Download Chemistry Mid-Term Review: 2015-2016

Chemistry Mid-Term Review: 2015-2016
This review is worth 5% extra credit to your mid-term exam. It is due the day you are scheduled to take your midterm. Your work must be hand written. Be sure to answer every question. For calculation questions, show your
work.
Chapter 1:
1. What are the steps of the scientific method?
2. Explain the difference between the variable and control in an experiment.
3. Why should a hypothesis be developed before experiments take place?
4. What is the difference between a theory and scientific law?
5. What is the difference between a manipulated variable and a responding variable?
6. Explain how the results of many experiments lead to the development of a theory and scientific law?
Chapter 2:
1. Compare the properties of solids, liquids, and gases.
2. Amongst solids, liquids, and gases, which have a definite volume?
3. Amongst solids, liquids, and gases, which take the shape of their container?
4. Amongst solids, liquids, and gases, which can be compressed?
5. How many phases exist in homogeneous and heterogeneous mixtures?
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6. How is a compound different from an element?
7. How can you distinguish a substance from a mixture?
8. In a chemical reaction, how does the mass of the reactants compare with the mass of the products?
9. What is the difference between chemical and physical properties? Give examples.
10. What is the main difference between physical and chemical changes?
11. Classify each of the following as physical or chemical changes:
a. Water boiling
d. A metal rusts
b. Milk turns sour
e. Wood burns
c. Salt dissolves in water
f. ice melts
12. What are the differences between extensive and intensive properties?
13. Is a solution considered homogeneous or heterogeneous? How many phases does a solution have?
14. Explain the process of distillation and describe the goal of distillation.
15. How can you distinguish between elements and compounds?
16. Explain the compositions of the following substances by looking at their chemical formulas:
a. H2O
c. C6H12O6
b. NH3
d. CH4
17. What is the difference between chemical and physical properties? Give examples.
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18. What are the properties of solids, liquids, and gases? Which have a definite volume? Which take the shape
of their container?
Chapter 3:
1. Count the significant figures in each length
a. 0.05730 meters
c. 8765 meters
b. 0.00073 meters
d. 40.007 meters
2. How many significant figures are in each measurement?
a. 143 grams
c. 0.074 meters
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b. 8.750 x 10 grams
d. 1.072 meters
3. Round each measurement to three significant figures. Write your answers in scientific notation.
a. 87.073
b. 4.3621 x 108 meters
c. 0.01552 meter
d. 9009 meters
e. 1.7777 x 10-3 meters
f. 629.55 meters
4. Perform each operation. Express your answers to the correct number of significant figures.
a. 61.2 meters + 9.35 meters + 8.6 meters
b. 9.44 meters - 2.11 meters
c. 1.36 meters + 10.17 meters
d. 34.61 meters – 17.3 meters
5. Solve each problem. Give your answers to the correct number of significant figures and in scientific
notation.
a. 8.3 meters x 2.22 meters
b. 8432 meters / 12.5
c. 35.2 seconds x (1 minute / 60 seconds)
6. Solve the following and express each answer in scientific notation and to the correct number of significant
figures.
a. (5.3 x 104) + (1.3 x 104)
b. (7.2 x 10-4) / (1.8 x 103)
c. 104 x 10-3 x 106
d. (9.12 x 10-1) – (4.7 x 10-2)
e. (5.4 x 104) x (3.5 x 109)
7. Name the quantity measured by each of the seven SI base units and give the SI symbol of the unit.
8. State the difference between mass and weight.
9. State the relationship between degrees Celsius and kelvins.
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10. Using tables from this chapter, convert the following.
a. 0.044 km to meters
b. 4.6 mg to grams
c. 0.107 g to centigrams
11. Covert the following
a. 15 cm3 to liters
b. 6.7 s to milliseconds
c. 94.5 g to micrograms
12. A 68-g bar of gold is cut into 3 equal pieces. How does the density of each piece compare to the density of
the original gold bar?
13. Round off each of these measurements to three significant figures.
a. 98.473 L
b. 0.00076321 cg
c. 57.048 m
d. 12.17 C
14. What equation is used to determine the density of an object?
15. What is the difference between accuracy and precision?
16. Which is the estimated digit when using significant figures?
17. Using the density formula, how do you solve for mass, density, and volume?
18. What is the difference between the accepted and experimental value?
19. How do you use conversion factors?
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20. What is dimensional analysis?
