Download Review - Final Exam

Review - Final Exam
1.
What is the difference between a theory and a law?
2. What is the difference between qualitative and quantitative? Give two examples of each.
3. What is matter? What else can we measure that is not matter?
4. What is meant by properties of matter? Name several intensive properties of water.
5. List as many properties that you can that would distinguish between a metal such as silver and a nonmetal such as sulphur.
6. What is the difference between homogeneous and heterogeneous? Which of these terms only applies to
pure substances? Explain. How can the other term apply to substances and mixtures? Use examples to
explain why.
7. What is the difference between: an element and a compound, an element and an atom, a compound and a
molecule, & an element and an ion? Is it possible to have a molecule of an element? Explain.
8. Explain how you would separate the following mixtures. List the differences in properties that allow for
the separation.
a.
two dyes of different color
b. a solution of alcohol and water
c.
fine white sand and salt
9. Name the following changes in state: Solid – liquid, liquid – solid,
liquid – gas, gas – liquid, solid – gas, gas – solid.
10. Sketch a cooling curve for a pure substance that cools from 100°C to
-50°C if its boiling point is 80°C and its freezing point is -20°C.
11. Calculate the Rf values for the chromatograph below. Show your work.
Solvent line
Temp (°C)
12. Is the heating curve shown below for a mixture or for a pure substance? Explain.
Time (min)
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13. Use this diagram to answer the questions below.
a
b
c
d
a.) Which diagram(s) represent pure substances? … mixtures? ….elements? ….compounds?
b.) Which state would each of these diagrams represent? Explain.
14. Who was the first person to suggest that all matter was made of small particles? Who suggested that
all matter was made out of only four elements? What were the four elements? State Dalton’s Atomic
Theory. Who discovered the electron? What was his model called and why? Who discovered the
nucleus and how was it done? What did Bohr contribute to our understanding of the atom?
15. What are the three subatomic particles, their symbols, their charge, and relative mass?
16. Complete the following table for the first four subshells.
Subshell
Starts at level #
Number of
orbitals
Number of
electrons
17. Complete the following table for the first five energy levels.
Energy level
Number of subshells
Number of electrons
18. What formula could you use to calculate the number of electrons that each energy level can hold? How
many electrons can fit into the eighth energy level?
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19. Draw an orbital diagram for V and P. Label the orbitals in each diagram. (On your exam, the orbitals will
be drawn for you. You will have to put in the electrons and label the orbitals.)
20. Do long form electron configurations for the following:
a.) Mg b.) Mn c.) Se d.) As
21. Do short form configurations for:
a.)
42Mo
b.)
51Sb
c.)
28Ni
d.)
55Cs
22. What are the valence electrons? What happens to the valence electrons across a row of representative
elements and down a group? What information is gained from the valence electrons? Why do chemical
families have similar properties?
23. Draw Lewis (Electron Dot) symbols for the elements across the second period and use these to predict
the combining number for each element.
24. What happens to the size of an atom across a period of representative elements and down a group. In
each case explain why.
25. What is ionization energy (IE)? In general, what happens to IE across a row of representative elements
and down a group? Explain why. What two elements in period 3 are exceptions to the general trends?
Explain why.
26. Which has a larger second ionization energy, Mg or Na? Explain why.
27. What happens to the size of an atom when electrons are lost and gained? Explain why.
28. Write the formulas for two positive ions and two negative ions that have the same number of electrons
as Ne and arrange them in order from smallest to largest. Write the long form configurations for each
of these ions.
29. Arrange the following groups in order from smallest size to largest:
a.) F, C, N b.) Na, K, Rb c.) Se, S, Cl d.) Cl1-, Ar, K1+ e.) F1-, F, F1+ f.) Co, Co2+, Co3+
30. Arrange the following groups from smallest ionization energy to largest.
a.) Mg, Ca, Sr b.) B, C, N
31. What electrons can the d-block metals use as valence electrons? Write short form electron
configurations for: a.) Cr3+, b.) Co2+, c.) Cu1+
32. What are the trends in metallic properties across a row of representative elements and down a group.
Which is the most metallic element on the periodic table? Give two reasons why. List three properties
of all metals.
33. Describe how you determine the number of electrons, protons and neutrons in a neutral element. What
is an isotope?
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34. How did Mendeleev arrange the elements in his periodic table? How is the modern periodic table
arranged?
35. What are the rows on the periodic table called? What are the columns on the periodic table called?
What are the names of the following groups; 1, 2 ,17 ,18? What is the name given to the d-block metals?
What is the name given to the elements in the s & p blocks?
