FINAL EXAM PRACTICE PROBLEMS SAMPLER NAME ______
FALL 2012
USE YOUR OBJECTIVES SHEETS FROM UNITS 1-4 TO HELP GUIDE YOUR STUDYING!!!
THIS IS NOT THE END-ALL-BE-ALL GUIDE!!! ALSO REVIEW YOUR NOTES and CLASSWORK!!!
UNIT 1: THE ATOM and PERIODIC TABLE
1. What does the Law of Conservation of Matter state? Matter is not created or destroyed, only rearranged.
2. What are the four postulates of Dalton’s atomic theory? All matter is made of atoms. Atoms of like elements are the same and atoms of different elements are different. The atom is indivisible. Atoms combine in whole number ratios.
3. Thomson used the cathode ray tube and discovered something about the structure of the atom. Explain what happened in his experiment. What did these results tell him about the structure of the atom? A negative cathode ray was attracted to a positive plate. This told him there was a particle with a negative charge (electron) in the atom. Idea of atom changed from the marble model to the plum pudding model.
4. Rutherford used alpha particles in the gold-foil experiment and discovered something about the structure of the atom. Explain what happened in his experiment. What did these results tell him about the structure of the atom? Alpha particles went through gold foil. Some were deflected and some bounced back toward the source. Alpha particles are positive, and since most went through the foil, Rutherford discovered that the atom is mostly empty space. Rutherford also figured out that there was a dense positive mass (nucleus) in the center of the atom. Idea of atom changed from plum pudding model to nuclear model.
5. How was light used to help determine Bohr’s model of the atom? Bohr saw colored line spectra (not continuous) when observing light omitted by atoms of different elements. This was the basis for the idea that electrons must exist in specific energy levels and orbitals around the nucleus. Bohr said that when electrons moved from a higher energy level to a lower energy level, energy was released (often in the form of visible light).
6. What is the current model of the atom called? What does it look like? What experiment(s) led to this model? Quantum mechanical model or electron cloud model. Has s, p, d, and f clouds/sublevels in multiple energy levels. See book for picture. Was a result of Bohr’s experiments with hydrogen and light.
7. Answer the following questions about the proton:
location? nucleus
Charge? Positive (+)
8. Answer the following questions about the neutron:
location? nucleus
Charge? Neutral- no charge
9. Answer the following questions about the electron:
location? Around/ outside the nucleus
Charge? Negative (-)
10. What does the atomic number of an element indicate? Number of protons
11. What does the mass number of an element indicate? Number of protons plus number of neutrons
12. What is an ion? Give an example. An atom with a charge. Ex. Ca2+
13. What is an isotope? Give an example. Two atoms of the same element with different masses (different number of neutrons). Ex. C-12 and C-14. Each has 6 protons, but C-12 has 6 neutrons and C-14 has 8 neutrons.
14. Calculate the average atomic mass of argon using the following relative atomic masses and abundances of each
of the isotopes: argon-36 (35.96 amu, 33.7%), argon-38 (37.96amu, 6.3%), and argon-40 (39.96 amu, 60.0%).
35.96 x .337= 12.1
37.96 x .063= 2.4 add up totals Average atomic mass = 38.5amu
39.96 x .600= 24.0
15. What are the shapes of s orbitals? p orbitals? d orbitals? Sphere, dumbbell, clover
16. Draw the Lewis dot diagram for: a. Argon Ar with 8 dots b. Calcium Ca with 2 dots c. Nitrogen N with 5 dots
(both unpaired) (3 unpaired)
17. How many electrons, protons, and neutrons in the following (use the periodic table if necessary):
a. 35Cl- #p __17____ #n ___18_____ #e __18_____
b. 15N #p __7___ #n ____8_____ #e __7______
c. Al -28 #p __13___ #n ___15_____ #e _13______
d. C-14 #p __6____ #n __8______#e __6______
e. 31P3- #p __15____ #n __16______#e __18_____
18. Where are metals located on the periodic table? (Left hand side of the periodic table/metalloid or stair step line)
19. Where are non-metals located on the periodic table? (Right hand side of the periodic table/metalloid or stair step line)
20. List at least 4 properties of metals: (conduct heat & electricity, malleable, ductile, solid room temp. not mercury, high melting & boiling point, have luster, high density, form cations, ionic compounds)
21. What is a row (horizontal) in the periodic table called? (period)
22. Why do elements within a group on the periodic table often times have similar properties? (same number of valence electrons)
23. Which family of elements on the periodic table is non-reactive? Explain why they are non-reactive. (Noble Gases, group 18, VIIIA) (8 valence electrons)
24. As you go down a family in the periodic table, do atoms get bigger or smaller? Why? (bigger) (shielding effect)
25. As you go across each row in the periodic table, do atoms get bigger or smaller? Why? (smaller) (nuclear charge)
26. When an electrons falls from an excited state to a lower energy state is it absorbing energy or giving off energy? How do you know? (giving off) (spectra lines)
