Download Chemistry Midterm Review 2006

Chemistry Midterm Review 2012 (Thematic)
Scientists Tools
Objective 1.1 – 1.3
1. Define chemistry.
2. What is the difference between an independent variable and a dependent variable?
3. Why is a constant important in some labs?
4. What is the difference between a hypothesis and a prediction?
5. What is the difference between a theory and a law?
6. State the difference between quantitative and qualitative data.
Objective 1.4
1. The metric system is based on the power of ____________
2. What is the basic metric unit of length, mass, and volume?
3. What is the SI unit for volume? 1cm3 = 1 ______
4. Define volume. What type of lab equipment measures approximate volume?
5. What are the prefixes for the metric units and what do they equate to?
6. Using dimensional analysis, convert 157 cg into g, convert 8.6 kg into cg, convert
100m into cm , and convert 7068 mm into m
Objective 1.5: HONORS ONLY
1. What is the difference between accuracy and precision?
2. Define significant figures. What rule is used to count sig figs?
3. How many significant figures are in each: 5.730 x 108, 3000, 0.01552, 9009,
629.55, 1.777 x 10-3?
4. What is the answer expressed in proper significant figures for the following: 6.54
+ 3.053 and 8.95 x .02?
Objective 1.6
1. Place the following numbers into scientific notation
a. 34,500
b. .098
c. 670
d. .000043
2. Take the numbers of scientific notation and place in ordinary notation
a. 4.56 x 103
b. 5.78 x 10-2
c. 1.678 x 10-5
Antacids
Objective 2.1
1. Define matter and list several examples
2. Matter can be broken down into _____________________ and
______________________________.
3. Differentiate characteristics of pure substances and mixtures.
4. Pure substances can be broken down into________________ and
___________________________.
5. Differentiate characteristics of an element and a compound. State examples.
6. Mixtures can be classified as _______________________ and
__________________________
7. Differentiate characteristics of a homogeneous and heterogeneous mixture. State
examples
8. What is another name for a homogenous mixture?
9. By what means can you separate a compound? A mixture?
10. State whether each is a compound or element: Fe, CO, CaCl2, Hg, Co, argon,
sodium chloride, manganese(II) nitrate, I2.
11. What is a chemical symbol? Write the symbols for the following. Mercury, gold,
iodine, calcium, barium, tin, magnesium, phosphorus.
12. State whether the following are homogeneous or heterogeneous:
a. rocky road ice cream
b. orange juice with pulp
c. fog
d. salt solution
d. Koolaid
e. air without pollution
f. brass
g. sandy water
13. Differentiate characteristics of a solution, colloid and suspension. State examples
Objectives 2.2-2.3
1. What are the characteristics of a covalent and ionic compound?
2. What is a chemical formula?
3. _________ are used to indicate the number of atoms in a chemical formula of a
compound.
4. Which of the following compounds are ionic? H2O, Na2O, CO2, CaS2, SO2,
CaCO3.
5. Know the difference between a formula unit and a molecule.
6. Define ion. What is the difference between a cation and an anion?
7. What is the chemical formula for the following; sulfide ion, sodium ion, fluoride
ion, mercury (II) ion?
8. What are the names of the following ions; Ba2+, Al3+, O2-, and Sn4+?
9. Metals form _______ions and nonmetals form ________ ions.
10. What are binary compounds? Ternary compounds?
11. _________ are used to indicate the charge of transition elements in a ternary ionic
compound.
12. Name the following chemical formulas; KClO3, BaSO4, MgBr2, Li2CO3, CoF2,
Ag3N and Na3PO4
13. Write formulas for the following: potassium nitrate, lithium oxide, calcium
phosphate, ammonium carbonate, barium chloride, and copper(I) chloride.
14. What is a polyatomic ion? What are the names of the following ions: PO4-3, SO3-2,
CO3-2, NO2-1, NH4+1
15. Prefixes are used to indicate the number of atoms in _________ compounds.
16. Name the following covalent compounds: CH4, N2O5, SO3, CO, H2O, Cl2P3
Objectives 2.4 – 2.5
1. List 3-4 characteristics of both an acid and base
2. What is the characteristic ion of an acid? For a Base?
3. What is the ph range for an acid? For a base?
a. Which acid is stronger, a pH of 4 or a pH of 1?
b. Which base is stronger, a pH of 11 or a pH of 8?
4. What color does red litmus paper turn in an acid? In a base?
5. What color does phenolphthalein turn in an acid? In a base?
6. What are the products of neutralization?
7. Know how to predict the products and balance neutralization (double
replacement) reactions.
H2CO3 + Fe(OH)2 →
HClO4 + NaOH →
HBr + Ba(OH)2 →
What is the difference between a strong and weak acid and base? What is the
difference between a concentrated and dilute acid & base?
a. Select the strong acids in the diagram
b. Select the dilute bases in the diagram
a)
b)
c)
8.
A
B
X
H
X
H
X
H
H

