South 고등학교 3학년 Chemistry I
Chapters: 4
1. First steps in chemistry
Chemistry and human problems · Useful carbon compounds · Avogadro's number and moles · Chemical equations · Molarity
- Chemistry and Society – Chemistry studies how atoms join and rearrange. By choosing which atoms to join and how, chemists make new materials, medicines, fertilisers and fuels. Atoms are never lost in a reaction, so every product has a cost and a waste that we must plan for.
- Carbon and Its Compounds: Bonding, Hydrocarbons and Naming – Carbon has 4 outer electrons, so it shares electrons (covalent bonds) instead of gaining or losing them. Because it bonds to itself (catenation) and always makes 4 bonds (tetravalency), it forms millions of compounds: chains, branches and rings, saturated or unsaturated. Compounds with the same functional group form a homologous series that differs by –CH₂–, and IUPAC names are built from the number of carbons + a suffix or prefix for the functional group.
- Mole Concept, Molar Mass and Chemical Formulas – Atoms are far too small to count one by one, so chemists count them by weighing. Atomic masses are given in u, where 1 u is one-twelfth the mass of a carbon-12 atom. One mole is 6.022 × 10²³ particles, and its mass in grams equals the formula mass in u. With moles we can find the percentage of each element in a compound, and work back from percentages to the empirical and molecular formulas.
- Stoichiometry, Limiting Reagent and Concentration Terms – Stoichiometry means measuring the amounts of substances in a reaction. A balanced equation acts like a recipe: its numbers give the mole ratio of reactants and products. With it you can change any mass into moles, use the ratio, and change back to the mass of any other substance. The reactant that runs out first is the limiting reagent; it decides how much product forms. For solutions, concentration is given as mass percent, mole fraction, molarity or molality.
- Solutions: Types, Concentration and Henry's Law – A solution is an even mix of two or more substances. The part in bigger amount is the solvent, the smaller part is the solute. We tell 'how strong' a solution is with concentration terms: mass %, volume %, ppm, mole fraction, molarity (per litre of solution) and molality (per kg of solvent). Solids usually dissolve more when hot. Gases dissolve more when their pressure is high (Henry's law, p = KH·x) and less when it is hot.
2. The atomic world
Atomic structure and isotopes · Quantum numbers and orbitals · Electron configuration rules · Development of the periodic table · Periodic trends
- Subatomic Particles and Early Atomic Models – An atom is made of three main particles. Electrons (charge −1, very light) were found with cathode ray tubes. Protons (charge +1, about 1836 times heavier) were found in canal rays. Neutrons (no charge, mass about a proton) were found by Chadwick in 1932. Thomson pictured the atom as a ball of positive charge with electrons stuck in it. Rutherford fired alpha particles at gold foil and found a tiny, heavy, positive nucleus with electrons moving outside. Atomic number Z = number of protons; mass number A = protons + neutrons. Rutherford's model could not explain why atoms are stable or why they give line spectra, which led to Bohr's model.
- Quantum Mechanical Model: Quantum Numbers, Orbitals and Configuration – Moving electrons also behave like waves: λ = h/mv (de Broglie). Because of this we cannot know exact position and momentum together: Δx·Δp ≥ h/4π (Heisenberg). So we drop sharp orbits and use orbitals: regions where the electron is likely to be, from the Schrödinger equation (ψ² gives probability). Four quantum numbers describe each electron: n (shell, size, energy), l (subshell, shape: 0 to n−1), mₗ (orientation: −l to +l) and mₛ (spin: +½ or −½). s is spherical, p is dumbbell, d is mostly four-lobed. Electrons fill by the Aufbau (n + l) rule, Pauli principle (max 2, opposite spins) and Hund's rule (single first). Half-filled and fully filled subshells are extra stable, so Cr is 3d⁵4s¹ and Cu is 3d¹⁰4s¹.
- Periodic Classification of Elements – Scientists sorted elements so that similar ones sit together. Döbereiner made triads, Newlands found that every 8th element repeats (octaves), Mendeleev arranged elements by atomic mass and left gaps for unknown ones. Moseley showed that atomic number is the real key. Modern periodic law: the properties of elements are a periodic function of their atomic numbers. The modern table has 7 periods and 18 groups. Elements with Z > 100 get temporary IUPAC names built from digit roots (nil, un, bi, tri, quad, pent, hex, sept, oct, enn) plus -ium.
- Periodic Trends in Properties – Two forces decide almost every trend: the pull of the nucleus (effective nuclear charge) and the distance of the outer shell. Across a period the nuclear pull grows while the shell stays the same, so atoms shrink, ionisation enthalpy rises, electron gain enthalpy becomes more negative and electronegativity rises. Down a group a new shell is added, so atoms grow and these values fall. Cations are smaller and anions bigger than their atoms. Valence follows the outer electrons; metallic reactivity is highest at the bottom left and non-metallic reactivity at the top right.
