Comprehensive Study Notes for General Chemistry: Entrance Exam Topics

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 three primary subatomic particles:

• Protons: Positively charged, located in the nucleus.

• Neutrons: Neutral, also in the nucleus.

• Electrons: Negatively charged, orbiting the nucleus in electron shells or orbitals.

The arrangement of these particles determines the atom's chemical behavior.

Atomic Theory

Atomic theory has evolved from ancient concepts to modern quantum mechanics:

• Dalton's theory: All matter is composed of atoms; atoms of an element are identical; atoms combine in fixed ratios; chemical reactions rearrange atoms.

• Thomson: Discovery of electrons, "plum pudding" model.

• Rutherford: Discovery of the nucleus via gold foil experiment.

• Bohr: Quantized electron orbits.

• Modern quantum mechanics: Electrons exist in probabilistic 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.

• Neutral atoms: Number of protons = number of electrons.

• Ions: Atoms can gain or lose electrons, forming charged species.

Isotopes and Atomic Weight

Isotopes are atoms of the same element with different numbers of neutrons. This affects their mass but not their chemical properties.

• Example: Carbon-12, Carbon-13, Carbon-14.

• Atomic weight: Weighted average of isotopic masses, reflecting natural abundance.

Additional info: Isotopes are important in radiological imaging and dating techniques.

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).

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

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

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

• Group 17: Halogens (seven valence electrons, form -1 ions).

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

Electronic Structure and Electron Configurations

Electrons occupy shells and subshells (s, p, d, f) around the nucleus. The maximum number of electrons per shell is given by:

$2n^2$

• s-subshell: 2 electrons

• p-subshell: 6 electrons

• d-subshell: 10 electrons

• f-subshell: 14 electrons

Electron Configurations and the Periodic Table

Electron configurations determine chemical behavior. The periodic table is divided into blocks:

• s-block: Groups 1 and 2

• p-block: Groups 13–18

• d-block: Transition metals

• f-block: Lanthanides and actinides

Electron-Dot (Lewis) Symbols

Lewis symbols represent valence electrons as dots around the element symbol, useful for visualizing bonding.

Ionic Compounds and Acids/Bases

Ionic Compounds

Ionic compounds form by electron transfer, creating cations (positive) and anions (negative) held together by ionic bonds.

• Octet rule: Atoms seek eight valence electrons.

• Metals form cations; nonmetals form anions.

• Ionic bonds: Electrostatic attraction between oppositely charged ions.

Periodic Properties and Naming

• Group 1, 2, 13: Lose electrons to form cations.

• Group 15, 16, 17: Gain electrons to form anions.

• Naming: Cation first, anion with "-ide" suffix; transition metals specify charge with Roman numerals.

Properties of Ionic Compounds

• High melting/boiling points

• Crystalline structure

• Conduct electricity when molten or dissolved

• Soluble in water

Acids and Bases: H+ and OH- Ions

• Acids: Increase H+ concentration

• Bases: Increase OH- concentration

• pH: Balance of H+ and OH- ions

Molecular Compounds and Covalent Bonding

Molecular Compounds

Formed by covalent bonds (electron sharing), usually between nonmetals.

• Group 14: Four bonds (e.g., carbon)

• Group 15: Three bonds (e.g., nitrogen)

• Group 16: Two bonds (e.g., oxygen)

• Group 17: One bond (e.g., fluorine)

Characteristics

• Low melting/boiling points

• Poor electrical conductivity

• Diverse structures

Molecular Formulas and Lewis Structures

• Molecular formula: Exact atom count

• Lewis structure: Shows valence electrons and bonds

Polar Covalent Bonds and Electronegativity

• Polar bonds: Unequal sharing due to electronegativity differences

• Polar molecules: Asymmetric shape, dipole moment

Naming Binary Molecular Compounds

• First element: Full name

• Second element: Root + "-ide"

• Prefixes: mono-, di-, tri-, tetra-, etc.

Classification and Balancing of Chemical Reactions

Classes of Chemical Reactions

• Combination (Synthesis)

• Decomposition

• Single Replacement

• Double Replacement

• Combustion

• Neutralization

• Redox (Oxidation-Reduction)

Chemical Equations and Balancing

• Law of conservation of mass: Atoms must balance

• Steps: Write equation, count atoms, use coefficients, verify balance

Acids, Bases, and Neutralization

• Acids: Release H+

• Bases: Release OH-

• Neutralization: Acid + base → salt + water

Redox Reactions

• Oxidation: Loss of electrons

• Reduction: Gain of electrons

Mole and Mass Relationships

The Mole and Avogadro’s Number

The mole is a counting unit for atoms/molecules:

$1\text{ mole} = 6.022 \times 10^{23}$ entities

Gram–Mole Conversions

Relate mass to moles using molar mass:

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

$\text{Mass} = \text{Number of moles} \times \text{Molar mass}$

Reaction Rates and Chemical Equilibria

Endothermic and Exothermic Reactions

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

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

Factors Influencing Reaction Rates

• Concentration

• Temperature

• Surface area

• Catalysts

• Nature of reactants

Chemical Equilibrium

At equilibrium, forward and reverse reaction rates are equal.

