AP Chemistry
Advanced Placement Chemistry aligned with the College Board CED: atomic structure and properties, molecular and ionic compound structure, intermolecular forces, chemical reactions, kinetics, thermodynamics, equilibrium, acids and bases, and applications of thermodynamics including electrochemistry.
Ämne: Kemi · Nivå: Gymnasium (16–19) · 399 kort
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- A mole is 6.022×10²³ entities (Avogadro's number). One mole of any substance contains the same number of particles as 12 g of pure ¹²C. The mole bridges the atomic and macroscopic scales.
- Molar mass is the mass of one mole of a substance in g/mol. Numerically it equals the atomic mass (for elements) or the sum of atomic masses (for compounds) from the periodic table. Example: H₂O = 2(1.008) + 16.00 = 18.02 g/mol.
- Atoms have a tiny dense nucleus (protons + neutrons) surrounded by electrons in a much larger volume. Protons carry +1 charge, electrons −1, neutrons 0. The atomic number Z (protons) defines the element; mass number A = protons + neutrons.
- Isotopes are atoms of the same element (same Z) with different numbers of neutrons. Example: ¹²C and ¹³C both have 6 protons but 6 and 7 neutrons. Isotopes have nearly identical chemistry but different masses.
- Mass spectrometry ionizes atoms/molecules, accelerates them through electric and magnetic fields, and separates them by mass-to-charge ratio (m/z). Peak heights give relative abundance; multiplying mass×abundance and summing gives the average atomic mass.
- Electromagnetic radiation has wave properties (λ × ν = c) and particle properties (E = hν). c = 3.00×10⁸ m/s; h (Planck constant) = 6.626×10⁻³⁴ J·s. Shorter wavelength = higher frequency = higher energy photons.
- The photoelectric effect: light striking a metal can eject electrons only if its frequency exceeds a threshold, regardless of intensity. This proved light is quantized into photons of energy E = hν. Einstein received the 1921 Nobel Prize for this explanation.
- The Bohr model describes electrons in quantized energy levels around a nucleus. Energy differences between levels equal photon energies absorbed or emitted (ΔE = hν). Although superseded by quantum mechanics, it correctly predicts hydrogen's spectrum.
- Electrons occupy orbitals described by four quantum numbers: n (principal, shell, 1,2,3...), ℓ (angular, subshell shape: s=0, p=1, d=2, f=3), mℓ (orientation), and mₛ (spin: +½ or −½). Each orbital holds at most 2 electrons (Pauli exclusion).
- Subshells fill in order of increasing energy (Aufbau principle): 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p. The 4s fills before 3d because of penetration and shielding effects.
- Hund's rule: when filling degenerate orbitals (same energy, like the three 2p orbitals), electrons enter each orbital singly with parallel spins before pairing. This minimizes electron-electron repulsion.
- Noble-gas core notation abbreviates filled inner shells. Example: Fe is [Ar] 4s² 3d⁶. Valence electrons (outermost shell) drive chemistry; for main-group elements they equal the group number.
- Coulomb's law: the force between charges is F ∝ (q₁·q₂)/r². Higher charge or shorter distance means stronger attraction (opposite signs) or repulsion (same signs). It underlies ionization energies, ionic-bond strength, and lattice energy.
- Effective nuclear charge (Z_eff) is the net positive charge an electron actually experiences — the full nuclear charge minus shielding by inner electrons. Z_eff increases left to right across a period as protons add without much extra shielding.
- Atomic radius decreases left to right across a period (higher Z_eff pulls electrons in) and increases top to bottom down a group (new shells add). Cations are smaller than their parent atoms; anions are larger.
- First ionization energy (IE₁) is the energy to remove one electron from a gaseous atom. It increases left-to-right across a period and decreases down a group. Successive IEs always rise; huge jumps occur when removing a core electron.
