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First Unit Review

Go to the Answer Key

  1. Basics
  2. What is an element?
  3. What is a compound?
  4. State the difference between a physical change and a chemical change.
  5. Review your sig figs and how to use them in calculations.
  6. a) A substance was found to have a mass of 35.94 g and a volume of 3.18 mL. What is its density?
    b) A sample of aluminum (density = 2.78 g/mL) was found to have a mass of 72.98 grams. What is its volume?
  7. Make the following conversions: a) 380 seconds to hours   b) 2500 mL/min to L/second

    The Atom

  8. State the four principles of Dalton's atomic theory.
  9. List the three main subatomic particles. Give their symbol, charge, approximate relative mass, and location in the atom.
  10. Describe Rutherford's experiment, his results, and conclusions he drew from them.
  11. What is an isotope?
  12. Describe J.J. Thomson's experiment and his results.
  13. Fill in the following table. (It includes ions, which are in a later chapter, but might as well cover it now.)
    Element or Ion Name Element or Ion Symbol Atomic Number Atomic Mass # of p+ # of e- # of no charge

    calcium ion





    sulfide ion





  14. The four isotopes of lead are 122Pb with an abundance of 1.37%, 124Pb with an abundance of 26.26%, 125Pb in 20.82% abundance, and 126Pb in 51.55% abundance. What is the atomic mass of lead?

    Atomic Spectra

  15. What was the difference between the atom as described by Thomson, Rutherford, Bohr, and the quantum mechanical model?
  16. List the four types of sublevels that an electron can occupy. State how many orbitals are in each sublevel and how many electrons can occupy that sublevel.
  17. State the Aufbau principle, the Pauli exclusion principle, and Hund's rule.
  18. Write electron configurations for the following elements:
    a) Na   b) Kr   c) Ta
  19. List the colors of the spectrum in order of increasing wavelength.
  20. What is the wavelength, in nm, and the energy of light whose frequency is 4.14 x 1014 Hz?
  21. Explain why we see lines in an atomic spectrum. A diagram may be useful.

    Chemical Formulas

  22. Identify each of the following as metals, nonmetals, or metalloids.
    a) sodium   b) carbon   c) radium   d) silicon   e) xenon
  23. Explain how calcium forms a +2 cation. Name the ion.
  24. Explain how phosphorus forms a -3 anion. Name the ion.
  25. What is the difference between a molecular compound and an ionic compound?
  27. Write formulas for the following compounds. State whether they are molecular or ionic.
    a) calcium iodide   b) phosphorus pentachloride   c) ammonium sulfide
    d) copper (I) hydrogen phosphite   e) dinitrogen pentoxide   f) tin (IV) oxalate
  28. Name the following compounds:
    a) KMnO4   b) SnCr2O7   c) S2O3   d) NaNO3   e) Pb(HCO3)4

    Mole Conversions

  29. How many molecules in 42.99 grams of carbon dioxide?
  30. How many grams are in one mole of calcium hydrogen carbonate?
  31. How many grams in 4.39 moles of magnesium nitrate?
  32. How many liters are in 2.3 grams of sulfur dioxide gas at STP?
  33. Determine the number of moles in 182.49 L of hydrogen gas at STP.
  34. What is the molar mass of a gas which has a density of 2.86 g/L at STP?
  35. How many atoms of O are in 2.33 grams of nitrogen dioxide?
  36. What is the percent carbon by mass in sucrose (C12H22O11)?
  37. A compound is composed of 58.8% C, 9.8% H, and 31.4% O.
    a) Find its empirical formula.
    b) What is its molecular formula if its molar mass is 204 g/mol?

    Chemical Reactions

  38. List the five types of reactions and briefly describe each one.
  39. Why must an equation be balanced?
  40. Write and balance the following equations (finish if needed). Identify each type.
    a) carbon and oxygen react to form carbon monoxide
    b) potassium chlorate decomposes to form potassium chloride and oxygen
    c) iron (III) chloride reacts with calcium hydroxide
    d) magnesium reacts with silver nitrate
    e) butane (C4H10) burns in air
    f) potassium reacts with oxygen
  41. Write a balanced net ionic equation for each reaction.
    a) lead (II) nitrate reacts with sodium chloride (lead (II) chloride is a precipitate)
    b) zinc is placed in hydrochloric acid solution

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