27. Draw the Bohr model for the sodium atom (be sure to include the # of protons and neutrons)
28. What is the name of group 2 on the periodic table? (Alkaline Earth Metals)
29. What is the name of group 17 on the periodic table? Review all group names. (Hallogens)
30. Classify each of the following as metal (M), non-metal (N), or metalloid (O)?
__N___ carbon __O___ boron ___M__ iron ___N__ hydrogen
__M___ potassium ___M__ gallium __N___ iodine __O___ silicon
31. I have one valence electron, am very reactive with water, and am found in the compound that you eat as table salt. What element am I? (Sodium)
32. I have 8 valence electrons, do not react with other elements, and am often times found in particular types of signs. What element am I? (Neon)
33. I have 6 valence electrons, am the most abundant element in the Earth, and am produced through the process of photosynthesis. What element am I? (Carbon)
34. Who developed the first periodic table? (Mendeleev)
35. Who was the first scientist to arrange elements according to atomic number? (Mosley)
36. Describe how ionization energy changes down a column and across a row of the periodic table. Why?
(Decreases DOWN column – shielding effect Increases ACROSS row – nuclear charge)
37. Which is bigger? Why?
a,__B__ B or C (nuclear charge) b. __I__ Cl or I (shielding effect) c _Ca___ Na or Ca (shielding effect)
d.__Cs_ K or Cs(shielding effect) e. _Mg_ Mg or Mg2+ (loss e-) f. _ Cl 1-___ Cl or Cl 1- (gain e-)
Calculation from unit one will not be on the final exam (E= hυ) or (c = υ λ)
UNIT 2: IONIC BONDING
38. Draw the orbital diagram for the following: Se, Ca2+ , S, Zn, and Rb Using your diagrams, explain why S tends to gain 2 electrons and Rb tends to lose 1 electron when forming ionic bonds.
Se / / / / / / / / / / / / / / / / / /1s / 2s / 2p / 3s / 3p / 4s / 3d / 4p /
Ca2+ / / / / / / / / /
1s / 2s / 2p / 3s / 3p
S / / / / / / / / /
1s / 2s / 2p / 3s / 3p
Zn / / / / / / / / / / / / / / /
1s / 2s / 2p / 3s / 3p / 4s / 3d
Rb / / / / / / / / / / / / / / / / / / /
1s / 2s / 2p / 3s / 3p / 4s / 3d / 4p / 5s
39. Write the complete electron configuration for the following: (also identify the number of valence electrons and highest principle energy level in each)
a. Rn 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d10 6p6
b. S2- 1s2 2s2 2p6 3s2 3p6
c. Mn 1s2 2s2 2p6 3s2 3p6 4s2 3d5
d. Nobelium (#102) 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d10 6p6 7s2 5f14
40. Write the abbreviated electron configuration for the following elements:
a. Cadmium [Kr] 5s2 4d10
b. Lead [Xe] 6s2 4f14 5d10 6p2
c. Sulfur [Ne] 3s2 3p4
41. a. Na loses 1; Na+ b. Al loses 3; Al3+ c. Cl gains 1; Cl1- d. O gains 2; O2-
e. K loses 1; K+ f. N loses 3; N3-
42. Metals lose electrons to non-metals which gain electrons in order to fulfill the octet rule. The resulting oppositely charged ions attract each other.
43. a. CuBr2 b. FeCl2 c. Hg(NO3)2 d. Mg(ClO3)2
e. Li2O f. PbS2 g. (NH4)2SO4 h.HCl
i. H2SO4 j. KOH k. Rb3P
44. a. copper (II) carbonate
b. calcium acetate
c. carbon dioxide
d. iron (II) chloride
e. sodium phosphate
f. potassium chloride
g. hydrofluoric acid
h. nitric acid
45. A chemical equation represents relative numbers of substances before and after a chemical reaction.
46. A reactant is a substance (usually written on the left side of a chemical equation) that enters into a chemical reaction. A product is what is formed in a chemical reaction.
