PRACTICE – Final Exam Name: ______
1) Complete the table below to convert between scientific notation and standard form 2 points
Standard Form Scientific Notation Standard Form Scientific Notation
62.3 g (example) / 6.23 x 101 g / 0.005300 cm1993 mL / 2.68023 x 102 kg
0.0000072 kg / 6.71 x 10-4 L
2) Use , , or = for each of the questions below. 3 points
a) 9.8 m ___ 9800 km b) 1 mL ___ 0.010 cL c) 3.5 g ___ .0035 kg
3) Complete the table below – Note that the particles are not is the same order: 3 points
Charge / Location / AMUNeutron
Electron
Proton
4) Draw the Lewis Structures for the following neutral atoms. 3 points
C F Si He Ga Ba
5) Fill in the blanks using one of the following elements. Please note that each element may be used once, more than once, or not at all. 14 points
K, Ca, Ga, C, As, O, Cl, Kr
______is an example of metal ______has 6 total electrons
______is an example of a nonmetal ______has 6 valence electrons
______is an example of a metalloid ______has 1 valence electrons
______is in period 4 ______is in group 5A
______is in period 2 ______is in group 4A
______is an Alkaline Earth Metal ______is a Halogen
______is a Nobel Gas. ______is an Alkali Metal
6) Look at each pair of elements below, select (circle) which element has the larger atomic radii:
Be or Ca Na or P N or B 2 points
7) Look at each pair of elements below, select (circle) which element has the higher ionization energy:
C or Si F or Br N or Li 2 points
8) Look at each pair of elements below, select (circle) which element has the higher electronegativity:
O or Se P or Mg N or F 2 points
9-10) Complete the following problems for Magnesium and Phosphorus 5 points
Identify each as either ionic or covalent compounds and provide either the name or formula 5 points
You may find these prefixes helpful for the questions below / mono- / di- / tri- / tetra- / penta / hexa- / hepta / octa- / nona- / deca-1 / 2 / 3 / 4 / 5 / 6 / 7 / 8 / 9 / 10
Ionic or Covalent / Name or Formula
11) dinitrogen trioxide
12) MgF2
13) phosphorus trifluoride
14) sulfur dioxide
15) N2O5
16) calcium nitrate
17) CO
18) lead (IV) oxide
19) FeSO4
For the following questions, put your answer in the box ___ at the beginning of the question.
20) ___ Which group of elements in the periodic table is known as the halogens?
21) ___ Use the periodic table to determine the number of electrons in a neutral atom of beryllium.
22) ___ Use the periodic table to determine the number of protons in an atom of bromine.
23) ___ What is the mass number for a carbon atom that has 7 neutrons in its nucleus?
Moving from left-to-right across a period (row) of the periodic table,
24) TRUE or FALSE - the ionization energy of the elements generally decreases
25) TRUE or FALSE - the atomic radius of the elements generally decreases
_____ 26) Which of the following is true about subatomic particles?
a. Electrons have no charge and have almost no mass.
b. Protons are negatively charged and the lightest subatomic particle.
c. Neutrons have a negative charge and are the lightest subatomic particle.
d. Electrons have almost no mass compared to the protons
_____27) All atoms are ____.
a. neutral, with the number of protons equaling the number of electrons
b. neutral, with the number of protons equaling the number of electrons, which is equal to
the number of neutrons
c. positively charged, with the number of protons exceeding the number of electrons
d. negatively charged, with the number of electrons exceeding the number of protons
_____28) The nucleus of all atoms ____.
a. always has the same number of neutrons and is considered neutral
b. are positively charged because of the positive charge of the protons
c. are negatively charged because of the negative charge of the neutrons
d. are positively charged and it occupies the vast majority of the volume of the atom.
