Chemistry Midterm Exam Review Sheet Spring 2012
–
1. Know your safety rules
2. A shopping mall wanted to determine whether the more expensive “Tough Stuff” floor wax was better
than the cheaper “Steel Seal” floor wax at protecting its floor tiles against scratches. One liter of each
brand of floor wax was applied to each of 5 test sections of the main hall of the mall. The test sections
were all the same size and were covered with the same kind of tiles. Five (5) other test sections received
no wax. After 3 weeks, the number of scratches in each of the test sections was counted.
Identify the following components of an experiment.
a.
Problem: The Effect of the (changes in the independent variable) on the (dependent variable) –
write it using this format
What is the effect of the “tough stuff” floor wax on the numberof
scratches on the floor?
b.
Hypothesis: If the (independent variable – describe how it will be changed), then the (dependent
variable – describe the effect). -- write it using this format If “tough stuff” is used then there will be
less scratches on the floor compared to “steel seal” or no wax at all.
c.
Independent variable - type of wax
d.
Dependent variable – number of scratches
e.
Control - - five tiles with no wax
f.
Repeated trials – used 5 different tiles for each type of wax and no wax
g.
Constants – used one liter of each substance, used same size tiles, tiles were in the main
hallway
1. State the law of conservation of mass. Matter/Mass cannot be created nor destroyed, or the mass
of substances produced (products) by a chemical reaction is always equal to the mass of the
reacting substances (reactants)
Atoms of reactants = Atoms of products; Mass of reactants = Mass of products
2H2 + O2  2H2O
2. Distinguish between the following word pairs (terms):
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a. homogeneous mixture and heterogeneous mixture – homogenous mixture is the same
throughout, heterogeneous mixture has distinguishable parts
b. element and compound – element cannot be separated, compound is 2 or more elements. a
compound can be separated chemically
3. Identify the following as a heterogeneous mixture, compound, element, or homogeneous mixture:
a. salt water
b. gold
c. water
d. Italian dressing
homogeneous
element
compound
heterogeneous
4. Distinguish between a physical and a chemical change.
Physical change is where a substance does not become something new (example tearing a piece of
paper) the product of the change is still the original substance with the same properties
(reversible)
Chemical change is where a substance does change into something new (breaking down water into
oxygen and hydrogen) products are new with different properties than the original substance.
5. Identify the following changes as physical or chemical:
a. water freezing - physical
b. paper burning - chemical
c. wax melting - physical
d. metal rusting - chemical
6. Identify the following properties as physical or chemical:
density - physical
b. flammability - chemical
c. reactivity – chemical (reacts with)
d. temperature - physical
7. Distinguish between an endothermic and an exothermic reaction.
Endothermic reactions absorb energy (feels cold to the touch)
Exothermic reactions release energy (feels hot to the touch)
8. List the six phase changes and state if they are endothermic or exothermic.
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Melting, sublimation and vaporization – endothermic
Freezing, deposition and condensation - exothermic
9. How do intensive properties differ from extensive properties?
An intensive property remains the same no matter how much of a substance is present, while an
extensive property is dependent upon the amount of substance present.
10. Label each of the following properties as intensive or extensive.
a. mass
b. density
c. length
d. melting point
extensive
intensive
extensive
intensive
e. volume
f. boiling point
g. color
extensive
intensive
intensive
13. Compounds can be separated __chemically_. Mixtures can be separated physically. List 3 separation
techniques that can be used to separate mixtures and explain how they work.
Filtration – separates an undissolved solid from a liquid (sand & water)
Distillation – separates two liquids based on difference of boiling points (alcohol and water)
Evaporation – separates a dissolved solid from a liquid (salt water)
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1. How does qualitative data differ from quantitative data? Give an example of each.
Qualitatitve – information describing some physical characteristic (color, odor, shape)
Quantitative – numerical information describing how much, how little, how big, how fast, etc.,
(mass, height, volume, speed)
2. What is the percent error equation? (experimental value – actual value) /actual value x 100
3. A student measured a piece of steel pipe and found its length to be 5.4 m. The accepted length of the steel
pipe is 5.1 m. Calculate the percent error.
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4. Convert the following temperatures.
a. 38oC = _______311______ K
b. 543 K = ______270_______oC
c. 50°C is equal to __50 + 273 = 323 Kelvin.
