2nd Semester CHEMISTRY Final Exam Review Guide 2014-2015
What will the test look like?
· 100 multiple choice problems
· All 2nd semester units, all material will be covered (see list below)
· 2 hour time period, no extra time. Some will finish early, some will use whole 2 hours
· Will go in test category, will be worth roughly 10% of overall grade
The units we have covered are listed below with the approximate percentage of questions on the final exam that will be from that unit:
Chapter Concept(s) Percentage of Q's
6 & 7 Ionic, Metallic, and Covalent Bonding 25
Formulas, Shapes, Polarity, Intermolecular Forces
End of 7 Empirical & Molecular Formulas 5
Percent Composition
8 & 9 Chemical Reactions 30-40
Stoichiometry
10 & 11 Gas Laws 25-30
Stoichiometry with Gases
· You will get all the charts, tables, and gas law equations and constants we have used in class so far: solubility table, activity series, electronegativity etc…
· What you were required to memorize, you will still be required to have memorized: polyatomic ions, common elements, etc....
· As with other tests, you will be allowed to use a scientific non-graphing calculator. There are a limited number of these you may borrow (not enough for everyone). Be prepared and bring one with you to class. You will also need #2 pencils for the scantron.
· You will not get extra time beyond the two hour time block reserved for this final. If you should finish your final early, please bring something with you to keep you occupied (a magazine, etc.)
· As usual, you may not have cell phones, ipods, headphones, etc… out during the test.
· You may not leave the room while you are taking the test. Therefore, go to the bathroom BEFORE the test starts. If you think you will get thirsty, bring a water bottle with you.
· If you know you will be absent, schedule a make-up time with me. Please make these arrangements as soon as possible.
Ch 6& 7: Bond Types, Properties, Structure of Molecules
Bond Types & Properties
· Be able to define and use the following terms: chemical bond, ion, molecule
· Know the 4 different bond types: ionic, metallic, network covalent, molecular covalent
· Given a chemical formula, identify what type of bonding is found in that substance.
· Compare and contrast the role of valence electrons in the 4 different types of bonds
· Identify what types of atoms are involved in the 4 bonding types (metals, non-metals, etc…)
· Describe the general properties of a certain type of bond (conducts electricity, dissolves, etc…)
· Given the properties of an unknown substance, identify type of bonding
Ionic Bonds
· What are ions? How do ions differ from neutral atoms?
· Predict ion charge for a given element. Relate ion charge to number of electrons gained or lost.
· Explain the relationship between ion charge, electron configurations, and the noble gases.
· Explain how cations and anions are formed.
· Predict the number of electrons an atom will gain or lose in an ionic bond
· Write a correct symbol for any ion (ex: calcium ion = Ca2+)
· Given a metal and non-metal combination, write the correct chemical formula for the ionic compound. (ex: when Ca and F combine, what compound will they make? Answer = CaF2)
· Given a chemical formula, correctly name any ionic compound
· Given a name, correctly write the chemical formula for any ionic compound
· Determine the charge of a metal cation based on its chemical formula (ex: what’s the charge of the iron atom in the compound FeO?)
· Be able to explain what a polyatomic ion is.
· Know the names, formulas, and charge of the following polyatomic ions:
o Sulfate
o Carbonate
o Nitrate
o Ammonium
o Hydroxide
o Phosphate
Covalent Bonds:.
· Write a Lewis dot symbol for a single atom
· Predict the number of covalent bonds an element will form. Explain how this is related to the atom’s Lewis dot structure.
· Write a Lewis dot structure for any molecule (using single, double, or triple bonds)
· Distinguish between bonded pairs and lone pairs in a Lewis Dot structure.
· Check Lewis Structures for correctness: make sure all electrons are paired, octet rule is satisfied, and each atom forms the correct number of bonds.
· Explain the role of electrons in determining shape. Explain what electron domains are, and how electron repulsion will create certain molecular shapes.
· Know the five molecular geometries we talked about in class:
o Linear
o Bent
o Trigonal planar
o Pyramidal
o Tetrahedral
· Explain what it means for a molecule to be “polar” or “non-polar”
· Explain what electronegativity is.