- Electron affinity is the energy change when a gaseous atom gains an electron. Halogens have the most negative (most favorable) values because adding an electron completes a noble-gas configuration. Noble gases and group 2 have positive (unfavorable) values.
- Electronegativity (Pauling scale) measures an atom's ability to attract bonding electrons. It increases left-to-right and bottom-to-top; F is the most electronegative element (3.98), Cs and Fr the least. Electronegativity differences predict bond polarity.
- Photoelectron spectroscopy (PES) measures binding energies of electrons in different subshells. Peak position = subshell energy; peak height = number of electrons in that subshell. PES spectra confirm the shell/subshell model directly.
- Ionic bonds form when a metal donates electrons to a nonmetal, producing cations and anions held together by electrostatic attraction. Ionic compounds form crystalline lattices with high melting points and conduct electricity only when molten or dissolved.
- Lattice energy is the energy released when gaseous ions form one mole of solid ionic compound. By Coulomb's law it grows with ion charge and shrinks with ion size: MgO (±², small) has much higher lattice energy than NaCl (±¹, larger).
- Covalent bonds form between nonmetals that share electrons. Single bonds share one pair, double bonds two pairs, triple bonds three pairs. As bond order rises, bonds become shorter and stronger.
- Bond polarity arises from electronegativity differences (ΔEN). ΔEN ≈ 0 → nonpolar covalent; 0 < ΔEN < ~1.7 → polar covalent (partial charges δ⁺ and δ⁻); ΔEN > ~1.7 → mostly ionic. These cutoffs are guidelines, not sharp boundaries.
- Lewis structures show valence electrons as dots/lines. Steps: count total valence e⁻, draw skeleton with central atom (least electronegative, not H), connect with single bonds, distribute lone pairs to satisfy octets, convert lone pairs to multiple bonds if needed.
- Formal charge = (valence e⁻) − (lone-pair e⁻) − ½(bonding e⁻). The best Lewis structure minimizes formal charges, places negative formal charge on the most electronegative atom, and has the sum of formal charges equal to the overall charge.
- Resonance occurs when two or more equivalent Lewis structures can be drawn that differ only in electron placement. The real molecule is a hybrid; bond orders are averaged. Examples: ozone (O₃), carbonate (CO₃²⁻), benzene (C₆H₆).
- Octet exceptions: H and He are stable with 2 valence electrons. B and Be can have fewer than 8 (BF₃ has 6 on B). Period-3+ central atoms (P, S, Cl, Xe) can have expanded octets (PCl₅, SF₆, XeF₄).
- VSEPR theory predicts molecular shape from electron-pair repulsion around a central atom. Count steric number = bonded atoms + lone pairs. Lone pairs occupy more space than bonding pairs, slightly compressing the angles between bonded atoms.
- Common VSEPR shapes: 2 = linear (180°); 3 = trigonal planar (120°) or bent (with 1 lone pair); 4 = tetrahedral (109.5°), trigonal pyramidal (1 LP), or bent (2 LP); 5 = trigonal bipyramidal; 6 = octahedral (90°).
- Molecular polarity depends on both bond polarity and geometry. CO₂ has polar C=O bonds but is linear, so dipoles cancel — nonpolar overall. H₂O has polar O–H bonds and is bent (104.5°), so dipoles add — strongly polar.
- Hybridization mixes atomic orbitals to form equivalent hybrid orbitals matching the molecule's geometry: 2 regions → sp (linear); 3 → sp² (trigonal planar); 4 → sp³ (tetrahedral); 5 → sp³d; 6 → sp³d².
- Sigma (σ) bonds form by head-on overlap of orbitals along the bond axis. Pi (π) bonds form by side-on overlap of unhybridized p orbitals above and below the axis. Single = 1σ; double = 1σ + 1π; triple = 1σ + 2π.
- Metallic bonding: metal cations sit in a sea of delocalized valence electrons. This explains high electrical and thermal conductivity, malleability, ductility, and luster. Alloys are mixtures of metals or metals with nonmetals that retain metallic properties.
