Download File first semester final study guide key

Chemistry I
First Semester Final Review Material
Unit I: Introduction to Chemistry
Short Answer/Fill in the Blank: Identify the letter of the choice that best completes the
statement or answers the question.
1. Chemistry can best be defined as…?
__the study of matter and all the changes it may undergo_________________________
2. Matter is anything that has __mass_______ and __volume________.
3. When writing the problem/statement of an experiment it is always best to write it in the
form of a __question________.
4. Must a hypothesis always be proven correct? Explain why or why not.
No; hypotheses are educated attempts to explain an observation but they may not explain
a scientific observation.
5. What is the purpose of a control group?
To provide a comparison of the experimental study to a norm.
6. Rank the three states of matter studied in this unit according to their respective densities.
Most dense
1. Solid
2. Liquid
Least dense
3. Gas
7. A pure substance that cannot be separated into simpler substance by a chemical change is
called a(n) __element___________. A substance that contains two or more elements
chemically combined in a fixed proportion is called a(n) ___compound______.
8. ___Distillation_________________ is a mixture separation technique for homogeneous
mixtures that utilizes differences in boiling points, while ____filtration____________ can
be used to separate heterogeneous mixtures such as sand and water.
Completion: Complete the following answers by following the directions given in each question.
9. Fill in the following chart to describe what happens to the shape and volume of each state
of matter when the sample is placed into a new container.
Shape
Volume
Gas
Takes the shape of the
container it is in.
Same volume as the
container that it is in.
Retains shape
Maintains its’ original
volume.
Takes the shape of the
container it is in.
Will fill the container
incompletely.
Solid
Liquid
10. Give an example of each of the following types of substances.
a. Pure substance (element)
____diamond_________
b. Pure substance (compound)
____water____________
c. Homogeneous mixture
____apple juice_______
d. Heterogeneous mixture
____________________
11. Give one example of each of the following processes.
a. Physical change
_____tearing paper_______________
b. Chemical change
_____digesting of food____________
Short Answer: Answer the following questions, being sure to include appropriate terminology
and/or definitions.
12. A student is given an unknown substance for analysis. The substance is an orange liquid
and appears to be uniform throughout. What would the student do to determine whether
the substance is a mixture or pure substance? If it is a mixture what separation technique
could be used to separate the mixture.
The student could use distillation or crystallization to separate the possible mixture.
13. What are constants? What is the importance of having constants in an experiment?
Constants are variables that are kept the same by the experimenter. They are important so
you can see the way the independent variable affects the dependent variable.
14. Describe two (2) ways that chemistry affects your daily life.
Answers will vary. Examples: Breathing and digestion of food.
15. What is the relationship between the independent variable and the dependent variable of
an experiment.
The independent or controlled variable is the variable that affects the dependent or
uncontrolled variable in an experiment.
16. A student is given a heterogeneous mixture of two unknown substances. One substance is
magnetic and the other substance is not. How might the student easily separate the
mixture?
Use a magnet to separate the magnetic substance from the non-magnetic substance.
17. What may happen if you do not communicate the results of your experiment to an
audience after it has been concluded?
The cure for rabies may not be found.
18. Using examples, compare qualitative and quantitative data.
Qualitative the candle was longer than the pencil.
Quantitative: the candle was 3 inches longer than the pencil.
Application: Use the following experiment description to answer questions 19 - 26.
A class of students decided to determine how various exercises affect the amount of calories they
burn while exercising for a 30 minute period. The amount of calories burned during a 30 minute
period was determined by using a heart rate monitor that could calculate the amount of calories
burned. Students then performed 5 different exercise routines including weightlifting, running,
walking, yoga, and biking. Students began the experiment by measuring the amount of calories
burned in a thirty minute period while resting. At the end of the experiment each student shared
their results with the other students in order to determine which exercise routine was the most
effective.
19. What is the problem for this experiment?
Which exercise routine will burn the most calories in a 30 minute period?