21. What is the order and values of the metric prefixes from largest to smallest?
Chapter 4:
1. How did Democritus characterize atoms?
2. According to Dalton’s Theory is it possible to convert atoms of one element into atoms of another? Explain!
3. What are the charges and relative masses of the three main subatomic particles?
4. Three isotopes of oxygen are oxygen-16, oxygen-17 and oxygen 18. Write the symbol for each including the
atomic number and mass number.
a.
c.
b.
5. What distinguishes the atoms of one element from the atoms of another?
6. What equation tells you how to calculate the number of neutrons in an atom?
7. What is the charge- positive or negative, of the nucleus of every atom?
8. Why is an atom electrically neutral?
9. What does the atomic number of each atom represent?
10. How are the elements on the modern periodic table arranged?
11. What are the parts of an atom? What are their charges? Where are they located?
12. What is the difference between mass number and atomic mass?
13. Which subatomic particle is mainly responsible for the properties of each element?
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14. Which subatomic particle makes each atom unique?
Chapter 5
1. How do two sublevels of the same principle energy level differ from each other?
2. How can electrons in an atom move from one energy level to another?
3. How many orbitals are in the following sublevels?
a. 3p sublevel- _______
c. 4p sublevel- _______ e. 4f sublevel- _______
b. 2s sublevel- _______
d. 3d sublevel- _______
4. What is an atomic orbital?
5. How many sublevels are contained in each of these principle energy levels?
a. 1st- _____
b. 2nd- _____
c. 3rd- _____
d. 4th- _____
6. How many electrons are in the highest occupied energy level of these atoms?
a. Barium -________
b. Sodium-________
c. Aluminum-________
d. Oxygen-________
7. What are the three rules that govern the filling of the atomic orbitals by electrons?
a.
b.
c.
8. Give electron configurations for atoms of these elements:
a. Na
b. S
c. Mg
d. Ne
e. K
9. What is the maximum number of electrons that can go into each of the following sublevels?
a. 2s- ______
c. 3p- ______
e. 4s- ______
g. 3d- ______
b. 4p- ______
d. 5s- ______
f. 4f- ______
h. 5p- ______
10. Give the symbol for the atom that corresponds to the following electron configurations:
a. 1s22s22p63s23p6 ________
b. 1s22s22p63s23p63d104s24p64d75s1 ________
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c. 1s22s22p63s23p63d104s24p64d104f145s25p65d16s2 _______
11. An atom of an element has two electrons in the first energy level and five electrons in the second energy
level. Write the electron configuration for this atom and name the element. How many unpaired electrons
does an atom of this element have?
a. Electron Configurationb. Namec. Unpaired electrons12. What is the difference between a cation and an anion?
13. How many electrons fit into each principal quantum number? (also known as an energy level.)
14. How many orbitals will the s, p, d, & f sublevels hold? How many electrons in each orbital?
15. Compare Hund’s rule, the Aufbau principle, and the Pauli Exclusion Principle.
16. Is it more stable for an atom to have 3 partially filled orbitals or 2 filled and 1 empty orbital?
Chapter 6
1. How are the elements arranged on the modern periodic table?
2. Identify each element as a metal, nonmetal or metalloid:
a. Gold-
c. Silicon-
b. Sulfur-
d. Barium-
3. Into what four classes can element be sorted based on their electron configurations?
4. Classify each element as a represented elements, transition metal or noble gas:
a. 1s22s22p63s23p63d104s24p6- ______________
b. 1s22s22p63s23p63d104s2- ______________
c. 1s22s22p63s23p2- ______________
5. How does atomic size change within groups and across the periods?
6. When do ions form?
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7. Compare the size of ions to the size of the atom from which they form.
8. How does electronegativity vary within groups and across periods?
9. Identify each property below as a characteristic of either a metal or nonmetal
a. A gas at room tempc. Malleableb. Brittle-
d. Shiny-
d. Poor conductor of electric current-
10. Where are the alkaline metals, the alkaline earth metals, the halogens and the noble gases located on the
periodic table?
