Download Chemistry Midterm Review 2006

Chemistry Midterm Exam Review 2013
These are topics from a traditional 1st quarter. If you used the thematic approach with me
this year, you will notice that the topics are not in order of their presentation throughout
the year. The final exam is cumulative so this review will help you with reviewing the
concepts and calculations. Notice that I have connected chapters to each unit from your
book. Don’t forget the importance of the review sheet.
Unit 1: Matter and Change (R= Ch 1&2 H= Ch 1 & 3)
1. Define chemistry.
2. State the difference between quantitative and qualitative data.
3. a. Define matter. b. Name the three states and list 2-3 characteristics for each of
them.
4. What is the difference between a physical and chemical property? Give examples
of each.
5. A chemical change is also known as a chemical _______________.
6. Name 5 buzz words that signify a physical change and 5 that signify a chemical
change.
7. a. Define pure substances and state examples. b. Define element and compound &
state examples.
8. a. By what means can you separate a compound? Give some examples. b. By
what means can you separate a mixture? Give some examples.
9. What is the difference between a homogeneous and heterogeneous mixture? List
some examples of each.
10. State whether each is a compound or element: Fe, CO, CaCl2, Hg, Co, argon,
sodium chloride, I2.
11. Write the symbols for the following. mercury, gold, iodine, calcium, barium, tin,
magnesium, phosphorus.
12. Name five indicators(observations) of a chemical reaction.
13. Define the words “reactant” and “product”. In a chemical equation, where are the
reactants located? Where are the products located? What separates them from
each other?
14. Classify each as a physical or chemical change: a. food spoiling b. water boils c.
nail rusting d. baking bread e. sugar dissolving in water f. tarnishing silver
Unit 1: Measurement (R= Ch 3 H= Ch 2)
1. The metric system is based on the power of ____________
2. What is the difference between accuracy and precision?
Honors Only:
3. Define significant figures. What “rules” are used to count sig figs?
4. How many significant figures are in each: a. 5.730 x 108 b. 3000 c. 0.01552 d.
9009 e. 629.55 f. 1.777 x 10-3?
5. Round the first 4 values above to 1 sig fig and the last 2 values above to 2 sig figs.
6. What is the answer expressed in proper significant figures for the following: a.
6.54 + 3.053 and b. 8.95 x .02?
7. a. Place .0178 in standard scientific notation. b. Put 8.931 x 104 into ordinary
notation.
8. What is the basic metric unit of length, mass, and volume?
9. What is the SI unit for volume? 1cm3 = 1 ______
10. a. Define volume. b. What type of lab equipment measures approximate volume?
11. a. Define density. b. What is the equation?
12. Ice floats because it is more or less dense than water?
13. Put the 3 phases of matter in order of increasing density.
14. A copper penny has a mass of 3.1 g and a volume of .35 cm3. What is the density?
15. A plastic ball has a volume of 19.7 cm3 and a density of .8029 g/cm3. What is the
mass?
16. The density of silicon is 2.33 g/cm3. What is the volume if its mass is 62.9g?
17. a. Convert 157 cg into g. b. Convert 8.6 kg into cg. C. Convert 100m into cm.
d. Convert 70.68 cl into Hl.
18. What are the 6 common prefixes for the metric units and what are their values?
19. Define temperature. What are the 3 temperature scales and what is the boiling and
freezing point of each scale?
20. a. Convert 25 ºC into K; b. Convert 35 ºC into ºF.
21. a. Sketch what a graph will look like that shows a directly proportional
relationship? b. What is an example of a directly proportional relationship that we
have talked about this year?
22. a. Sketch what a graph will look like that shows an inversely proportional
relationship? b. What is an example of an inversely proportional relationship that
we have talked about this year?
Unit 2: Atomic Structure (R= Ch 4 H= Ch 4)
1. Define an “atom”.
2. Describe the law of conservation of mass.
3. List the assumptions of Dalton’s atomic theory. What were the revisions?
4. What are the 3 subatomic particles and state their charge and where they are
found in the atom?
5. Who is Rutherford and what did he discover about the atom?
6. a. What is the dense center of an atom called? b. What subatomic particles are
found in the center?