H
X
X


H
M
MOH
OH 
OH 
X
MOH
X

X

M

M
OH 
OH
M
H
E

X
H
H
MOH
OH 
X
MOH
MOH
OH 
OH 
D
MOH
M
H
X

H
X
H
C
M
MOH
F
X
X
MOH
H
M
H
X

H
G
H
OH 
M
OH

M
M
OH 
M
M
OH 
M
HX
MOH
OH 

M
HX
X
MOH
X
HX
HX
H
MOH
OH 
M
OH
HX
X
OH 
OH 
M
OH 
M
OH
HX
HX
H
HX
HX

MOH
M
HX
H
HX
HX
9. Classify the substances as either Arrhenius acids or Arrhenius bases and give their
chemical name (STANDARD ONLY: only need to name a-g)
a) HCl
b) H2SO4
c) NaOH
d) HC2H3O2
e) HNO3
f) Fe(OH)2
g) H2CO3
h) HNO2
i) H3PO4
Objective 2.6
1. What is a chemical equation?
a) Identify the reactants and the products of the following equation:
Al + O2 → Al2O3
2. What is the arrow called? Where would a catalyst’s chemical formula be placed? If
Δ was above the arrow, what would that mean?
3. What are the 5 types of chemical reactions?
4. Identify the different types of chemical reactions.
Examples:
b) FeCl3 + NaOH → Fe(OH)3 + NaCl
c) Al + O2 → Al2O3
d) C2H2 + O2 → CO2 + H2O
e) Pb(NO3)2 + NaCl → PbCl2 + NaNO3
f) Na + H2O → NaOH + H2
g) KClO3 → KCl + O2
5. Know how to predict products of a chemical reaction. Predict and balance the
following equations(STANDARD ONLY complete a-g)
Example:
a) Al + N2 →
b) Li2CO3 
c) H2O →
d) K + HCl →
e) Cl2 + NaBr →
f) CaS + FeCl2 →
g) CH4 + O2 
h) SO3 + H2O →
i) Mg(ClO3)2 
j) Mg + H2O(l) 
6. What is an activity series chart? What type of reaction do you use it for?
a) Using the activity chart, why can sodium replace hydrogen?
7. What are 5 indicators of a chemical reaction?
8. List the chemical formulas for the 7 diatomic molecules.
9. Know how to translate chemical equations and balance them appropriately.
Example:
a) ammonium chloride reacts with calcium hydroxide to form calcium chloride &
nitrogen trihydride (ammonia) & water
b) sodium oxide and water yield sodium hydroxide
10. What is the precipitate for the following reaction
a. CaS + FeCl2 → CaCl2 + FeS
b. AgNO3 + PbCl2  AgCl + Pb(NO3)2
11. HONORS ONLY:Write a complete ionic and net ionic equation for the 2 examples
above.
Objective 2.7
1. What is the law of conservation of mass?
2. Differentiate between a subscript and a coefficient? Which one is used to balance
a chemical equation?
3. To satisfy the conservation of mass, we balance chemical equations.
Example:
a) __ FeCl3 + __ NaOH → __ Fe(OH)3 + __ NaCl
b) __ Al + __ O2 → __ Al2O3
c) __ C2H2 + __ O2 → __ CO2 + __ H2O
d) __ Pb(NO3)2 + __ NaCl → __ PbCl2 + __ NaNO3
e) __ Na + __ H2O → __ NaOH + __ H2
f) __ KClO3 → __ KCl + __ O2
Objective 2.8
1. Define reaction rate
2. Explain the three criteria of the collision theory.
3. Draw in the activation energy for each line. Label the pathways as “with
catalyst” & “without the catalyst”. Which line represents the faster
reaction.
4. On the pathway to the left, label the activated complex,
activation energy, reactants, products, and enthalpy released
or absorbed by the reaction. Is this reaction Endo or
Exothermic?
5. List the 5 factors that affect the rate of a reaction and why the rate changes with
each factor?
6. How does rate change if you increase the concentration of the reactants?
7. How does rate change if you increase the surface area?
8. How does rate change if you decrease the temperature?
9. How does rate change if you add a catalyst?