20. Write an appropriate hypothesis for this experiment.
Running for 30 minutes will burn the most calories when compared to weightlifting,
walking, yoga, and biking.
21. What is the independent variable?
Exercise activity
22. What is the dependent variable?
Calories burned
23. What is the control group?
Calories burned during rest
24. Identify 2 constants for the experiment.
Answers will vary. Examples: Temperature and clothing.
25. Was the data collected in this experiment qualitative or quantitative?
Quantitative
26. After conducting this experiment, Allison concludes that “running is the best exercise.”
Would this be an appropriate conclusion for this experiment? Explain your answer.
No; although running may burn the most calories it may not be the “best” exercise.
Unit II: Chemical Foundations
Fill in the Blank: Fill in the blanks with the appropriate term listed below.
atom
electron
atomic number
isotope
average atomic mass protons
nucleus
cation
anion
periods
neutrons
groups
The periodic table organizes elements by putting them in __groups__________ or the elements
in a vertical column on the periodic table and by _periods____ which are the elements in a
horizontal row on the periodic table. Elements are placed in the periodic table in numerical order
according to their ____atomic number_________ , which tells us the number of protons in the
nucleus of an element. The ____atom____________ is the fundamental unit of an element. The
central core of an atom is the ____nucleus_________, which contains ___protons_______,
which are positively charged subatomic particles, and ___neutrons_______, which are subatomic
particles with no charge. ____Electrons____ are negatively charged particles that are found
outside the nucleus. Atoms can become charged when they gain or lose electrons. A
__anion_______ is a negatively charged ion and a ____cation______ is a positively charged ion.
When an atom has a different number of neutrons in the nucleus it is then called an
___isotope___.
Their presence is what allows us to calculate the ___average atomic mass____ which is our final
number on the periodic table.
Short Answer: Using a brief statement, answer each of the following questions about atomic
structure.
1. John Dalton’s atomic theory states that, “atoms of a given element are identical in size,
mass, and other properties; atoms of different elements differ in size, mass, and other
properties.” The underlined portion of this statement is incorrect. Why?
Because of the presence of ions and isotopes.
2. How did Rutherford’s model of the atom differ from Thomson’s plum pudding model?
Rutherford had a positively charged nucleus at the center of the atom. Thomson had
positive charges spread throughout the entire atom.
3. What were two main conclusions of Rutherford’s gold foil experiment?
1. The atom is mostly empty space.
2. At the center of the atom is a dense, positive nucleus.
4. J.J. Thomson used a cathode ray tube to discover which negatively charged subatomic
particle?
Electron
5. Which scientist built off of Thomson’s work with the electron? What did he determine?
Millikan. He determined the mass-to-charge ratio of the electron.
6. What is the atomic number and mass number of an atom of an element with 13 protons
and 16 neutrons?
Atomic number -13 Mass Number - 29
7. List, in order, from least massive (has the least mass) to most massive (has the most
mass) the following particles: hydrogen atom, electron, helium atom, proton
Electron; proton; Hydrogen atom; Helium atom
8. What is the name of the number, that is used to identify an element?
Atomic number
9. An ion always contains an unequal/equal (circle one) number of protons and
electrons/neutrons (circle one).
10. An atom of uranium that contains 92 protons and 143 neutrons would be called:
Uranium-235
11. How many protons, neutrons and electrons does a neutral atom of an element with atomic
number 50 and mass number 120 contain? Identify the element.
Protons – 50; neutrons – 70; electrons - 50
Tin is the element
12. What is the family/group name for the elements in column 6A (16)?
Oxygen family/Chalcogens
13. What is the family/group name for the elements in column 4A (14)?
Carbon family
14. What is one way that Mendeleev organized his periodic table?
Increasing atomic mass or chemical reactivity.
15. How are metals and nonmetals distinguished on the periodic table?
Metals are mainly on the left side of the periodic table and non-metals are on the
right
side of the table.
16. Why are the noble gases called “inert gases”?
They do not react.
Problems: Solve the following problems in the space provided. Show your work where
necessary.