5. You report four measurements of mass, 7.52 g, 7.50 g, 7.51 g, 7.50 g. If the correct measurement is 7.85 g,
are these measurements accurate? __No___ Are they precise? _Yes___ Explain your answer.
Not accurate because the measurements are not close to the known value, but they are precise because the
measurements are close to each other.
4. Describe how you add or subtract numbers using significant figures?
When you add and subtract you have to report your answers in the least number of decimal places.
5. Describe how you multiply or divide numbers using significant figures?
When you multiply and divide you have to report your answers in the least number of significant figures.
6. Identify the number of significant figures in the following measurements:
a. 1.00 g 3 trailing 0s count when a decimal is present
c. 0.004500 g 4 for reason in a and b above
b. 0.0035 g 2 leading 0s never count
d. 101 g 3 middle 0s always count
7. Perform the following operations. Keep the correct number of significant figures.
a. 4.56 cm x 9.1 cm = 41 cm
c. 7.1 + 6.00 =
2
b. 20.0 ÷ 4 = 5
d. 8.0 – 4 = 4
13.1
8. List the appropriate SI unit for the following measurements. Define temperature and mass.
a. Length
b. mass
c. volume
meter
kilogram
Liter
d. temperature
e. time
f. quantity
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Kelvin
second
mole
Mass – amount of matter in a substance
Temperature – measure of the average kinetic energy
9. Perform the following unit conversions: Use the factor label (dimensional analysis) method.
a. 44 mm = _____4.4 x 10-5____ km
b. 15 lbs = __6.8 x 106_______ mg (1lb = 454 g)
c. 5.0 x 106 sec = __58___ days
d. 14,110 ft = ____4301___ m (1m = 1.0936 yards)
10. How should all equipment be read? Read the following balance to the appropriate precision.
All known values and one estimate. 373.35g
11. List the equipment that can be used to measure volume.
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Ruler for regular solid
Graduated cylinder (liquids & water displacement)
12. Read the following graduated cylinder to the appropriate precision.
1.25 mL
13. Describe how you could determine the volume of an irregularly shaped object.
Use water displacement – fill a graduated cylinder to a given value (record this measurement)
place the object in the graduated cylinder then record the new value. Subtract the original value
from the new value and that is equal to the volume of the irregular object
14. What is the formula for density? Density = mass ÷ volume
15. What is the density of an unknown substance if the mass is 100g and the volume is 100 ml? 1 g/mL
16. What is the volume of an object if its density is 1.0 g/cm 3 and the mass is 150 g? 150 mL
17. What is the mass of an object if its density is 4.5 g/ml and the volume is 75 ml? 340 g
18. Given two substances, how do you know which one will float?
The substance with the lower density will float.
19. What type of physical property is density, intensive or extensive? What does that mean?
Density is an intensive property which means that it will not change as long as the substance is the same and
the state (phase) doesn’t change.
1. What is an atom? – smallest part of substance that still has the properties of that substance
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2. List the 2 regions of an atom. What do they give to the atom? – nucleus gives mass to the atom and
electron cloud gives size to the atom
3. List the 3 subatomic particles, their charges, their relative masses and their location in the atom.
Proton (+) nucleus, 1 amu
neutron (o) nucleus, 1 amu
electron (-) electron cloud, 0 (1/1840) amu
4. Who discovered the nucleus in the famous “gold foil” experiment? What conclusions were drawn from this
experiment? Rutherford He concluded that the atom was mostly empty space and had a small very dense
center with a positive charge.
5. Who proposed the first atomic theory in the 1800’s? What has been changed? Dalton
He said that the atom was indivisible. But that has been proven wrong with the existence of protons,
electrons and neutrons.
He said that all atoms of the same element are the same. But that has been proven wrong with the existence
of isotopes.
6. Who discovered the electron? - Thomson
7. Which number identifies the element? – atomic number
8. What is the mass number equal to? Mass number = protons + neutrons
9. How are the numbers of neutrons determined? Neutrons = mass number – atomic number(protons)
10. How are the numbers of electrons determined in a neutral atom? Neutral atom means no charge the two
particles that have charge are protons and electrons. Therefore in a neutral atom #electrons=#protons.
11. How are the numbers of electrons determined in a cation? Charge = #protons- #electrons to find electrons =
#protons - charge. So for Ca+2 #electons = protons – charge = 20 – (+2) = 18. Or for instance given Ca+2
you can say calcium has two more protons than electrons. #protons = 20 therefore #electrons = 18.