36. Complete the following table:
Particle
37
17
Atomic
Number
Atomic
Mass
Number of
Protons
Number of
Electrons
29
14
14
Number of
Neutrons
Cl
33
16
S
2−
35
79
36
15
18
16
37. How are ionic compounds formed?
38. What is an ionic bond?
39. Are ionic bonds strong or weak? How do we know this?
40. What two factors determine how strong an ionic bond is? Explain.
41. Which has the stronger ionic bonds, BeF2 or LiF? Give two reasons.
42. Why can’t ionic compounds conduct an electric current in the solid state?
43. Why can ionic compounds conduct electricity when dissolved in water?
44. Describe a small experiment that could be done to determine whether a compound is ionic or covalent?
45. From the following list of compounds identify those that are ionic and write dissociation equations for
them. CH4, K2Cr2O7, NH4Br, SO2, Mg(NO2)2, PCl3, Fe2(CO3)3
46. How are covalent compounds formed?
47. What is a covalent bond?
48. Draw Lewis structures for the following compounds, name their shapes, & predict which are polar or nonpolar. Which of these compounds will be expected to dissolve in water? Explain. CH4, NH3, SO2, Cl2,
BH3, H2S, CO2
49. Draw Lewis diagrams for the following: CO32-, NH41+, PO43-
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50. Name the following compounds: Li2SO3, CoSO4, CF4, N2O5, HCl
(aq),
Al2(SO4)3, FePO4•5H2O, Hg(HCO3),
CO, CH4
51. Write the formulas for the following compounds: ammonia, manganese(IV) oxide,
Cobalt(II) chloride hexahydrate, Hydrochloric acid, nitric acid, sulphur trioxide,
mercury(II) phosphate, zinc dihydrogen phosphate, hydrogen peroxide, aluminum carbide
52. Equations
a.) Balance the following equations
i. CaC2 + O2 → Ca + CO2
ii. C + SO2 → CS2 + CO
iii. BN + F2 → BF2 + N2
iv. Al2C6 + H2O → Al(OH)3 + C2H2
v. NO2 + H2O → HNO3 + NO
b.) Write the products for the following reactions and balance them.
vi. Cu + FeSO4 →
vii. C5H10 + O2 →
viii. Al + I2 →
ix. Mg(OH)2 + HBr →
x. Al2(SO4)3 + Na2CO3 →
c.) Write the ionic and net ionic equation for iv & x in 52.b) above.
53. Explain the difference between exothermic and endothermic reactions. Draw and label energy diagrams.
54. What is the molar mass of Ni2(SO3)3?
55. What is the mass of 3.65 of CO2?
56. How many moles of SO2 are there in 12.6 SO2?
57. How many moles of O2 are there in 2.45 × 10 of H2O?
58. How many molecules of Ca are there in 1.34 of Ca?
59. How many mol of N2 (g) are there in 46.1 of N2 @ STP?
60. What is the mass of 16.9 CH4 (g) at STP?
61. What volume of F2 (g) at STP would 6.19 × 10 molecules of F2 have?
62. How many molecules of NO2 are there in 4.87 NO2?
63. What is the mass in grams of 1 of K?
64. A certain amount of P2O3 has 3.98 P. How many grams of O are there?
65. How many molecules of CO2 are there in 68.2 NO2?
66. What is the density of NH3 (g) at STP?
67. The density of = 1.59
"
#$
. How many mol of CCl4 are there in 87.1 of CCl4?
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68. What is the % composition of Ca(NO3)2?
69. What is the % of H2O in CoSO4⋅6H2O?
70. Determine the empirical formula for a compound consisting of 50.0% C, 10.0% H, and 40% O.
71. 1.99 of Al combined with oxygen to form an oxide of aluminum with a mass of 3.75. What is the
empirical formula for this compound?
72. A compound has an empirical formula of C2H4O and a molecular mass of 220.0
"
#&'
. What is its molecular
formula?
73. A compound containing only C and H was burned in a combustion reaction to produce 11.0 CO2 and
5.40H2O. What is the empirical formula of the compound? If the molar mass of the compound is
144.0
"
#&'
what is the molecular formula?
74. How many grams of K2SO4 are needed to prepare 250.0 of 0.350( K2SO4?
75. What is the molarity of a solution in which 45.6 of solution contained 8.34 MgCl2?
76. A chemist diluted 27.9 of 12.00( HCl to a volume of 500.0. What is the new concentration of
HCl?
77. How many mL of 5.34( NaCl are needed to produce 234 of 1.74( NaCl?
78. If 34.8 of 0.74( KBr are mixed with 51.2 of 0.63( KBr, what is the resulting concentration of
KBr?
79. What are the concentrations of each of the ions in 0.35( Al2(SO4)3? Include the dissociation equation.
80. Given - 4 Al
a.
(s)
+ 3 O2 (g) → 2 Al2O3 (s)
How many grams of Al2O3 can be produced from 34.2 Al?
b. How many grams of Al are needed to completely react with 13.8 of O2 at STP?
81. Given - Sr(OH)2 (aq) + 2 HCl
a.
(aq)
→ SrCl2 (aq) + 2 H2
If it takes 28.39 )0.760( HCl to neutralize 25.00 of Sr(OH)2 what is the concentration
of Sr(OH)2?
b. What volume of 0.450( Sr(OH)2 is needed to neutralize 245.0 of 0.800( HCl?
82. Given - 4 NiS
a.
(s)
+ 7 O2 (g) → 2 Ni2O3 (s) + 4 SO2 (g)
If 45.2 NiS is reacted with 37.6 O2 how many grams of Ni2O3 will be produced?
b. What is the limiting reactant?
c.
How many g of the excess reactant will remain unconsumed?
83. Given - 16 Fe + 3 S8 → 8 Fe2S3
If 35.0 of Fe were reacted with excess S8 to produce 49.4 Fe2S3 what is the % yield for the
reaction?
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