17. Determine the number of protons, neutrons, and electrons in the ion
Protons: 23; neutrons: 28; electrons: 20
51
V3+.
23
18. A neutral atom has 10 protons and 11 neutrons in its nucleus. Write out its nuclear
symbol. (1.5 points)
21
10
Ne
19. An ion contains 50 protons, 69 neutrons, and 46 electrons. What is its nuclear symbol?
119
50
Sn+4
20. Complete the following table.
Symbol
# protons
# neutrons
112
Cd
48
48
64
79 2Se
34
34
45
189 3Bi
83
83
106
35 1Cl
17
17
18
137
2+
Ba
56
56
81
# electrons
Net Charge
48
0
36
2-
86
-3
18
-1
54
2+
Short Answer: Answer the following questions, being sure to include appropriate terminology
and/or definitions.
21. Compare and contrast metalloids and non-metals.
Metalloids are mainly solids at room temperature and exhibit properties of both metals
and
non-metals. Non-metals are solids, liquids, and mainly gases at room temperature.
22. List 2 nonmetals from the periodic table (using their symbols) and describe how they may
be used.
Answers will vary. Examples: Oxygen is used to breathe. Neon is used in lighting.
23. List 2 metals from the periodic table (using their symbols) and describe how they may be
used.
Answers will vary. Examples: Gold is used in jewelry. Lithium is used in batteries.
24. Are isotopes and ions considered atoms? Why or why not?
Yes, they just have a different number of electrons and neutrons.
Unit III: Compound Formulas and Nomenclature
Questions: Write the correct answer to each question or statement in the space provided.
1. Why do most elements tend to gain or lose electrons when they become involved in the
bonding process? To satisfy the octet rule.
2. How many electrons will an atom of any element in the alkali metal family lose when
they bond with a non metal? one
3. What is one difference between an ionic compound and a covalent compound?
Answers will vary.
4. Which of the following compounds is not ionic: SiCl2 , NaCl, N2O5 , S2F2
5. Which of the following compounds is covalent: NaNO3, CaI2, SO2, RbO
6. During the formation of covalent compounds, electrons are __shared_______ between
atoms. An ionic bond forms when electrons are ___transferred_______ from one atom to
another. (2 points)
7. Metals will typically _lose_______ electrons to attain a noble gas electron configuration.
8. What are valence electrons and why are they important? (2 points)
valence electrons are the outermost electrons and are involved in bond formation
9.
What does the octet rule state?
Atoms will gain, lose or share electrons in order to attain a complete set of 8.
Short Answer: Answer the following questions in complete sentences.
10. Compare and contrast ionic and covalent compounds in terms of their properties and the
elements that make up the compounds. Be specific and detailed in your answer.
Ionic – metal and nonmetal; electrons are transferred; brittle; high melting points;
conduct electricity in water.
Covalent – nonmetal and nonmetal; electrons are shared; soft; low melting points; do not
conduct electricity in water.
11. What do the prefixes in covalent compounds tell us?
Exact number of atoms in the compound.
12. Why do we use roman numerals when naming some ionic compounds but not others?
Some metals have multiple charges. The roman numeral tells the charge that the metal
has.
13. Give an example of both an empirical formula and a molecular formula. What are the
differences between the two?
CH2O – empirical (lowest whole number ratio)
C6H12O6 – molecular (true formula)
Problems: Solve the following problems in the space provided. Show your work.