47. The arrow means “to yield” or “to produce”.
48. A diatomic element is one that, when not bonded to another element, exists in a molecule of two atoms of the same element bonded together. (H2, O2, N2, Cl2, Br2, I2, F2)
49. Equations must be balanced to fulfill the Law of Conservation of Mass which states that matter cannot be created or destroyed.
50. a. synthesis Na2O + H2O à 2NaOH
b. synthesis H2 + F2 à 2HF
c. combustion C6H12O6 + 6O2 à 6CO2 + 6H2O
d. synthesis Ca + S à CaS
e. synthesis 2H2 + O2 à 2H2O
f. decomposition 2HgO à 2Hg + O2
g. double replacement NaOH + HCl à H2O + NaCl
h. single replacement Ca + 2HBr à CaBr2 + H2
i. single replacement 2Al + 3Pb(NO3)2 à 2Al(NO3)3 + 3Pb
j. single replacement No reaction- gold is not reactive enough to take hydrogen’s
place in the HBr.
k. double replacement AgNO3 + KI à KNO3 + AgI
l. single replacement Rb + AgNO3 à RbNO3 + Ag
m. double replacement Pb(NO3)2 + NaOH à NaNO3 + Pb(OH)2
n. Combustion 2C2H6 + 7O2 à 4CO2 + 6H2O
Beakers Illustration:
BaCl2 (aq) + K2(SO4)(aq) à BaSO4 (s) + 2 KCl (aq)
+ à
Beakers Illustration:
Mg (s) + 2 HCl (aq) à MgCl2 (aq) + H2 (g)
+ à
Not on Final
2 Ag(NO3) (aq) + CaI2(aq) à Ca(NO3)2(aq) + 2 AgI (s)
Complete Ionic (shows all ions if aqueous dissociate (separated)
2 Ag1+ + 2 (NO3)1- + Ca2+ + 2 I1- à Ca2+ + 2 (NO3)1- + 2 AgI (s)
Net Ionic (only shows the ions involved in making a change (in phase or charge).
2 Ag1+ + 2 I1- à 2 AgI (s) Spectators: Ca2+ + 2 (NO3)1-
UNIT 4: STOICHIOMETRY
51. How many sig figs? a. 100500 (4) b. 0.0043210 (5) c. 32.000 (5) d. 0.4510 (4) e. 11000 (2)
52. Round to 3 sig figs: a. 100500 (1.01 x 105) b. 0.0043210 (4.32 x 10-3) c. 32.000 (3.20 x 101)
d. 0.4510 (4.51 x 10-1) e. 11000 (1.10 X 104)
53. Quantitative analysis shows that a compound contains 32.28% sodium, 22.65% sulfur, and 44.99% oxygen. Find the empirical formula of this compound.
List % / Assume 100 g sample / Covert to Moles / Mole Answer / Mole to Mole Ratio*32.28 % Na / 32.28 g Na / x 1.00 mol Na
23.0 g Na / = 1.403 mol Na / 1.403 mol = 1.988 = 2
0.7056 mol
22.65% S / 22.65 g S / x 1.00 mol S
32.1 g S / = 0.7056 mol S / 0.7056 mol = 1.00 = 1
0.7056 mol
44.99% O / 44.99 g O / x 1.00 mol O
16.0 g O / = 2.81 mol O / 2.81 mol = 3.98 = 4
0.7056 mol
So the correct Empirical Formula is Na2S1O4 but one is understood so Na2SO4
· The mole to mole ratio is calculated by taking the smallest number of moles and dividing it into itself (result 1) and into to all other mole values. The Empirical Formula is the lowest whole number ratio of atoms in a compound.
54. Analysis of a 10.150g sample of a compound known to contain phosphorus and oxygen indicate a phosphorus content of 4.433g.
a. What is the % composition of the compound? The % of each element is determined by taking the mass of each element and dividing by the total mass.
4.433 g P x 100 = 43.67 % P 5.717 g O x 100 = 56.33 % O
10.150 g total 10.150 g total
b. What is the empirical formula of this compound? Using the format above and changing the % to grams or just using the grams provided the student will need to convert to moles and calculate the mole to mole ratio. The moles are 0.143 mol P [1.41 mol P] to 0.357 mol O [3.52 mol O], when the ratio is determined it is a (P) 1 : (O) 2.5. This is impossible, since the Empirical Formula is the lowest whole number ratio of atoms in a compound. So the final answer is P2O5 (which is doubling each value ( x 2 )