AP Chemistry Quarter 1 Review

Answer and outline of separate paper.

  1. Stoichiometry and Reactions
  • Memorize solubility rules and polyatomic ions.
  • Use the stock system and drop and swap to write chemical formulas and equations.
  • Memorize the rules for significant figures.
  • Understand scientific notation and the metric system.
  • Memorize the percent error formula.
  • Calculate mole conversions using gram formula mass, 22.4L and Avogadro’s number.
  • Understand and calculate empirical and molecular formulas.
  • Balance and identify the type of reactions (S, D, SR, DR, C)
  • Calculate limiters and excess reactants with percent yield.
  • Calculate Molarity and Dilutions.
  1. Name and determine the solubility: NaNO3 KI MgSO4 CaClO3 PbS FeCl3
  2. Give formula & solubility: Strontium sulfite, ammonium nitrite, Aluminum chlorite, Chromium (III) oxide.
  3. How many significant figures: 2536 20.300.02260
  4. Calculate and report to the correct number of significant figures: 1.52-.265= _____ 5.2*6.98= ______
  5. Put into scientific notation: 562000000.00000258
  6. Take out of scientific notation: 5.26x109 6.25x10-5
  7. Write a number line outlining the metric system from Pico to Tera.
  8. Convert: 5.6 mL to L, 6.9 m to nm, 596 kg to g.
  9. A student finds the density of water to be 1.09g/mL. Find the percent error.
  10. Define: conservation of mass, qualitative, quantitative, extensive, intensive, solute, solvent, solution.
  11. Calculate the number of moles, liters and molecules of 65g of oxygen at STP.
  12. A compound with an empirical formula of CH4 has a mass of 64. What is the molecular formula?
  13. Find the empirical formula: N2H3C6H12O6 HCl.
  14. Find the molecular formula of a compound of nitrogen and oxygen that weighs150g/mol and has 46.7% Nitrogen.
  15. Balance and ID the type of reaction:AlPO4 + Mg(NO3) 2 Mg3(PO4)2 + Al(NO3)3

Na + N2 NaN3

H2O  H2 + O2

C3H8 + O2 CO2 + H2O

Na + KI  K + NaI

  1. Describe how to prepare a 100ml and 2M solution of HCl from a 5M solution with the equipment you’d use.
  2. Calculate the volume of 2M Phosphoric acid needed to react to react with 3.5g of Al.
  3. Calculate the mass of calcium acetate formed when 2.34 grams of calcium reacts with 0.345 moles of acetic acid.
  1. Atomics
  • Memorize theories, experiments and discoveries of Dalton, Thompson, Rutherford, Bohr, and QMM.
  • Understand the subatomic particles, atomic number, mass number, isotopes, ions, atomic mass calculations.
  • Calculate wavelength, frequency, and Energy of light.Explain ground versus excited states.
  • Write and explain spdf configurations and notations for atoms and ions with exceptions.
  • Explain the laws of Heisenberg, Hund, Pauli, and Aufbau.
  • Draw orbital shapes, explain nodes, and identify isoelectric pairs and para/diamagnetism.
  • Identify and explain the Mass spec and PES diagrams.
  1. Isotopes have the same number of ______but different number of ______.
  2. Write out the configuration and valence orbital notation for: C, Cu, Mo, N3-, Cr+2.
  3. When an electron is excited and falls down to the ground state, energy is ______in the form of ______. Which releases the most energy: n=1n=3 , n=2n=5, n=6n=2 or n=5n=4?
  4. Write a configuration for Boron in the excited state. How is it different than the ground state?
  5. Describe the nature of electrons in terms of waves and particles.
  6. Calculate the frequency and energy of a photon of light with a wavelength of 750nm. Convert energy in J to KJ/mol.
  7. Calculate the wavelength (in nm) and energy of a photon of light with a velocity of 2.50x1014s-1. Use your notes to explain what part of the light spectrum this photon would be in (you’d be given a light spectrum diagram on the exam if asked).
  8. Sketch the mass spec diagram for Oxygen knowing the isotopes that exist are O-16 and O-18.
  9. Sketch the PES diagram for Phosphorus. Explain peak size and location on the x-axis.
  1. Periodic Table
  • Explain the Periodic law, periods, group names/properties (alkali metals, alkaline earth metals, halogens, noble gases)
  • Explain the metal, metalloid, nonmetal, transition metals- properties and locations
  • Identify and explain atomic and ionic radii, ionization energy, electro-negativity, definitions and patterns
  • Explain covalent, metallic, and ionic bonds (octet rule),
  • Calculate and explain Lattice energy, heat of reactions, PE diagrams, bond strength and length.
  1. Elements in the same ______have similar properties and are also known as ______.
  2. Describe metals, metalloids and nonmetals in terms of properties, IE, EA, ions, and conductivity.
  3. Describe the general trend associated with IE, Radius, Electronegativity and Electron Affinity.
  4. Write equations to describe the electron affinity of Fluorine and the IE of Sodium.
  5. What is the name given to elements that form colored solutions and compounds?
  6. List the monatomic elements. Why don’t these elements react?
  7. Describe ionic, covalent and metallic bonds in terms of conductivity, strength, electrons, mp, and bp.
  8. Draw a picture to show how NaCl is dissolved in water based on polarity.
  9. Define lattice energy. Which has more lattice energy: NaCl or Li2O? Explain.
  10. Calculate the heat of formation of the reaction between CH4 and oxygen. (Use a textbook).
  1. Bonding
  • Draw Lewis structures, provide geometric shapes and polarity with exceptions (linear, bent, pyramidal, planar, tetrahedral, trigonal bipyramidal, T shaped, see saw, octahedral, square pyramidal, square planar), angles, polarity
  • Explain and depict Resonance, formal charge, hybridizations, sigma and pi bonds, bond order.
  • Explain properties and examples of IMF (dipole-dipole, ion-dipole, H Bonds, network solids, LDF).
  1. Draw Lewis structures for all resonance structures of the following and calculate formal charge to find the best structure: CO2, NF3, CH4, OF2, PF5, AsF3, SF6, XeF4, PF4-, SF3-, IF5, I3-
  2. Define the octet rule. Which elements can exceed the octet? Which can fall under? Why?
  3. Using the structures above, what are their geometries? What are the central atom’s hybridizations? Show any dipole moments. Give the bond angles.
  4. Draw 1,2 pentadiene and count the sigma and pi bonds. Which bonds are strongest: sigma or pi? Explain.
  5. List the diatomic molecules. Which have a bond order of 2? Which has the shortest bond length? Explain.
  6. Give properties and examples of the following IMF: Dipole-Dipole, LDF, Hydrogen bonds, Ion-Dipole, Network Solids. List the forces in order from weakest to strongest.
  7. Which are stronger: bonds or forces? Which break when a chemical change occurs? Which are involved when physical changes occur?
  8. Explain in terms of bond or forces why water evaporates when heated.

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