10. Define activation energy? Will a reaction that proceeds faster have a higher or
lower activation energy?
11. Give the following values in kilojoules:
a.
b.
c.
d.
e.
f.
g.
Potential energy of reactants
Potential energy of products
Heat of reaction(∆H)
Activation energy of forward reaction
Activation energy of the Reverse reaction
Potential energy of the activated complex
Is the overall forward reaction endo or
exothermic?
12. What are the 2 reasons why an increase in
temperature can increase the reaction rate?
Airbags
Objective 3.1
1. Define the 6 changes of state. Which ones are exothermic and which ones
are endothermic?
2. Create a chart comparing& contrasting characteristics of solids, liquids,
and gases in regards to density, compressibility, particle size, shape,
volume, kinetic energy, attractive forces and movement
3. Explain how a solid melts into a liquid using kinetic energy in your
explanation.
4. Define boiling point?
5. What is the difference between normal boiling point and boiling point?
6. What happens to vapor pressure as temperature increases?
7. What happens to boiling point as altitude increases?
8. Who has a higher boiling point; Mount McKinley or Charlotte, NC?
9. What 2 temperatures measure the same amount during a phase change of a
pure solvent?
10. Know how to read phase change graphs.
a. In what state is compound X at a temperature of
600oC and a pressure of 80 atm?
b. State the temperature and pressure at the critical
point of compound X?
c. What phase change does X undergo if a sample at
500◦C and 40 atm has its temperature changed to
100◦C?
-30
F
-40
11. Know how to read a heating or cooling curve
a. State the boiling point of the substance:
-50
b. Which segment represents an increase in the
kinetic energy of the liquid particles?
c. Which segment represents condensation?
Temperature (oC)
-60
-70
-80
E
D
-90
-100
-110
C
B
-120
-130
-140
A
-150
0
2
4
6
8
Time (min.)
10
12
14
16
d. How many minutes does it take from the time
the substance begins to vaporize until it is
completely gaseous?
Objective 3.2 – 3.4
1. What is the difference between a physical and chemical property? Give examples
2. What is the difference between an intensive and extensive property? Give
examples.
3. A chemical change is also known as a chemical _______________.
4. Name 5 buzz words that signify a physical change and 5 that signify a chemical
change.
5. Name five indicators/observations of a chemical change(reaction).
6. Classify each as a physical or chemical change: food spoils, water boils, nail
rusting, baking bread, sugar dissolving in water, tarnishing silver, acid
neutralizing a base, drying a wet towel
7. Define density. What is the equation?
8. Ice floats because it is more or less dense than water?
9. Put the 3 states of matter in order of increasing density.
10. A copper penny has a mass of 3.1 g and a volume of .35 cm3. What is the density?
11. A plastic ball has a volume of 19.7 cm3 and a density of .8029 g/cm3. What is the
mass?
12. The density of silicon is 2.33 g/cm3. What is the volume if its mass is 62.9g?
Objective 3.6
1. Know the representative particle for an ionic compound, covalent compound, element,
and a diatomic molecule?
2. Calculate the molar mass of
a) NaCl