· What is the difference between a polar-covalent, non-polar covalent, and ionic bond? What happens to the electrons in each type of bond?
o Example: in a non-polar bond the electrons are shared equally between atoms, whereas in a polar covalent bond the electrons are shared unequally- one atom “pulls harder” on the electrons
· Be able to use the difference in electronegativity to predict which type of bond a pair of elements will form (polar, non-polar, or ionic)
· Identify whether a molecule is polar or non-polar, and be able to explain why.
· Identify dipole charges (which sided of the molecule of bond are negative and positive)
· Given a chemical formula or Lewis dot structure, identify (name and/or draw) the 3-dimensional shape a given molecule will have
· Know the 3 intermolecular forces and their relative strengths.
· Be able to predict which intermolecular forces are present in a given molecule.
· How do intermolecular forces (IMF) affect the properties of a molecule? Relate properties such as solubility and boiling point to IMFs such as dipoles and hydrogen bonds
· Predict how molecules will orient themselves based on polarity and IMF.
· Be able to name covalent molecules and write formulas based on names
End of Ch 7: Percent Composition, Empirical & Molecular Formulas
· Be able to find the percent composition of a compound
· Be able to use percent composition to determine the empirical formula of a compound
· Be able to use formula (molar) mass and percent composition to determine the molecular formula of a compound
Ch. 8 & 9: Reactions & Stoichiometry
Reactions:
§ Translate a chemical equation into words, and vice versa
§ Use appropriate symbols when writing a reaction (à, (s), etc…) and know what they represent
§ Explain what the following symbols represent: (s), (g), (l), and (aq).
§ Describe what a reaction will look like based on a chemical equation
o Describe what you will see- will you bubbles? Will you see a cloudy mixture form?
§ Identify reaction type, and/or predict the products of any of these reaction types:
o Single displacement
o Double displacement
o Synthesis
o Combustion
o Decomposition
§ Use coefficients to balance chemical equations
§ Know the 7 diatomic molecules (HOFBrINCl), and apply this knowledge appropriately when writing chemical equations
§ Use the activity series to determine whether a single displacement reaction will occur
o Solitary metal must be higher in activity than the one it is trying to displace
Stoichiometry:
§ Determine the molar mass of an element or compound using the periodic table
§ Use the coefficients in a balanced chemical equation to determine the mole ratios in a reaction.
§ Be able to do any stoichiometry calculation!
o grams ↔ grams, moles ↔ moles, moles ↔ grams, etc…
§ Given a chemical reaction and actual lab data, be able to calculate percent yield.
o Know terms: theoretical yield, actual yield, percent yield
Ch 10 & 11: Kinetic Molecular Theory and Gas Laws:
· Explain the Kinetic molecular theory of gases (know different properties of gases)
· What is the effect of temperature on gas particles?
· What is gas pressure? What is it caused by?
· What happens to pressure when you change volume? Using what you know about the motion of
o molecules, explain why this happens.
· How does changing temperature affect volume? How does changing temperature affect pressure? What causes these relationships? Explain by discussing the motion of gas molecules.
· Be able to convert between Celsius and Kelvin
· Be able to draw examples of barometers at different altitudes.
· What is absolute zero? Be able to theoretically define it in terms of gas particles.
· Be able to use these formulas to find the unknown pressure, volume, or temperature in different scenarios.
· Be able to explain a given phenomenon by discussing the movement of gas molecules and the relevant variables (P, T, V) that are involved.
