, transmission, and expression

Nuclear Chemistry

Radioactivity and Nuclear Decay

• Spontaneous emission from unstable nuclei: alpha (α), beta (β), gamma (γ) decay

• Applications: Radiotherapy, imaging (PET scans), radiocarbon dating

Radioactive Half-Life

• Time for half of a radioactive sample to decay

• Formula: $N(t) = N_0 \times (1/2)^{t/t_{1/2}}$

• Medical relevance: Choosing isotopes for diagnosis/treatment

Reaction Rates and Chemical Equilibria

Endothermic and Exothermic Reactions

• Exothermic: Release energy (e.g., combustion)

• Endothermic: Absorb energy (e.g., photosynthesis)

Factors Affecting Reaction Rates

• Concentration, temperature, surface area, catalysts, nature of reactants

Chemical Equilibrium and Equilibrium Constants

• Dynamic state: Forward and reverse rates equal

• Equilibrium constant (K): $K = \frac{[\text{Products}]^c}{[\text{Reactants}]^a}$

• K >> 1: Products favored; K << 1: Reactants favored

Tables

Group	Representative Elements	Valence Electrons	Typical Ion Formed
1 (Alkali Metals)	Li, Na, K	1	+1
2 (Alkaline Earth Metals)	Mg, Ca	2	+2
17 (Halogens)	F, Cl, I	7	-1
18 (Noble Gases)	He, Ne, Ar	8 (2 for He)	0

Additional info: Table summarizes periodic group properties and typical ions formed.

Summary

This guide covers the foundational topics of general chemistry, including atomic structure, the periodic table, chemical bonding, reactions, stoichiometry, solutions, acids and bases, organic and biomolecules, nuclear chemistry, and reaction kinetics. Mastery of these concepts is essential for further studies in chemistry, biology, and medicine.