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12. How are the numbers of electrons determined in an anion? Charge = #protons- #electrons to find electrons
= #protons - charge. So for O-2 #electrons = 8 – (-2) = 10. Or for instance given O-2 you can say oxygen
has two more electrons than protons. #protons = 8 therefore #electrons = 10.
13. Besides the atomic number, what other number is found on the periodic table. What is it and give its
definition. Atomic mass number is found on the periodic table. Atomic mass number is the weighted
average of all the isotopes of an element. It can be calculated by
isotope 1
+
isotope 2
+ … = average atomic mass
14. Define isotope. Given the following nuclide, explain all the numbers.
+3
. Tell how you would find
the number of protons, neutrons and electrons and then do so.
Isotopes have the same number of protons but different number of neutrons, so different mass number. 58 =
mass number; 26 = atomic number; +3 is the charge. #protons = atomic number = 26; #electrons = #protons
minus charge = 26 - +3 = 23; #neutrons = mass number – atomic number = 58 – 26 = 32.
15. Calculate the average atomic mass for element Z, if it has the following isotopes:
Z-56
Z-57
25%
75%
55.98 amu
56.99 amu
Using equation in #13 above:
isotope 1 +
isotope 2 = 56.75 amu
16. Define the following
a. ion – atom or group of atom with a charge
b. cation – a positive ion (protons+>electrons-)
c. anion – a negative ion (electrons->protons+)
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17. Complete the following chart:
Element
Atomic # Mass # # Protons # Electrons
Carbon-14
Chlorine-36
6
14
6
6
17
36
17
18
# Neutrons
Symbol (Nuclide)
8
19
-1
Uranium-238
92
238
92
92
146
Magnesium-24
12
24
12
10
12
+2
PERIODIC TABLE
1. The periodic table is divided into three main categories. Name them and give their location on the periodic
table and their properties.
Metals – found to the left of the stair step on the periodic table, these elements are electron donors,
have a metallic luster, are good conductors, malleable, ductile, typically a solid at room
temperature
Nonmetals – found to the right of the stair step, these elements are electron acceptors, are
dull/earthy, brittle solids or gases at room temperature, poor conductors
Metalloids - found along the stairstep (except aluminum), have properties of metals and nonmetals,
used as semi-conductors
2. Mendeleev arranged the periodic table by increasing atomic _mass number__. Moseley arranged the
modern periodic table by increasing atomic _number__.
3. Distinguish between groups and periods. Groups are the columns (vertical) on the periodic table where
the elements have similar chemical properties due to the number of valence electrons Periods are the
rows (horizontal) on the periodic table where chemical properties differ as you move across
4. Which of the following elements have similar properties? Why?
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Ca, K, Kr, Cl, Br
Cl and Br have similar properties because they are in the same column and have same number of
valence electrons,
5. Give the names of the following groups.
a. Group 1
alkali metals
b. Group17
halogens
c. Group 2
alkaline earth metals
d. Group 18
noble gases
6. Where are the transition metals? Groups 3-12 Inner transition metals? Bottom two rows
7. Which Group contains the most reactive metals? Alkali metals Reactive nonmetals? Halogens
8. Give the group and period for the following elements:
a. helium –
18(noble gas), 1
b. bromine - 17(halogen), 4
c. calcium -
2(alkaline earth metal), 4
d. copper – 11(transition metal), 3
9. Identify the number of valence electrons present in the following elements:
a. Radium
2
b. Iodine
c. Cesium
7
1
d. Aluminum
e. vanadium
3
10. Draw an electron dot structure for the following then predict the oxidation number or charge:
a. Nitrogen
b. argon
c. magnesium
d. sodium
e. aluminum
f. sulfur
g. silicon
h. fluorine
11. Which group of elements on the periodic table is most likely to donate one electron? Alkali metals
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2
UNIT 4 IONIC BONDING AND ACID NAMING
1. Which elements lose electrons? Metals Gain electrons? Nonmetals Why? To become more stable obey the
octet rule
2. Why do ions have a + or – charge? + charge because a metal loses electrons so p+>e- Charge because a nonmetal gains electrons so e->p+
3. How many electrons are lost or gained if an ion has the following charge:
a. +1
b. +2
c. +3
d. +4
1 electron lost
2 electrons lost
3 electrons lost
4 electrons lost
e. –4
f. –3
g. –2
h. –1
4 electrons gained
3 electrons gained
2 electrons gained
1 electron gained
4. What elements are “Stable” (lowest energy state)? Noble gases
5. How many valence electrons are needed for an element to be considered stable? ___8______ This known
as the _octet___ Rule.