14. Write the name of each of the following compounds in the space provided.
a. Li3PO4
_____lithium phosphate__________________________
b. S2F8
_____disulfur octafluoride________________________
c. Cu(OH)2
_____copper (II) hydroxide_______________________
d. (NH4)2SO4
_____ammonium sulfate__________________________
e. AuF3
_____gold (III) fluoride___________________________
f. N3O6
_____trinitrogen hexaoxide________________________
g. SiO4
_____silicon tetraoxide___________________________
h. Mg(NO3) 2
_____magnesium nitrate__________________________
i. NaCN
_____sodium cyanide_____________________________
j. CBr4
_____carbon tetrabromide__________________________
15. Write the chemical formula that represents each of the following compounds on the space
provided.
a. silver sulfate
_____Ag2SO4_________________________
b. dihydrogen monoxide _____H2O_________________________
c. sodium nitride
_____Na3N_________________________
d. lead (III) oxide
_____Pb2O3_________________________
e. lithium chloride
_____LiCl_________________________
f. manganese phosphate _____Mn3(PO4)2_________________________
g. carbon tetrachloride _____CCl4_________________________
h. nickel (II) cyanide
_____Ni(CN)2_________________________
i. iron (II) acetate
_____Fe(C2H3O2)2_________________________
j. nitrogen dioxide
_____NO2_________________________
Unit IV: Measurements and Calculations
Short Answer/Multiple Choice: Circle the letter of the answer that correctly answers the
question or completes the sentence.
1. Why is the SI unit for mass the kilogram and not the gram?
Because of large items
2. Write the measurement 3.46 x 10-3 g in standard notation.
0.00346 g
3. How many decimeters are in 25 kilometers?
250,000 dm
4. The smallest volume from among the following is
a. 0.014 L.
b. 1.8 cL.
c. 250 mL.
d. 1.6 x 10-2 L.
5. What is the SI unit for time?
seconds
6. The largest mass from among the following is
a. 1.8 g.
b. 1.6 x 102 dg.
c. 0.012 mg.
d. 25 ng.
7. Which of the following statements is true?
a. All zeros in numbers are not significant.
b. Numbers trapped between non-zero numbers are significant.
c. Leading zeros are significant only if a decimal is absent.
d. Trailing zeros are never significant.
8. In scientific notation, what does the sign (+/-) of the power of ten indicate?
which way to move the decimal
9. What is the rule for an addition and subtraction problem? (How do we determine how
many significant digits our answer will have?)
Based off of the least precise measurement
10. What is the process of using conversion factors to change from one unit to another called
Dimensional Analysis
Short Answer: Answer the following questions in complete sentences.
11. What is the purpose of using significant figures in science? Why are they more important
to use in science than in math?
Significant figures illustrate the uncertainty that every measurement has with it.
12. Why is the metric system considered easier to use than the English system of
measurement?
Based off of powers of 10.
13. What is wrong with the following number written in scientific notation?
45.68 x 10-6
45.68 is not a number between 1 and 10. It should be 4.568
Problems: Solve the following problems in the space provided. Show your work.
14. Write the following numbers in scientific notation.
a. 51,200,000 cm
____5.12 x 107___________
b. 0.0458 mL
____4.58 x 10-2___________
c. 0.000000781 km
____7.81 x 10-7___________
d. 540,000 s
____5.4 x 105___________
15. Convert the following numbers to standard notation.
a. 1.64 x 10-4 m
_____0.000164 m_______________
b. 7.2 x 107 dL
_____72,000,000 dL_______________
c. 3.25 x 10-2 nm
_____0.0325 nm_______________
d. 6.5 x 101 ms
_____65 ms_______________
16. How many significant figures are there in each of the following measurements?
a. 34.000 g
__5_____
b. 0.0006789 cm
__4____
c. 2108.45 km
__6_____
d. 0.870 mm
__3_____
e. 560. L
__3_____
17. Perform the following operations. Make sure that your answers have the correct number
of significant figures. (No Naked Numbers!)
a. 26.5 cm x 200.50 cm
5310 cm2
b. 307.506 km – 156 km
152 km
c. (4.650 X 109 mm)/ (8.2 X 10-6 s)
5.7 x 1014 mm/s
d. 268.500 g + 86.47 g
354.97 g
18. Convert the following metric measurements. (You do not have to show dimensional
analysis.)
a. 5.0 x 103 mm to km
0.0050 km
b. .0257 dL to cL
0.257 cL
c. 620 pm to nm
.620 nm
d. 9.9 kg to cg
990,000 cg
19. A teacher brings his students 36 dozen doughnuts to eat. However, he decides to give 1.5
dozen doughnuts to the office staff. How many total doughnuts does he have left for the
students? Show your work in dimensional analysis.
