CHEM 1A Exam Review: Stoichiometry, Atomic Structure, Bonding, Gases and Equilibrium

Terms in this set (35)

1. What is a mole?The amount of substance containing Avogadro's number of particles, 6.022 x 10^23. The molar mass in g/mol equals the mass of one mole of that substance.
2. What is a limiting reactant?The reactant that is used up first, so it determines the maximum amount of product. Find it by converting each reactant to moles of product and taking the smallest.
3. Percent yield formulaPercent yield = (actual yield / theoretical yield) x 100%. Theoretical yield comes from stoichiometry using the limiting reactant.
4. Mini-problem: How many moles are in 36.0 g of water (M = 18.0 g/mol)?n = m / M = 36.0 g / 18.0 g/mol = 2.00 mol.
5. Relationship between energy, frequency and wavelength of lightE = h*nu and c = lambda*nu, so E = h*c/lambda. h = 6.626 x 10^-34 J*s, c = 3.00 x 10^8 m/s. Shorter wavelength means higher energy.
6. Bohr energy levels for hydrogenE_n = -2.18 x 10^-18 J / n^2 (about -13.6 eV / n^2). A photon is emitted when an electron drops from higher n to lower n; its energy equals the difference.
7. de Broglie wavelengthlambda = h / (m*v). Any moving particle has wave character; it matters for tiny masses like electrons.
8. Heisenberg uncertainty principlePosition and momentum cannot both be known exactly: delta_x * delta_p >= h/(4*pi). This is why electrons are described by orbitals (probability distributions), not orbits.
9. The four quantum numbers and their allowed valuesn = 1, 2, 3...; l = 0 to n-1 (s, p, d, f); m_l = -l to +l; m_s = +1/2 or -1/2. No two electrons in an atom share all four (Pauli exclusion).
10. Hund's ruleIn degenerate orbitals, electrons occupy separate orbitals with parallel spins before pairing up, which minimizes electron repulsion.
11. Electron configuration of Fe (Z = 26)1s2 2s2 2p6 3s2 3p6 4s2 3d6, or [Ar] 4s2 3d6.
12. Periodic trends: atomic radius and ionization energyAtomic radius decreases across a period (higher effective nuclear charge) and increases down a group. Ionization energy shows the opposite trend.
13. What is effective nuclear charge (Z_eff)?The net positive charge felt by an electron after shielding by other electrons, roughly Z minus the number of core electrons. It rises across a period.
14. Ionic vs covalent bondingIonic bonds form by electron transfer between atoms with a large electronegativity difference (often metal plus nonmetal). Covalent bonds share electron pairs, usually between nonmetals.
15. How do you calculate formal charge?Formal charge = valence electrons - nonbonding electrons - (bonding electrons / 2). The best Lewis structure keeps formal charges near zero, with negative charge on the more electronegative atom.
16. VSEPR: geometry of a central atom with 4 electron domains and 1 lone pairElectron geometry is tetrahedral; molecular geometry is trigonal pyramidal (example: NH3), bond angle a bit less than 109.5 degrees.
17. What makes a molecule polar?It has polar bonds arranged so their dipoles do not cancel. CO2 is nonpolar (linear, dipoles cancel) while H2O is polar (bent).
18. Hybridization for 2, 3 and 4 electron domains2 domains: sp (linear); 3 domains: sp2 (trigonal planar); 4 domains: sp3 (tetrahedral).
19. Sigma vs pi bondsA single bond is one sigma bond; a double bond is one sigma plus one pi; a triple bond is one sigma plus two pi. Pi bonds form from side-on overlap of unhybridized p orbitals.
20. Ideal gas law and its symbolsPV = nRT. P = pressure, V = volume, n = moles, T = temperature in kelvin, R = 0.08206 L*atm/(mol*K) or 8.314 J/(mol*K).
21. When do real gases deviate most from ideal behavior?At high pressure (molecule volume matters) and low temperature (attractions matter). The van der Waals equation corrects with the a (attraction) and b (volume) terms.
22. Dalton's law of partial pressuresP_total = P_1 + P_2 + ...; each partial pressure P_i = x_i * P_total, where x_i is the mole fraction.
23. Hess's lawThe enthalpy change for a reaction is the same whatever path is taken, so delta H values of steps can be added. Reversing a reaction flips the sign; multiplying it scales delta H.
24. Standard enthalpy of reaction from enthalpies of formationdelta H_rxn = sum(n * delta H_f products) - sum(n * delta H_f reactants). delta H_f of an element in its standard state is zero.
25. Gibbs free energy and spontaneitydelta G = delta H - T*delta S. A reaction is spontaneous when delta G < 0. With delta H < 0 and delta S > 0 it is spontaneous at all temperatures.
26. Link between delta G standard and Kdelta G standard = -R*T*ln K. If K > 1, delta G standard is negative and products are favored at equilibrium.
27. Reaction quotient Q vs KIf Q < K the reaction shifts toward products; if Q > K it shifts toward reactants; if Q = K it is at equilibrium.
28. Le Chatelier's principle: effect of adding heat to an exothermic reactionHeat acts like a product, so the equilibrium shifts toward reactants and K decreases. Temperature is the only change that alters the value of K.
29. pH, pOH and Kw at 25 CpH = -log[H3O+], pOH = -log[OH-], pH + pOH = 14, and Kw = [H3O+][OH-] = 1.0 x 10^-14.
30. Henderson-Hasselbalch equationpH = pKa + log([A-]/[HA]). It applies to buffers; when [A-] = [HA], pH = pKa.
31. Solubility product Ksp: write it for CaF2CaF2(s) <=> Ca2+ + 2F-, so Ksp = [Ca2+][F-]^2. If molar solubility is s, Ksp = s(2s)^2 = 4s^3.
32. Oxidation vs reductionOxidation is loss of electrons (oxidation number rises); reduction is gain of electrons (oxidation number falls). The oxidizing agent is the species that gets reduced.
33. Rate law and reaction orderRate = k[A]^m[B]^n, where m and n are found from experiment, not from coefficients. The overall order is m + n.
34. First-order half-lifet_1/2 = 0.693 / k, independent of starting concentration. Integrated law: ln[A] = ln[A]0 - k*t.
35. Arrhenius equationk = A * e^(-Ea/(R*T)). Raising T or lowering activation energy Ea (as a catalyst does) increases the rate constant.