6. Name the following ionic compounds:
a. CaBr2 calcium bromide
e. SnBr4 tin (IV) bromide
b. KF
f.
potassium fluoride
Na2CO3 sodium carbonate
c. Mn(OH)2 manganese (II) hydroxide
g. Cu2O copper (I) oxide
d. BeS
h.
beryllium sulfide
PbCl2 lead (II) chloride
7. Write the formulas for the following ionic compounds.
a. Sodium fluoride NaF
d. Iron (III) sulfate Fe2(SO4)3
b. Calcium nitrate
e. Sodium bicarbonate NaHCO3
c.
Ca(NO3)2
Magnesium nitride Mg3N2
f. Lead (IV) Phosphate Pb3(PO4)4
8. List the rules for naming acids.
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Determine if Binary – 2 capital letters If binary name like normal for instance HCl hydrogen chloride
forget the gen and change the –ide to –ic and make one big word then add the word acid
Hydrochloric acid
Determine if Ternary – more than 2 capital letters. Put a box around everything but the H, then name that
ion (found on periodic table). If the ion ends in –ate then change to –ic. If the ion ends in –ite then change
to –ous. When you ate to much you get ic and when ite gets married it becomes ous.
9. Write the formulas for the following acids.
chloric acid
phosphorous
sulfuric acid
_HClO3__________________
acid
_H3PO4__________________
__H2SO4_________________
10. Name the following acids.
HCl
__hydrochloric acid_______________
H2S
___hydrosulfuric acid______________
HNO2
____nitrous acid_____________
11. Identify the properties of ionic compounds.
Conduct electricity when dissolved in water (aqueous) or melted
High melting point and boiling point
12. Draw the Lewis Structures for the following ionic compounds: sodium chloride, magnesium fluoride,
Lithium oxide, Aluminum sulfide
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UNIT 5 – CHEMICAL QUANTITIES
1. What is the value of Avogadro’s number?
6.02 x 1023 representative particles/ 1 mole
2. How many moles of chromate ions are in 5.50 moles of Pb(CrO4)2 ?
2 moles CrO42-
5.50 moles Pb(CrO4)2
= 11.0 moles CrO42-
1 mol Pb(CrO4)2
3. Work the following problems.
a. Calculate the mass of 3.57 moles of aluminum.
3.57 mol Al
27.0 g Al
= 96.4 g Al
1 mol Al
b. How many moles are in 25.0 g of Fe2O3?
25.0 g Fe2O3
1 mol Fe2O3
= 0.156 mol Fe2O3
160 g Fe2O3
c. Calculate the number atoms in 2.50 moles of Zinc. (moles to particles) representative particles =
molecules, atoms, ions, formula units
2.50 mol Zn
6.02 x 1023 atoms Zn = 1.51 x 1024 atoms Zn
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1 mol Zn
d. Determine the number of moles for 3.58 x 1023 formula units ZnCl2?
3.58 x 1023 formula units ZnCl2
1 mol ZnCl2
= 0.595 mol ZnCl2
23
6.02 x 10 formula units ZnCl2
e. Determine the number of molecules in 25.34 g of CO?
25.34 g CO 1 mol CO
6.02 x 1023 molecules CO = 5.448 x 1023 molecules CO
28.0 g CO
1 mol CO
f. Determine the mass in 6.11 x 1025 atoms of sulfur?
6.11 x 1025 atoms S
1 mole S
6.02 x 1023 atoms S
32.1 g S = 3250 g
1 mole S
4. Calculate molar mass for the following.
a. AlBr3
(1x27.0) + (3 x 79.9) = 267 g/mole
b. H2O2
c. Mg3(PO4)2
(2x1.01) + (2x16.0) = 34.0 g/mole
(3x24.3) + (2x31.0) + ( 8x16.0) = 263 g/mole
5. Caffeine, a stimulant found in coffee, has the chemical formula C8H10N4O2. Determine the percent
composition of caffeine.