7. Where are the electrons found in an atom?
8. Define atomic number and atomic mass.
9. Define average atomic mass.
Honors Only
10. Boron-10 has a mass of 10.013 amu and a % abundance of 19.8. Boron -11’s
abundance is 80.2% with a mass of 11.009 amu. Calculate the average atomic
mass.
11. What does “electrically neutral” mean in terms of the atom?
12. How many protons and electrons are in a carbon atom?
13. How many neutrons in beryllium?
14. a. Define an isotope. b. Write the nuclear symbol for nitrogen-15. c. Write the
hyphen notation for 38Cl.
15. How many protons, electrons, and neutrons are in oxygen-16?
16. Determine the number of neutrons in 226Ra and 15N.
17. Fill in the chart below.
Atomic #
Mass #
protons
electrons
neutrons
7
7
9
10
39
19
59
27
18. How many protons and electrons are in the following? a. Na+1 b. N3- c. F-1
d. Al3+?
Unit 3: Electron Arrangement (R= Ch 5 H= Ch 5)
1. Transition elements are in the __________ block and the inner transitions are in
the ________ block.
2. Group 1A and 2A are in the ___________ block and groups 3A to 8A are in the
__________ block.
3. Write the electron configuration for the following: a. boron b. magnesium, c.
vanadium,d. strontium, e. iron f. arsenic
4. a. What is an atomic orbital? b. What shape is the s sublevel? c. The shape of the
p sublevel? d. What are the maximum number of electrons allowed in each
sublevel?
5. What is the difference between the Bohr model and the Quantum mechanical
model?
6. a. What are flame tests? b. What area of the electromagnetic radiation spectrum
allows us to observe flame tests? c. Is energy released or absorbed when an
electron falls from a higher energy level to a lower energy level?
7. What is the difference between a ground state and an excited state?
8. What is the lowest energy level? The lowest sublevel?
9. What is the maximum number of electrons in the 4th energy level?
10. State the three rules for filling atomic orbitals with electrons and describe them.
11. How many unpaired electrons are in the following: a. boron b. fluorine?
12. How many valence electrons are in the highest energy level (valence)? a. barium
b. sodium c. aluminum d. oxygen.
13. What is the symbol of the following configurations?a. 1s2 2s2 2p6 3s1 b. 1s2 2s2
2p6 3s2 3p6 4s2 3d2 c. 1s2 2s2 2p6 3s2 3p2 d. [Kr] 5s2 4d10.
14. Write the shorthand configurations for a. barium b. aluminum c. arsenic.
15. a. Define wavelength and frequency. B. What kind of relationship do they have?
c. What kind of relationship does the energy and frequency of a wave have?
Unit 3: Periodic Table (R= Ch 6 H= Ch 6-7)
1. a. Define the periodic table. b. What is the difference between a group and a
period? c. What do columns of elements have in common?
2. Group A elements are called ___________. Group B elements are called
__________.
3. a. What side of the periodic table are the metals? b. The nonmetals? c. Where are
the metalloids?
4. Name 4 characteristics of both metals and nonmetals.
5. Define malleable and ductile.
6. Identify each as a metal, nonmetal, or metalloid: K, B, Mo, iodine, uranium, and
aluminum.
7. a. Who is Mendeleev? b. Who is Moseley? c. How is the modern periodic table
arranged?
8. Name the a. group 1A metals b. 2A metals c. 7A nonmetals d. 8A nonmetals.
9. State 3-4 properties of each of the families above.
10. Which family is the most stable?
11. Which family reacts vigorously with water?
12. Which family is extracted from mineral ores?
13. Which family are the most reactive metals?
14. Which family of nonmetals combines with 1A and 2Ametals to make salts?
15. a. What is electronegativity? b. What is the period and group trend? c. Which one
has a higher electronegativity; C, N, or K?
16. a. Define ionization energy. b. What is the period and group trend? c. Which has
a higher ionization energy; Na, K, Mg, or P?
17. a. Define atomic radius. b. What is the period and group trend? c. Which has a
higher atomic radius; C, N, Mg, P, Na, or K?
