01:160:161 · Flashcards

01:160:161 General Chemistry Review: Stoichiometry, Gases, Atomic Structure, Bonding and Thermochemistry

30 termsRutgers Universityby @studyboardUpdated October 2026

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A study set written by the Studyboard team for the topics in this course's published outline. It isn't official course material from Rutgers University and doesn't contain real exam questions.

Terms in this set (30)

  1. What is a mole?A mole is 6.022 x 10^23 particles (Avogadro's number). The molar mass in grams per mole equals the formula mass in atomic mass units.
  2. Mini-problem: how many moles are in 36.0 g of water (H2O, 18.0 g/mol)?36.0 g / 18.0 g/mol = 2.00 mol.
  3. Limiting reactantThe reactant that is completely used up first; it determines the maximum amount of product that can form (the theoretical yield).
  4. Percent yield formulaPercent yield = (actual yield / theoretical yield) x 100.
  5. Molarity definitionM = moles of solute / liters of solution.
  6. Dilution equationM1V1 = M2V2: moles of solute stay the same when a solution is diluted.
  7. Empirical vs molecular formulaThe empirical formula is the simplest whole-number ratio of atoms (CH2O for glucose); the molecular formula is the actual number of each atom (C6H12O6).
  8. Oxidation vs reductionOxidation is loss of electrons (oxidation number increases); reduction is gain of electrons (oxidation number decreases). Mnemonic: OIL RIG.
  9. Ideal gas law with symbolsPV = nRT: P pressure, V volume, n moles, R gas constant (0.08206 L atm / mol K), T temperature in kelvin.
  10. Pitfall in gas law problemsTemperature must be in kelvin (K = degrees C + 273.15); using Celsius gives wrong answers.
  11. Dalton's law of partial pressuresThe total pressure of a gas mixture equals the sum of the partial pressures of each gas; partial pressure of gas A = mole fraction of A x total pressure.
  12. Kinetic molecular theory: what does average kinetic energy depend on?Only absolute temperature. At the same temperature, lighter gas molecules move faster on average than heavier ones.
  13. Graham's law of effusionrate1 / rate2 = square root of (M2 / M1): lighter gases effuse faster.
  14. Relationship between wavelength, frequency and photon energyc = wavelength x frequency, and E = h x frequency, so shorter wavelength means higher frequency and higher energy photons.
  15. Four quantum numbersn (principal: energy level and size), l (angular momentum: subshell shape, 0 to n - 1), ml (magnetic: orbital orientation) and ms (spin, +1/2 or -1/2).
  16. Pauli exclusion principle and Hund's rulePauli: no two electrons in an atom have the same four quantum numbers, so an orbital holds at most two electrons of opposite spin. Hund: electrons fill degenerate orbitals singly before pairing.
  17. Electron configuration of oxygen (Z = 8)1s2 2s2 2p4.
  18. Periodic trends in atomic radiusAtomic radius decreases across a period (greater effective nuclear charge) and increases down a group (more electron shells).
  19. Periodic trends in ionization energy and electronegativityBoth generally increase across a period and decrease down a group. Fluorine is the most electronegative element.
  20. Ionic vs covalent bondsIonic bonds form from electron transfer and electrostatic attraction between ions (often metal + nonmetal). Covalent bonds form from shared electron pairs (usually between nonmetals).
  21. Formal charge formulaFormal charge = valence electrons - nonbonding electrons - (bonding electrons / 2). The best Lewis structure minimizes formal charges.
  22. VSEPR: geometry of a molecule with 4 electron domains and 2 lone pairs (like H2O)Electron geometry is tetrahedral, molecular geometry is bent, with a bond angle slightly less than 109.5 degrees (about 104.5 for water).
  23. Hybridization for 2, 3 and 4 electron domains2 domains: sp (linear). 3 domains: sp2 (trigonal planar). 4 domains: sp3 (tetrahedral).
  24. Sigma vs pi bondsA single bond is one sigma bond; a double bond is one sigma and one pi; a triple bond is one sigma and two pi. Pi bonds come from side-by-side p orbital overlap.
  25. How can a molecule with polar bonds be nonpolar? Give an example.If its geometry is symmetric, the bond dipoles cancel. CO2 is linear with two polar C=O bonds but no net dipole.
  26. Rank the intermolecular forces from weakest to strongestLondon dispersion forces < dipole-dipole forces < hydrogen bonding (H bonded to N, O or F). Ion-dipole forces are stronger still.
  27. Endothermic vs exothermic reactionsExothermic reactions release heat (delta H < 0); endothermic reactions absorb heat (delta H > 0).
  28. Heat equation q = m c delta Tq is heat, m is mass, c is specific heat capacity and delta T is the temperature change. For water, c = 4.184 J/g C.
  29. Hess's lawThe enthalpy change of a reaction is the same regardless of path, so delta H values of steps can be added; reversing a reaction flips the sign and multiplying it scales delta H.
  30. Enthalpy of reaction from enthalpies of formationDelta H rxn = sum of n x delta Hf(products) - sum of n x delta Hf(reactants). Elements in their standard states have delta Hf = 0.

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