Example: $\text{CO}_2 + \text{H}_2\text{O} \rightleftharpoons \text{H}_2\text{CO}_3$

Equilibrium Constants

$K = \frac{[\text{Products}]^c}{[\text{Reactants}]^a}$

• K >> 1: Products favored

• K << 1: Reactants favored

Nuclear Chemistry

Radioactivity

• Alpha decay: Emission of $^4\text{He}$

• Beta decay: Neutron → proton + electron

• Gamma decay: Emission of gamma rays

Radioactive Half-Life

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

• Half-life: Time for half the substance to decay

Physical Quantities and Metric System

Metric System of 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 there is a decimal point

Fundamental Chemical Laws

Law of Conservation of Mass

Mass is neither created nor destroyed in a chemical reaction.

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.

Chemical Calculations

Mole Concept and Chemical Formulas

Use the mole and molar mass for stoichiometric calculations.

Calculations Involving Chemical Equations

• Balance equation

• Convert quantities to moles

• Use mole ratios

• Convert moles to grams or other units

Calculations Involving Volume and Concentration

Molarity:

$\text{Molarity (M)} = \frac{\text{moles of solute}}{\text{volume of solution in liters}}$

Dilution:

$M_1V_1 = M_2V_2$

Solutions and Electrolytes

Mixtures and Solutions

• Homogeneous: Solutions

• Heterogeneous: Distinct phases

Units of Concentration

• Molarity (M)

• Mass/volume percent (% m/v)

• Parts per million (ppm)

• Molality (m)

Dilution

$M_1V_1 = M_2V_2$

Ions in Solution: Electrolytes

• Strong electrolytes: Complete dissociation

• Weak electrolytes: Partial dissociation

• Non-electrolytes: No dissociation

Acids and Bases

Acids and Bases in Aqueous Solution

• Acids: Release H+

• Bases: Release OH-

• pH scale: 0 (acidic) to 14 (basic), 7 neutral

Brønsted–Lowry Definition

• Acid: Proton donor

• Base: Proton acceptor

Acid Dissociation Constants and Strength

• Strong acids/bases: Complete dissociation

• Weak acids/bases: Partial dissociation

• Acid dissociation constant ($K_a$): Quantifies strength

Acid-Base Reactions and Salt Solutions

• Neutralization: Acid + base → salt + water

• Salt solutions can be neutral, acidic, or basic

Buffers and pH Measurement

Measuring Acidity: pH

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

• Acidic: pH < 7

• Neutral: pH = 7

• Basic: pH > 7

Buffer Solutions

• Mixture of weak acid and conjugate base (or weak base and conjugate acid)

• Resist pH changes

• Henderson–Hasselbalch equation:

$\text{pH} = \text{p}K_a + \log \left( \frac{[\text{A}^-]}{[\text{HA}]} \right)$

Organic Chemistry: Alkanes, Alkenes, Alkynes, Aromatics

Alkanes

• Saturated hydrocarbons: Only single bonds

• General formula: $\text{C}_n\text{H}_{2n+2}$

• Isomers: Same formula, different structure

• Tetrahedral geometry, 109.5° bond angles

• Naming: Longest chain, substituents, prefixes

• Properties: Nonpolar, low boiling/melting points, insoluble in water

• Reactions: Combustion, halogenation

Alkenes and Alkynes

• Alkenes: Double bonds, $\text{C}_n\text{H}_{2n}$

• Alkynes: Triple bonds, $\text{C}_n\text{H}_{2n-2}$

• Cis-trans isomerism in alkenes

• Properties: More reactive than alkanes

• Reactions: Addition (hydrogenation, halogenation, hydration)

Aromatic Compounds

• Benzene: Delocalized π-electrons, resonance

• Hückel’s Rule: 4n+2 π-electrons

• Naming: Substituents, ortho/meta/para positions

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

Alcohols, Phenols, Ethers, Thiols, Halides

Alcohols

• Contain -OH group

• Naming: Replace "-e" with "-ol"

• Properties: Polar, hydrogen bonding, high boiling points

• Acidity: Weak acids, phenols more acidic

• Reactions: Oxidation, dehydration, esterification

Phenols

• -OH group attached to aromatic ring

• More acidic than alcohols (resonance stabilization)