_____ 29) The sum of the protons and electrons in an atom equals the ____.
a. atomic number b. charge of the atom c. atomic mass d. mass number
_____ 30) Isotopes of the same element have different ____.
a. numbers of neutrons b. numbers of protons c. numbers of electrons d. atomic numbers
_____31) All atoms of the same element have the same ____.
a. number of neutrons b. number of protons c. mass numbers d. mass
_____ 32) Which of the following elements is in the same period as krypton?
a. helium b. magnesium c. nitrogen d. bromine
_____ 33) Of the elements Fr, Sb, Al, and Rn, which is a metalloid?
a. Fr b. Sb c. Al d. Rn
_____ 34) Which of the following statements is NOT true?
a. Atoms of the same element must always have the same mass
b. Atoms of isotopes of an element have different numbers of neutrons.
c. The nucleus of an atom has a positive charge.
d. Atoms are mostly empty space
____ 35) Which of the following particles are free to drift in metals?
a. protons b. electrons c. neutrons d. cations
____ 36) Which of the following pairs of elements is most likely to form an ionic compound?
a. chlorine and oxygen c. aluminum and chlorine
b. nitrogen and sulfur d. sodium and lithium
____ 37) What characteristic of metals makes them good electrical conductors?
a. They have mobile valence electrons. c. They have mobile cations.
b. They have mobile protons. d. Their crystal structures can be rearranged easily.
____ 38) Which of these elements does not exist as a diatomic element?
a. H b. F c. Ar d. O
____ 39) How do atoms achieve noble-gas electron configurations in double covalent bonds?
a. Two atoms share one electron. c. Two atoms share two pairs of electrons.
b. Two atoms share two electrons. d. One atom completely loses two electrons to the other atom.
____ 40) When Group 6A elements form ions, they ____.
a. lose two protons b. gain two protons c. lose two electrons d. gain two electrons
____ 41) Which of the following is true about the composition of ionic compounds?
a. They are composed of anions and cations. c. They are composed of cations only.
b. They are composed of anions only. d. They are formed from two or more nonmetallic elements.
____ 42) Which element, when combined with bromine, would most likely form an ionic compound?
a. lithium b. carbon c. phosphorus d. chlorine
____ 43) Which of the following occurs in an ionic bond?
a. Oppositely charged ions attract. c. Two atoms share more than two electrons.
b. Two atoms share two electrons. d. Like-charged ions attract.
____ 44) Which of the following pairs of elements is most likely to form an ionic compound?
a. magnesium and fluorine c. oxygen and chlorine
b. nitrogen and sulfur d. sodium and aluminum
____ 45) A bond formed between a calcium atom and an oxygen atom is likely to be ____.
a. ionic b. coordinate covalent c. polar covalent d. nonpolar covalent
____ 46) Which of the following covalent bonds is the most polar?
a. H—F b. H—C c. H—H d. H—N
____ 47) Which of the following atoms acquires the most negative charge in a covalent bond with hydrogen?
a. C b. Na c. O d. S
____ 48) Which of the forces of molecular attraction is the weakest?
a. dipole interaction b. dispersion c. hydrogen bond d. single covalent bond
49) Compare and Contrast Ionic and covalent bonds. 3 points
a) how are they similar?
b) how are they different?
50) Complete the Lewis Structures for each Compound (there will be two of these) 4 points
(NOTE: I will provide you with a copy of the electronegativities of the atoms)
Identify each of the Bond Types as either: Nonpolar Covalent, Polar Covalent, or Ionic
(Do not forget to include the identification of cations and anions or partial positive and partial negative where appropriate.)
4 points
51) HO 52) NaI
53) HC 54) H-H
55) FO 56) HN
Identify each of the following as either intermolecular or intramolecular 2 points
57) ______Polar Covalent Bond 60) ______Ionic Bond
58) ______Hydrogen Bond 61) ______Dipole-Dipole Interact
59) ______London Dispersion 62) ______Van der Waals
63) Explain “why” atoms become partially positive and partially negative in a polar bond.