18. a. What is the period and group trend for the ionic size of cations? Of anions?
b. How does the size of a neutral atom compare with the cation and the anion?
Unit 4: Ionic & Covalent Bonding (R= Ch 7 & 9 H= Ch 8-9)
1. a. What is a valence electron? b. How many valence electrons are in potassium
and oxygen?
2. What is the shorthand notation for a. P b. Sr c. I?
3. Draw the Lewis structure for a. Mg b. Si c. Cl.
4. State the octet rule. The duet rule.
5. a. Define “ion”. b. What is the difference between a cation and an anion?
6. What is a chemical formula?
7. What is the chemical formula for the following; a. sulfide ion b. sodium ion
c. fluoride ion d. mercury (II) ion?
8. What are the names of the following ions; a. Ba2+ b. Al3+ c. O2- d. Sn4+?
9. Metals form _______ions and nonmetals form ________ ions.
10. a. What is the difference between ionic and covalent bonds? b. How does
electronegativity difference determine bond type?
11. Write the electron configurations for a. Al+3 b. O-2 c. Ti+2.
12. Draw Lewis structures for the following; a. K2O b. MgCl2 c. KI d. Na3P
13. Which of the following compounds are ionic? A. H2O b. Na2O c. CO2
d. CaS, e. SO2 f. CaCO3.
14. What is the difference between a nonpolar covalent bond and a polar covalent
bond?
15. What are the properties of a covalent and ionic compound in terms of state of
matter, ability to conduct electricity, hardness & solubility in water?
16. a. Define chemical bond.
b. What is a lone pair of electrons?
17. What is the difference between a single, double, and triple bond?
18. Draw Lewis diagrams for a. PBr3 b. N2 c. CF4 d. HBr e. SO2
19. a. What is a dipole? b. What direction does it travel?
20. What is the difference between a polar and nonpolar molecule?
21. Using the formulas from question 16, list ones which are polar and nonpolar.
22. Describe the model for a metallic bond.
23. Using the model described above, explain why a metal is a great conductor of
electricity.
Unit 5: The MOLE
1. Calculate molar mass of a) C6H12O6
b) Ca3(PO4)2
2. Convert from moles into grams, moles into liters of gas, moles into molecules, atoms,
or formula units, and liters to liters
a) Convert 100.0 grams of HCl into moles.
b) Find the mass in grams of 36.5 moles of S.
c) How many moles of Ca are in 5.43 x 1023 atoms of Ca?
d) How many moles of O2 at STP are in 34.5 L?
e) What is the mass in grams of 1.20 x 108 atoms of copper?
3. Calculate percent composition of
a) nitrogen in NH3
b) both elements in BaCl2
Honors Only
4. Calculate the empirical formula and molecular formulas.
a) What is the empirical formula of a compound that is 25.9% nitrogen and 74.1%
oxygen?
b) What is the empirical formula of a compound that has a mass of 10.150 grams
and contains 4.433 grams of P and 5.717 g of O?
c) What is the molecular formula if the empirical formula is Na2O and the gram
formula mass is 62 g?
5. Which of the following is a molecular formula of XY3?
X2Y3, XY4, X2Y5, X2Y6
6. Know the representative particle for an ionic compound, covalent compound,
element, and a diatomic molecule?
Unit 6: Naming Compounds/ Formula Writing (R= Ch 8
1. Define formula unit and define molecule.
H= Ch 8-9)
2. Name the following chemical formulas; a. KCl b. BaSO4 c. MgBr2 d. Li2CO3
e. CoF2 f. NaBr
3. Write formulas for the following: a. potassium nitrate b. lithium oxide c.
calcium phosphate d. ammonium carbonate e. barium chloride.
4. Prefixes are used to indicate the number of atoms in _________ compounds.
_________ are used to indicate the number of atoms in a compound.
5. _________ are used to indicate the charge of transition elements in a compound.
6. a. What is a polyatomic ion? b. What are the names of the following ions: PO4-3,
SO3-2, CO3-2, NO3-1, NH4+1
7. What are binary compounds? Ternary compounds?
8. Name the following covalent compounds: a. CO b. N2O5 c. SO3
9. What are the chemical formulas for the seven diatomic molecules?