· Be able to calculate volume, moles, and mass of gases using molar volume
· Be able to calculate temperature, pressure, volume, moles, and mass using the Ideal Gas Law: PV = nRT
· Be able to solve stoichiometry problems involving gases
Name:______Date:______Period:______
Semester 2 Final Exam Review Practice Problems
Ch 6 & 7:
1. How many electrons are lost or gained in forming each ion?
a. Mg2+ b. Br- c. Ag+ d. Fe3+
2. Classify each of the following as a covalent compound or an ionic compound.
a. CO2 b. NaCl c. MgCl2 d. N2 e. H2O
3. What types of elements tend to combine to form covalent compounds? What exactly is a covalent bond?
4. What types of elements tend to combine to form ionic compounds? What exactly is an ionic bond?
5. Will solid copper sulfate conduct electricity? Will aqueous copper sulfate conduct electricity? Explain.
6. Which type of bond results in a “sea of electrons”? Why do we call it that?
7. What type of bonding do you expect there to be
in the following molecules?
a. NH3 (g)
b. NaCl (s)
c. Mg (s)
d. H2O (l)
e. a hard solid which does not dissolve in water and does not conduct electricity
f. a substance which conducts electricity as a solid.
8. What charge of ion will each of the following elements form? Write the symbol for the ion they will form.
a. Potassium b. Magnesium c. Bromine
9. Write the formula (including charge) for each ion.
a. carbonate ion ______c. sulfate ion ______e. phosphate ion ______
b. nitrate ion ______d. hydroxide ion ______f. ammonium ion ______
10. Explain why transition metals need roman numerals in their compound names.
11. List 3 elements that will combine with Oxygen in a 1:1 ratio. Explain why this is.
12. Write the formulas for these ionic compounds.
a. magnesium oxide______d. sodium sulfide ______
b. potassium iodide______e. aluminum chloride ______
c. tin(II) fluoride ______f. Iron (II) nitride ______
g. Ba2+, Cl- ______h. Ca2+, S2- ______i. Al3+, O2- ______
j. Ag+, I- ______k. K+, Br- ______l. Fe2+, O2-______
13. Name the following ionic compounds.
a. MnO2 ______e. SrBr2 ______
b. CaCl2 ______f. K2S ______
c. NiCl2 ______g. CuCl2______
14. Write formulas for the following ionic compounds.
a. sodium phosphate ______d. potassium nitrate ______
b. sodium hydroxide ______e. ammonium chloride ______
c. magnesium sulfate ______f. potassium carbonate______
15. Write formulas for compounds formed from these pairs of ions.
a. NH4+, SO42- ______c. barium ion and hydroxide ion ______
b. K+, NO3- ______d. lithium ion and carbonate ion ______
16. Name the following compounds.
a. Na2CO3 ______d. FeCl3______
b. Li2SO4 ______e. K2CO3______
c. Cu(OH)2 ______f. LiNO3______
17. Draw the Lewis dot structure for nitrogen trichloride.
18. Draw the electron dot configuration for acetylene, C2H2.
19. How many bonding pairs of electrons are in a Lewis dot diagram of PH3?
(A) 1 (B) 2 (C) 3 (D) 4
20. In a single bond, the atoms share
(A) 1 electron (B) 2 electrons (C) 3 electrons (D) 4 electrons
21. If the electronegativity difference between two atoms is extremely large, what type of bond will they form? What if the electronegativity difference is very small?
22. If the bond in the molecule HI is polar, on which end of the hydrogen iodide molecule would you find a partial negative charge?
23. Describe how you determine the shape of a molecule. What do “electron domains” have to do with it? How and why do the “electron domains” affect shape?
24. Draw a structural formula for C4H10O and C2H4. Describe the shape around each central atom. Are the molecules polar or nonpolar?
25. Draw & name the shape for carbon dioxide, CO2. Is this molecule polar or nonpolar?
26. Draw & name the shape of fluorine monoxide, F2O. Is this molecule polar or nonpolar?
27. What type of bond; nonpolar covalent, polar covalent, or ionic will form between each pair of atoms?
a. Na and O b. O and O c. P and O
28. What is the strongest Intermolecular force acting between the following molecules?
a. CO2 ______
b. CH4O ______
c. CH2O ______
d. NF3 ______
29. Which of the molecules from question 28 would evaporate the fastest? The slowest?
30. Which of the following molecules would have the highest boiling point? Circle and explain why.
31. For the molecules below, answer each of the following questions:
A) Draw the Lewis dot structure
B) Name the shape of the molecule
C) Identify the molecule as polar or non-polar
D) Identify the intermolecular force present