2 points
1 Atm = 101.3 kPa = 760 mmHg = 14.7 PSI 1 m3 = 1,000 L 1 lb = 454 grams 1 gallon = 4 quarts
1 km = 1,000 m 1 m = 100 cm = 1,000 mm 1 mL = 1 cm3 1oz = 28.35 g 1 quart = 2 pints
1 kg = 1,000 g 1 g = 100 cg = 1,000 mg 1 inch = 2.54 cm 1 lb = 16 oz 1 pint = 2 cups
1 kL = 1,000 L 1 L = 100 cL = 1,000 mL 1 mile = 5280 ft 1 quart = 0.95 L 1 cup = 8 oz. (fl)
Perform the following conversions: (Show your work to earn partial credit) Missing units = loss of credit
NOTE: There will be five problems like those below (64-71)
64) 147.2 kPa = ____ mmHg
65) 3.4 Atm = ____ kPa
66) 894.2 cm = ____ km
67) 287,520 mg = ____ kg
68) 97 mL to gallons
69) 133.5 mm to inches
NOTE: You will need to know Avagadro’s Number for the following two questions
70) How many atoms of platinum are in the 4.2 g ring?
71) If there are 5.45 x 1025 atoms of aluminum (Al), how many grams would this be?
72) TRUE or FALSE: The motion/speed of the particles increases as the temperature of particles decreases.
73) TRUE or FALSE: Because the shape of a solid never changes, the particles of the solid do not move.
74) TRUE or FALSE: Particles can stick together because of the intermolecular forces between the particles.
75-78) Balance the following equations 2 points
[ ] NO2 + [ ] O2 à [ ] N2 [ ] NaI + [ ] Pb(SO4)2 à [ ] PbI4 + [ ] Na2SO4
[ ] Fe(OH)3 à [ ] Fe2O3 + [ ] H2O [ ] H3PO4 + [ ] CaBr2 à [ ] HBr + [ ] Ca3(PO4)2
3 points
79) Look at the equations in questions 75-78 above. List each chemical as either a reactant (R) or a product (P).
C PbI4 Fe(OH)3 NaBr NaI H2O C3H8 H2
80) How does the Law of Conservation of Mass relate to balancing chemical equations? 2 points
82-85) Write balanced equations for the following word equations. Indicate the type of reaction on the line to the left of the equation by classifying each reaction as single replacement (SR), double replacement (DR), decomposition (D), synthesis (S), or combustion (C). (NO PHASE LABELS) 4 * 3 points each = 12 points
____ 82) lithium nitride + barium nitrate à lithium nitrate + barium nitride .
____ 83) sodium bromide + calcium hydroxide à ______.
____ 84) When a balloon filled with hydrogen gas is ignited, it forms ______. (write the balanced eq)
____ 85) Aqueous ammonium hydroxide reacts with aqueous copper (II) nitrate to produce solid copper (II) hydroxide and a solution of ammonium nitrate.
87) How did you know that a chemical reaction occurred when you added the magnesium metal to the solution of hydrochloric acid? 2 points
88) How many grams of lead (II) phosphate will be formed when 85.39 grams of phosphoric acid (H3PO4) react with an excess of lead metal in the following reaction? 5 points
____ Pb + ____ H3PO4 à ____ H2 + ____ Pb3(PO4)2
MM =
MM =
x ------x ------x ------=
89) How many grams of hydrofluoric acid (HF) are required to react completely with 23.68 grams of calcium hydroxide in the following reaction? 5 points
____ Ca(OH)2 + ____ HF à ____ CaF2 + ____H2O
MM =
MM =
x ------x ------x ------=
90) TRUE or FALSE: Based upon the chemical equation in # 89 above, if more CaF2 where added, then the reaction would shift right.
91) TRUE or FALSE: Based upon the chemical equation in # 89 above, if more HF is added, then the concentration of H2O will increase.
92) TRUE or FALSE: Some chemical reactions can go in reverse, forming reactants.
93) EXPLAIN what it means when a reaction is in equilibrium.
94) List four things that affect the rate of a chemical reaction.
95) What must occur in order for a chemical reaction to occur?
96) How does the addition of a catalyst change the activation energy of a chemical reaction?
97) Calculate the percent yield of an experiment if the stoichiometry calculations stated that it should have produced 2.97 g of H2O, but during the experiment, only 2.47 g of H2O was produced.