414 doughnuts
20. Perform the following conversion using dimensional analysis: 3.6 days to centiseconds.
Show your work in dimensional analysis.
3.1 x 107 cs
21. A student measures her foot to be 5 inches long. How long is her foot in meters? In
kilometers? Show your work in dimensional analysis and express your answer to the
appropriate number of significant figures. (There are 2.54 cm in one inch.)
0.127 m
1.27 x 10-4 km
22. Convert 25 km/hr to miles/min. (There are 1.609 km in a mile.)
0.26 miles/min
Unit V: The Mole
Short Answer/ Multiple Choice: When appropriate circle the letter of the answer that correctly
answers the question or write a brief answer in your own words.
1. What is the importance of relative mass when determining the average atomic mass of
an element?
Relative mass is the mass of the individual isotopes multiplied by its abundance and tells
how much that isotope contributes to the overall mass.
2. What two things are necessary in order to calculate the average atomic mass of an
element?
mass of the isotope and percent abundance
3. The molar mass of a compound is defined as…?
the mass of one mole of the substance
4. The molar mass of C6H6 (also known as benzene) is
a. 125.00 g/mol
b. 120.00 g/mol
c. 78.00 g/mol
d. 72.00 g/mol
5. Representative particles include all of the following EXCEPT…
a. formula units
b. molecules
c. moles
d. atoms
6. A mole of tungsten contains 6.02 x 1023
a. formula units.
b. molecules.
c. moles.
d. atoms.
7. A mole of silver nitrate contains 6.02 x 1023
a. formula units.
b. molecules.
c. moles.
d. atoms.
8. It is possible to convert moles to molecules by
a. dividing by Avogadro’s number.
b. multiplying by Avogadro’s number.
c. dividing by molar mass.
d. multiplying by molar mass.
9. The percentage composition of CH4 is:
a. 75% C, 25% H
b. 92% C, 8% H
c. 20% C, 80% H
d. 25% C, 75% H
10. Which of the following is an empirical formula?
a. C6H6.
b. H2O2.
c. N2O4.
d. CH2O.
11. What is the difference between an empirical formula and a molecular formula?
Empirical – lowest whole number ratio
Molecular – true formula
12. What is the empirical formula of a compound whose molecular formula is C10H22O12?
a. CH12O2
b. C20H44O24
c. C5H11O6
d. C4H2O3
13. How are one mole of Mercury and one mole of Iron similar? How are they different? (2
points) same number of atoms; different masses
14. Match the representative particles to its’ correct partner by connecting them with a line.
(3 points)
i.
Covalent compounds
a. atoms
ii.
Ionic compounds
b. formula units
iii.
Elements
c. molecules
Problems: Solve the following problems in the space provided. Show your work for full credit!
15. The substance Teflon, the nonstick coating on many frying pans, is made from the C2F4
molecule. Calculate the number of C2F4 units present in 135 g of Teflon. (4 points)
8.13 x 1023 molecules
16. Element X has two naturally occurring isotopes. They have masses of 34.9689 amu and
36.9659 amu, with percent abundances of 75.78% and 24.22% respectively. What is the
average atomic mass of element X? Identify element X. (4 points)
35.45 amu
Chlorine
17. Fat makes up a major portion of all soaps. A fat used in many soaps is 76.5% carbon,
12.2% hydrogen, and 11.3% oxygen. What is its empirical formula? (5 points)
C9H17O
18. How many moles of phosphorus pentabromide are present in 225.7 g of the compound?
(3 points)
0.5237 mol
19. Strychnine, a deadly poison, has a molecular mass of 334 amu and a percentage
composition of 75.42% carbon, 6.63% hydrogen, 8.38% nitrogen, and the balance
oxygen. What is the molecular formula of strychnine? (6 points)
C21H22N2O2