C 8x12.0 = 96.0
%C = 96.0/194 x 100 = 49.5%
H 10x1.01 = 10.1 %H = 10.1/194 x 100 = 5.21%
N 4x14.0 = 56.0 %N = 56.0/194 x 100 = 28.9%
O 2x16.0 = 32.0 %O = 32.0/194 x 100 = 16.5%
Total = 194 g/mol
6. Compare molecular and empirical formulas. Be able to determine which compounds have the same
empirical formula.
Molecular – actual atom combination within a formula
Empirical – lowest whole number ratio of atoms of a formula
7. Determine the molecular formula for ibuprofen, a common headache remedy. Analysis of ibuprofen yields a
molar mass of 206 g/mol and a percent composition of 75.7% C, 8.80% H an d15.5% O. (find the empirical
formula first)
75.7 g C = 6.3 mol
÷ .96 = 6.56 x 2 = 13
14
12.01
8.80 g H = 8.71 mol ÷ .96 = 9.07 = 9 x 2 = 18
1.01
15.5 g O = 0.96 mol ÷ .96 = 1 x 2 = 2
16g/mol
molar mass of O = 16
molar mass of Empirical Formula = 206.31 g/mol
n = 206 g/mol
= .998 = 1
206.31 g/mol
Empirical Formula 1 (C13H18O2) = C13H18O2
Empirical Formula C13H18O2
UNIT 6 – CHEMICAL REACTIONS
1. List the 7 diatomic elements.
Br I N Cl H O F
2 2 2 2 2 2 2
2. Define reactants and products and know where they are located.
Reactants – starting substances located to the left of the arrow
Products – substances formed located to the right of the arrow
3. List the types of reactions and how to identify them.
Synthesis – 1 Product
Decomposition – 1 Reactant
Combustion - + ___O2 as a reactant
Single Replacement (also known as Redox) – E + C → C + E
Double Replacement – C + C→ C + C
4. Know the symbols for all the states including aqueous and define aqueous.
(s) = solid
(l) = liquid
(g) = gas
(aq) = aqueous (dissolved in water)
5. Write and balance the following chemical equations, showing the states, and identify the type. Can be
more than one.
a. Solid calcium reacts with solid sulfur to produce solid calcium sulfide.
Synthesis
Ca(s) + S(s) → CaS(s)
b. Solid aluminum metal reacts with aqueous zinc chloride to produce solid zinc metal and aqueous
aluminum chloride.
Single
2 Al(s) + 3 ZnCl2(aq) → 3 Zn(s) + 2 AlCl3(aq)
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and Redox
c. Aqueous aluminum sulfate reacts with aqueous barium hydroxide to produce solid aluminum
hydroxide and solid barium sulfate.
Double
Al2(SO4)3 (aq) + 3Ba(OH)2 (aq) → 2Al(OH)3 (s) + 3BaSO4 (s)
d. Propane gas (C3H8) reacts with oxygen gas to form carbon dioxide gas (CO2) and water vapor.
Combustion C3H8 (g) + 5 O2 (g) → 3 CO2 (g) + 4 H2O (g)
e. Solid silver oxide decomposes to produce solid silver and oxygen gas
Decomposition 2Ag2O (s) → 4Ag (s) + O2 (g)
2. Mixed Naming & Formula Practice:
Name
Formula
Formula
Name
copper (II) hydroxide
Cu(OH)2
FePO4
Iron(III) phosphate
silver chloride
AgCl
K2CO3
Potassium carbonate
hypochlorous acid
HClO
HClO3
Chloric acid
iron (II) sulfate
FeSO4
Cd(NO3)2
Cadmium nitrate
ammonium hydrogen sulfate
NH4HSO4
Cs2C2O4
Cesium oxalate
iron (II) oxide
FeO
Na3N
Sodium nitride
ammonium sulfide
(NH4)2S
SnI4
Tin(IV)iodide
magnesium phosphate
Mg3(PO4)2
HClO3
Chloric acid
nickel (II) bicarbonate
Ni(HCO3)2
SrSO3
Strontium sulfite
potassium dichromate
K2Cr2O7
HNO2
Nitrous acid
phosphoric acid
H3PO4
Ba(OH)2
Barium hydroxide
cobalt (III) nitrate
Co(NO3)3
H3PO3
Phosphorous acid
tin (IV) bromide
SnBr4
H2S
Hydrosulfuric acid
sodium hydrogen sulfate
NaHSO4
PbO
Lead (II) oxide
silver chromate
Ag2CrO4
HNO2
Nitrous acid
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