11. Which noble gas does not have eight electrons in its highest occupied energy level?
12. Explain the difference between the first and second ionization energy of an element.
13. For groups 1A – 7A, how many electrons fill the outermost sublevel?
14. How are the electron configurations for the atoms in each group of representative elements similar?
15. Where are the noble gases located on the periodic table? Why are these elements noble?
16. What is the difference between ionization energy and electronegativity?
17. Why does atomic size decrease across a period?
18. What are the group and period trends for atomic size, ionization energy, electronegativity, shielding effect,
and ionic radius?
19. What is the difference between shielding electrons and valence electrons?
20. What is the difference between an anion and a cation? Why do they form?
21. How are the electron configurations of ions different than those of their neutral version?
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22. What is the octet rule?
23. Why are copper, silver, and zinc likely to form pseudo-noble gas electron configurations?
Chapter 7
1. How many electrons are lost or gained in the forming of each ion?
a. Ba2+
b. As3c. Cu2+
2. How can you determine the number of valence electrons in an atom of a representative element?
3. Atoms of which element tend to gain electrons? Atoms of which elements tend to lose electrons?
4. How many valence electrons are in each atom?
a. Potassium
b. Carbon
c. Magnesium
d. Oxygen
5. Draw the electron dot structure for each atom.
a. Potassium
c. Carbon
b. Magnesium
d. Oxygen
6. Use the dot structure to determine formulas of the ionic compounds formed when
a. Potassium reacts with iodine.
b. Aluminum reacts with oxygen.
7. What properties characterize ionic compounds?
8. How many electrons must each atom lose to attain a noble-gas electron configuration?
a. Ca
b. Al
c. Li
d. Ba
9. How many electrons must each atom gain to attain a noble-gas electron configuration?
a. N
b. S
c. Cl
d. P
10. Which of the following pairs of atoms would you expect to combine chemically to form an ionic compound?
a. Li and S
b. O and S
c. Al and O
d. F and Cl
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e. I and K
f. H and N
11. Explain why ionic compounds are electrically neutral.
12. Explain briefly why metals are good conductors of electricity.
13. Which of the following substances are most likely not ionic?
a. H2O
b. Na2O
c. CO2
d. CaS
e. SO2
f. NH3
14. Which particle plays the most important role in chemistry?
15. What is an ionic bond? An ionic compound?
16. What are the characteristics of ionic compounds?
17. What are the properties of metals and metallic bonds?
18. What is a monatomic ion?
19. What is true about the charges of transition metal ions?
Chapter 8
1. How are the melting points and boiling points of molecular compounds usually different from ionic
compounds?
2. What information does the molecular formula provide?
3. The following molecules have single covalent bonds. Draw an electron dot structure for each.
a. H2O2
b. PCl3
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4. What electron configurations do atoms usually achieve by sharing electrons to form covalent bonds?
5. How is a coordinate covalent bond different from other covalent bonds?
6. Draw the electron dot resonance structure for ozone (O3) and explain how they describe its bonding.
7. Draw electron dot structures for the following molecules, which have only single covalent bonds.
a. H2S
b. PH3
c. ClF
8. Identify the bonds between atoms of each pair of elements as nonpolar covalent, moderately polar
covalent, very covalent, or ionic.
a. H and Br
b. K and Cl
c. C and O
d. Cl and F
e. Li and O
f.
Br and Br
9. Place the following covalent bonds in order from least to most polar.
a. H-Cl
b. H-Br
c. H-S
d. H-C
10. Which of the following gases in Earth’s atmosphere would you expect to find as molecules and which as
individual atoms?
a. Nitrogen
b. Oxygen
c. Argon
11. Classify the following compounds as ionic or covalent.
a. MgCl2
b. Na2S
c. H2O
d. H2S
12. How many electrons do two atoms in a double covalent bond share? How many in a triple covalent bond?
13. Characterize a coordinate covalent bond and give an example.
14. Using electron dot structure, draw at least two resonance structures for the nitrite ion (NO2-). The oxygen
in NO2- are attached to the nitrogen.
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15. How must the electronegativities of two atoms compare if a covalent bond between them is to be polar?