3. Bonding and molecules
Electrical nature of bonds · Ionic bonding · Covalent and metallic bonding · Electronegativity · Lewis structures · VSEPR shapes · Properties from shape
- Kossel-Lewis Approach and the Ionic Bond – Atoms join so that each gets a stable outer shell of 8 electrons (an octet), like a noble gas. Kossel said atoms can give or take electrons to make ions (ionic bond). Lewis said atoms can also share pairs of electrons (covalent bond). Lewis structures show these electrons as dots and lines. Formal charge (V − L − B/2) helps pick the best Lewis structure. Ions pack into a crystal, and the energy released is linked to the lattice enthalpy.
- Bond Parameters, Resonance and Polarity – A covalent bond is described by four numbers: bond length (distance between nuclei), bond angle (angle between bonds at an atom), bond enthalpy (energy to break 1 mol of bonds) and bond order (number of shared pairs). Higher bond order means a shorter, stronger bond. When one Lewis structure cannot describe a molecule, the real molecule is a resonance hybrid of several structures. Unequal sharing makes a bond polar; the dipole moment μ = q × d measures it, and the shape decides whether bond dipoles cancel.
- VSEPR Theory: Shapes of Molecules – VSEPR stands for Valence Shell Electron Pair Repulsion. The electron pairs around a central atom repel each other and move as far apart as possible. The number of pairs sets the basic arrangement: 2 linear, 3 trigonal planar, 4 tetrahedral, 5 trigonal bipyramidal, 6 octahedral. Lone pairs take more room than bond pairs (lp–lp > lp–bp > bp–bp), so they squeeze bond angles and change the shape we see, for example NH₃ (pyramidal, 107°) and H₂O (bent, 104.5°).
- Intermolecular Forces – Intermolecular forces are weak attractions BETWEEN molecules. From weakest to strongest (for similar-size molecules): London dispersion < dipole–dipole < hydrogen bond. Stronger forces mean higher melting and boiling points, higher viscosity and surface tension, and they decide what dissolves in what.
4. Dynamic reactions
Dynamic equilibrium · Autoionisation of water and pH · Neutralisation · Titration · Oxidation and reduction · Heat in reactions
- Chemical Equilibrium: Kc, Kp, Q and Gibbs Energy – In a closed container a reversible reaction goes both ways. After some time the forward and backward rates become equal, so amounts stop changing, but the reaction does not stop. This is dynamic equilibrium. At equilibrium the ratio of products to reactants (each raised to its coefficient) is a fixed number, the equilibrium constant K. Kc uses concentrations, Kp uses partial pressures, and Kp = Kc(RT)^Δn. Pure solids and liquids are left out of K. The reaction quotient Q tells the direction: Q < K goes forward, Q > K goes backward, Q = K is equilibrium. K and Gibbs energy are linked: ΔG = ΔG° + RT ln Q and ΔG° = −RT ln K.
- Ionic Equilibrium: Acids, Bases, Ionisation and pH – Acids and bases can be defined in three ways: Arrhenius (give H⁺ or OH⁻ in water), Brønsted–Lowry (proton donor or acceptor) and Lewis (electron-pair acceptor or donor). Strong acids and bases ionise fully; weak ones ionise only a little and set up an equilibrium with constant Ka or Kb. Water itself ionises: Kw = [H⁺][OH⁻] = 1.0 × 10⁻¹⁴ at 25 °C. pH = −log[H⁺] and pH + pOH = 14. For a weak electrolyte, α ≈ √(K/C), so dilution increases α (Ostwald's dilution law). For a conjugate pair, Ka × Kb = Kw. Salts of weak acids or weak bases react with water (hydrolysis) and give non-neutral solutions.
- Titration: Finding an Unknown Concentration – Titration finds the unknown concentration of a solution. A solution of known concentration (the titrant) is added slowly from a burette to a measured volume of the unknown (the analyte) until the reaction is just complete. An indicator or a meter shows this point. At the equivalence point the moles react in the ratio of the equation, so for a 1 : 1 reaction C₁V₁ = C₂V₂.
- Redox Reactions: From Oxygen to Electron Transfer – Oxidation first meant adding oxygen or removing hydrogen. Reduction meant the opposite. Today we use a bigger idea: oxidation is losing electrons and reduction is gaining electrons. Both always happen together, so we call them redox reactions. A more active metal gives electrons to the ion of a less active metal.
- Enthalpy, Calorimetry and Hess's Law – Enthalpy H = U + pV. At constant pressure the heat taken in or given out equals ΔH; at constant volume it equals ΔU. We measure ΔU in a sealed bomb calorimeter and ΔH in an open cup calorimeter, using q = C ΔT or q = m c ΔT. For gases, ΔH = ΔU + Δn_g RT. The reaction enthalpy ΔrH is negative for exothermic and positive for endothermic reactions, and ΔrH° = ΣΔfH°(products) − ΣΔfH°(reactants). Hess's law: the total enthalpy change is the same whether a reaction happens in one step or many, so thermochemical equations can be added like algebra.