- Network covalent solids (diamond, graphite, SiO₂, silicon) consist of atoms covalently bonded in extended 3D or 2D networks. They have very high melting points and great hardness; most do not conduct electricity (graphite is an exception, with delocalized π electrons).
- Intermolecular forces (IMFs) are attractive forces between molecules — not within them. IMFs determine melting point, boiling point, vapor pressure, viscosity, and surface tension. They are much weaker than covalent or ionic bonds.
- London dispersion forces (LDFs) arise from instantaneous, temporary dipoles in any molecule's electron cloud. They occur in all molecules but are the only IMF between nonpolar molecules. Strength grows with molecular size (more electrons → more polarizable).
- Dipole-dipole forces operate between polar molecules whose permanent dipoles align δ⁺ to δ⁻. They are stronger than LDFs of comparable size but weaker than hydrogen bonds. Example: HCl molecules in the liquid state.
- Hydrogen bonding is an unusually strong dipole-dipole force when H is covalently bonded to N, O, or F. The small H, large ΔEN, and lone pairs on the acceptor create attractive forces strong enough to give water its anomalously high boiling point.
- Ion-dipole forces hold ions to polar molecules — the basis of ionic compounds dissolving in water. Ion-induced dipole and dipole-induced dipole forces help nonpolar gases dissolve slightly in polar solvents.
- States of matter differ by particle motion and IMFs. Solid: fixed positions, vibrating; strong IMFs. Liquid: tumbling but in contact; moderate. Gas: free, far apart; negligible IMFs. Adding heat overcomes IMFs and changes state.
- Ideal gas law: PV = nRT, with P in atm or kPa, V in L, n in mol, T in K. R = 0.0821 L·atm/(mol·K) or 8.314 J/(mol·K). Real gases deviate at high pressure and low temperature where molecular volume and IMFs become significant.
- Kinetic molecular theory assumes: gas particles are tiny and far apart; they move in random straight lines; collisions are elastic; no IMFs; average kinetic energy is proportional to T (in K). KE_avg = (3/2)kT per molecule.
- Dalton's law: in a mixture of gases, P_total = P₁ + P₂ + .... Partial pressure of gas i: P_i = x_i × P_total, where x_i is the mole fraction. Useful when collecting gas over water (subtract water vapor pressure).
- Maxwell-Boltzmann distribution shows the spread of speeds among gas particles at a given T. Higher T shifts the curve right (faster, broader) and flattens the peak. At any T some fraction of molecules has enough energy to react (relates to activation energy).
- Solutions: a homogeneous mixture of solute(s) in a solvent. Concentration units: molarity (M = mol solute / L solution), molality (m = mol solute / kg solvent), mole fraction (x), mass percent. Molarity changes with T (volume expands); molality does not.
- 'Like dissolves like': polar solvents dissolve polar and ionic solutes; nonpolar solvents dissolve nonpolar solutes. Dissolving requires that solute-solvent IMFs replace solute-solute and solvent-solvent IMFs with comparable strength.
- Chromatography separates components by differential affinity for a stationary phase vs. a mobile phase. Paper, thin-layer, column, and gas chromatography all rely on this. R_f = distance traveled by spot / distance traveled by solvent front.
- Distillation separates liquids by boiling point: the lower-boiling component vaporizes first, condenses, and is collected. Fractional distillation uses a packed column to refine separations of liquids with similar boiling points.
- Beer-Lambert law: A = εbc, where A is absorbance (unitless), ε is molar absorptivity (L/(mol·cm)), b is path length (cm), and c is concentration (M). A plot of A vs. c at fixed λ is linear and used to determine unknown concentrations.
- UV-Vis spectroscopy probes electronic transitions; infrared (IR) spectroscopy probes molecular vibrations (each bond type absorbs at a characteristic wavenumber); microwave spectroscopy probes rotations; NMR probes nuclear spin in magnetic fields.