16. What are the diatomic elements?
17. What is the difference between single, double, and triple covalent bonds?
18. What does it mean when resonance structures occur?
19. Explain how phosphorus and sulfur are sometimes exceptions to the octet rule.
20. What is a polar covalent bond?
21. What is a polar molecule?
22. What is hydrogen bonding and how does it work?
23. How can elements achieve a noble gas configuration?
24. What are the bond angles of some common molecules using the VSEPR theory?
Organic
1. What is an isomer? Give an example.
2.
Define and give an example of each of the following.
a. Alkane
b. Alkene
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c. Alkyne
3. What is the molecular formula of each of the following?
a. Methane
b. Ethane
c. Propane
d. Butane
Chapter 9
1. Write the formulas for compounds formed from these pairs of ions.
a. Ba2+, S2b. Li+, O2c. Ca2+, N3d. Cu2+, I2. Write formulas for these compounds.
a. Sodium iodide
b. Tin (II) chloride
c. Potassium sulfide
d. Calcium iodide
3. Write formulas for compounds formed from these pairs of ions.
a. NH4+, SO32b. Calcium ion, phosphate ion
4. Write formulas for these compounds.
a. Lithium hydrogen sulfate
b. Chromium (III) nitrite
5. Give the expected charge of the cations of these elements.
a. Sr
b. Ca
c. Al
d. Cs
6. What is the net ionic charge of every ionic compounds? Explain.
7. Which of the following compounds are binary ionic compounds?
a. KBr
b. K3PO4
c. Sodium nitride
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d. Calcium sulfate
8. Complete the table by writing correct formulas for the compounds formed by combining positive and
negative ions. Then name each compound.
NO3-
CO32-
CN-
PO43-
NH4+
a.
b.
c.
d.
Sn4+
e.
f.
g.
h.
Fe3+
i.
j.
k.
l.
Mg2+
m.
n.
o.
p.
a.
h.
n.
b.
i.
o.
c.
j.
p.
d.
k.
e.
l
f.
m.
9. What prefix indicates each of the following numbers of atoms in the formula of a molecular compound?
a. 3
d. 6
b. 1
e. 5
c. 2
f. 4
10. Write the formula or name for these compounds.
a. Boron trichloride
b. N2O5
c. Dinitrogen tetrahydride
d. CCl4
11. Give the name or the formula for these acids.
a. HCl
b. Sulfuric acid
c. HNO3
d. Acetic acid
12. Write names or formulas for these compounds.
a. Iron (II) hydroxide
b. Pb(OH)2
c. Copper (II) hydroxide
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d. Co(OH)2
13. Write formulas for these compounds.
a. Potassium permanganate
b. Calcium hydrogen carbonate
c. Dichlorine heptaoxide
d. Trisilicon tetranitride
e. Sodium dihydrogen phosphate
f.
Phosphours pentabromide
g. Carbon tetrachloride
14. Name these compound.
a. NaClO3
b. Hg2Br2
c. K2CrO4
d. HClO4
e. SnO2
f.
Fe(C2H3O2)3
g. KHSO4
h. Ca(OH)2
i.
BaS
15. Name each compound.
a. Mg(MnO4)2
b. Be(NO3)2
c. K2CO3
d. N2H4
e. LiOH
f.
BaF2
g. PI3
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h. ZnO
i.
H3PO3
16. A compound of general formula QxRy contains no hydrogen and Q and R are both elements. Neither Q nor R
is a metal. Is QxRy an acid, a binary ionic compound, an ionic compound containing a polyatomic anion, or a
binary molecular compound?
17. Explain what is wrong with each formula.
a. CsCl2
b. ZnO2
c. LiNe
d. Ba2S2
18. What is the difference between a molecular and an ionic compound?
19. How can you tell from looking at the formula that a compound is ionic or molecular?
20. What is a salt?
21. When given the name of the compound, how do you write formulas for molecular compounds?
22. When given the formula, how do you name ionic and molecular compounds?
23. How do you name and write formulas for acids and bases?
24. What is a binary compound? When given the name of the compound or the formula of the ions, how do you
write formulas for ionic compounds?
Chapter 10
1. How many moles is 2.80 x 1024 atoms of silicon?
2. How many moles is 2.17 x 1023 representative particles of bromine?
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3. How many moles are in 4.65 x 1024 molecules of NO2?