• Reactions: Form phenoxide salts with strong bases

Ethers

• R-O-R' structure

• Nonpolar, low boiling points, good solvents

• Can form peroxides

Thiols and Disulfides

• Thiols: R-SH, strong odor, weak acids

• Disulfides: R-S-S-R', important in protein structure

Halogen-Containing Compounds

• Alkyl/aryl halides: R-X

• Properties: Higher boiling points, polar bonds

• Reactions: Substitution, elimination

Amines, Aldehydes, Ketones

Amines

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

• Basicity: Lone pair on nitrogen accepts protons

• Form amine salts with acids

• Heterocyclic amines: Nitrogen in ring structure (e.g., pyridine)

Aldehydes

• Contain carbonyl group (C=O) bonded to hydrogen

• Naming: Replace "-e" with "-al"

• Properties: Intermediate boiling points, reactive

• Reactions: Oxidation to carboxylic acids, reduction to alcohols

Ketones

• Carbonyl group bonded to two alkyl/aryl groups

• Naming: Replace "-e" with "-one"

• Properties: Intermediate boiling points, soluble in water

• Reactions: Reduction to secondary alcohols

Carboxylic Acids and Derivatives

Carboxylic Acids

• Contain carboxyl group (-COOH)

• Naming: Replace "-e" with "-oic acid"

• Properties: High boiling points, soluble in water, weak acids

• Acidity influenced by substituents

Derivatives

• Esters: Carboxylic acid + alcohol

• Amides: Carboxylic acid + amine

• Anhydrides: Dehydration of two acids

Reactions

• Esterification, amide formation

• Hydrolysis of esters and amides

Amino Acids and Proteins

Amino Acids

• General structure: H2N–CHR–COOH

• Classified by side chain: nonpolar, polar, acidic, basic

• Essential vs. non-essential

• Chirality: L-form is biologically active

Acid–Base Properties

• Amphoteric: Can act as acid or base

• Zwitterion: Both positive and negative charges

• Isoelectric point (pI): pH where net charge is zero

Proteins

• Peptide bonds link amino acids

• Structure: Primary, secondary, tertiary, quaternary

• Denaturation: Loss of structure/function

• Enzymatic activity: Proteins as catalysts

Enzymes and Vitamins

Enzymes

• Biological catalysts, lower activation energy

• Highly specific for substrates

• Factors: Temperature, pH, substrate concentration

• Inhibition: Competitive, non-competitive

• Coenzymes (organic, from vitamins), cofactors (inorganic)

Vitamins and Minerals

• Vitamins: Water-soluble (B, C), fat-soluble (A, D, E, K)

• Minerals: Macrominerals (Ca, K, Na), trace elements (Fe, Zn, Se)

• Deficiency diseases: Scurvy, rickets, anemia

Carbohydrates

Classification

• Monosaccharides: Glucose, fructose

• Disaccharides: Sucrose, lactose, maltose

• Oligosaccharides: 3–10 units

• Polysaccharides: Starch, glycogen, cellulose, chitin

D and L Families, Structure of Glucose

• Chirality: D- and L- forms

• Fischer and Haworth projections

• Glucose: C6H12O6, pyranose ring, α/β anomers

Properties

• Solubility: Monosaccharides/disaccharides are water-soluble

• Reducing properties: Free aldehyde/ketone group

• Energy source: 4 kcal/g

• Biological functions: Structural, storage, communication

Lipids

Structure and Classification

• Simple lipids: Fats, waxes

• Complex lipids: Phospholipids, glycolipids

• Derived lipids: Steroids, fat-soluble vitamins

Fatty Acids and Esters

• Saturated: No double bonds

• Unsaturated: One or more double bonds

• Triglycerides: Ester of three fatty acids and glycerol

Properties of Fats and Oils

• Fats: Solid at room temperature (saturated)

• Oils: Liquid at room temperature (unsaturated)

• Hydrophobic, soluble in organic solvents

• Hydrolysis, saponification, hydrogenation

• Essential fatty acids: Must be obtained from diet

Nucleic Acids and Protein Synthesis

DNA, Chromosomes, and Genes

• DNA: Stores genetic information

• Chromosomes: DNA + proteins, 23 pairs in humans

• Genes: DNA sequences encoding proteins

Composition of Nucleic Acids

• Nucleotides: Phosphate, pentose sugar (deoxyribose/ribose), nitrogenous base

• Pyrimidines: Cytosine, thymine (DNA), uracil (RNA)

• Purines: Adenine, guanine

Structure and Base Pairing

• Phosphodiester bonds form backbone

• Watson–Crick model: Double helix, antiparallel strands

• Base pairing: A-T (2 H bonds), G-C (3 H bonds)

Additional info: Understanding nucleic acids is essential for molecular biology, genetics, and medical diagnostics.