4. What is the mass of 1.00 mol of sodium hydrogen carbonate?
5. Find the number of moles in 3.70 x 10-1 g of boron.
6. Calculate the number of moles in 75.0 g of dinitrogen trioxide.
7. What is the volume of these gases at STP?
a. 3.20 x 10-3 mol CO2
b. 3.70 mol N2
8. At STP, what volume do these gases occupy?
a. 1.25 mol He
b. 0.335 mol C2H6
9. A gaseous compound of sulfur and oxygen, which is linked to the formation of acid rain, has a density of 3.58
g/L at STP. What is the molar mass of this gas?
10. What is the density of krypton gas at STP?
11. A compound is formed when 9.03 g Mg combines completely with 3.48 g N. What is the percent composition of
this compound?
12. When a 14.2 g sample of mercury (II) oxide is decomposed into its elements by heating, 13.2 g Hg is obtained.
What is the percent composition of the compound?
13. Calculate the empirical formula of each compound.
a. 94.1% O, 5.9% H
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b. 67.6% Hg, 10.8% S, 21.6% O
14. 1,6-diaminohexane is used to make nylon. What is the empirical formula of this compound if it is 62.1% C,
13.8% H, and 24.1% N?
15. Find the molecular formula of ethylene glycol, which is used as antifreeze. The molar mass is 62 g/mol and the
empirical formula is CH3O.
16. Which pair of molecules has the same empirical formula?
a. C2H4O2, C6H12O6
b. NaCrO4, Na2Cr2O7
17. Calculate the molar mass of each substance.
a. H2PO4
b. N2O3
c. CaCO3
d. (NH4)2SO4
e. C4H9O2
f.
Br2
18. How many moles is each of the following?
a. 15.5 g SiO2
b. 0.0688 g AgCl
c. 79.3 G Cl2
d. 5.96 g KOH
e. 937 g Ca(C2H3O2)2
f.
0.800 g Ca
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19. Find the mass of each substance.
a. 1.50 mol C5H12
b. 14.4 mol F2
c. 0.780 mol Ca(CN)2
d. 7.00 mol H2O2
e. 5.60 mol NaOH
f.
3.21 x 10-2 mol Ni
20. Calculate the percent composition of each compound.
a. H2S
b. (NH4)2C2O4
c. Mg(OH)2
d. Na3PO4
21. Which of the following can be classified as an empirical formula?
a. S2Cl2
b. C6H10O4
c. Na2SO3
22. What is the molecular formula for each compound? Each compound’s empirical formula and molar mass is
given.
a. CH2O, 90 g/mol
b. HgCl, 472.2 g/mol
23. Determine the molecular formula for each compound.
a. 94.1% O and 5.9% H; molar mass = 34 g
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b. 50.7% C, 4.2% H, and 45.1% O; molar mass = 142 g
24. Determine the empirical formula of compounds with the following percent compositions.
a. 42.9% C and 57.1% O
b. 32.00% C, 42.66% O, 18.67% N, and 6.67% H
c. 71.72% Cl, 16.16% O, and 12.12% C
25. What does the term STP mean?
26. What is the volume of 1 mole of any gas at STP?
27. When given the name of the compound or the formula of the ions, how do you write formulas for ionic
compounds?
28. What is Avogadro’s number?
29. What is a mole?
30. What are representative particles of elements? Molecular compounds? Ionic compounds?
31. How do you convert from representative particles to moles?
32. How do you convert from moles to representative particles?
33. What is molar mass?
34. How do you convert moles to mass? Mass to moles?
35. What is standard temperature and pressure?
36. How do you convert moles of a gas to liters @ STP?
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37. How do you find the density of a gas @ STP when you know the formula?
38. How do you find percent composition?
39. How do you determine the empirical formula and the molecular formula of a compound?
Additional Topics to Review
 Review the correct use of parentheses ( ) in formulas
 Know these molecular shapes: tetrahedral, pyramidal, bent, linear
 Study the stock and classical names for transition metal ions (p 255)
 Review the polyatomic ions (p. 257)
 Study p. 277 and p. 278 for extra help in naming and writing formulas of compounds
 Look over the note cards for all chapters.
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