b) C6H12O6 c) Ca3(PO4)2
3. Be able to convert moles into grams using molar mass, moles into particles using
Avogadro’s number
a) Convert 100.1 grams of HCl into moles.
b) Find the mass in grams of 365.8 moles of SO2.
c) How many moles are in 5.43 x 1023 atoms of Ca?
d) What is the mass in grams of 1.20 x 108 formula units of CuO?
4. Be able to find percent composition.
a) Find the percent composition of NH3
b) Find the percent composition for 80.0 g Ba and 32.0 g of Cl.
5. Which of the following is a molecular formula of XY3? X2Y3, XY4, X2Y5, X2Y6
6. What is the empirical formula of the following:
a) C4H16
b) P4O10
Honors Only:
7. Be able to find the empirical formula and molecular formulas.
a) What is the empirical formula of a compound that is 25.9% nitrogen and 74.1%
oxygen?
b) What is the empirical formula of a compound that has a mass of 10.150 grams and
contains 4.433 grams of P and 5.717 g of O?
c) What is the molecular formula if the empirical formula is NaO and the gram
formula mass is 78 g?
Objective 3.5 and 3.7
1. Know the 5 assumptions of the kinetic molecular theory.
2. Know the difference between an ideal gas and a real gas. What conditions does a
real gas deviate from an ideal gas.
3. What is the difference between effusion and diffusion of a gas?
4. What happens to average kinetic energy when Kelvin temperature doubles?
5. Both methane gas(CH4) and hydrogen gas(H2) are in a container at 67 C.
a. Which gas will have the greatest speed?
b. Do they have the same average kinetic energy of different average kinetic
energy?
4. Know the relationship between the variables of a gas
a) As the temperature of a gas decreases, the volume of a gas will __________.
b) As the temperature of a gas decreases, the pressure of the gas will _________.
c) As the volume of the gas decreases, the pressure of the gas will ___________.
d) As the moles of the gas decreases, the pressure of the gas will ____________.
e) As the volume of the gas decreases, the moles of the gas will ____________.
6. Define pressure. What are some common pressure units?
7. Know how to convert pressure units:
a) convert .200 atm to mmHg
b) convert 345.8 kPa to atm
c) convert 760 mmHg to kPa
8. What is standard temperature and standard pressure?
9. Know how to solve problems using Boyle’s law, Charles law, Gay-Lussac,
Combined, Ideal, Density and Molar mass using the Ideal gas law and Dalton’s
law of partial pressure.
a) A gas occupies a volume of 200. ml at 100. mmHg. What volume will the gas
occupy at 300. mmHg?
b) Air has a total pressure of 20.6 atm and contains carbon monoxide, oxygen, and
nitrogen. If air is made up of 0.6 atm of carbon monoxide, 12.6 atm of oxygen,
what would be the partial pressure of nitrogen?
c) If a sample of gas occupies 15.9 L at 34 C, what will its volume be at 27 C if the
pressure does not change?
d) The volume of a sample of oxygen gas is 300.0 ml when the pressure is 1.00 atm
and the temperature is 27.0 C. At what temperature would the volume change to
1.00 L and the pressure change to 0.500 atm?
e) A sample of gas at 25.0 C has a volume of 11.0 L and exerts a pressure of
660.0 mmHg. How many moles of gas are in the sample?
f) A sample of gas in a closed container at a temperature of 100. C and a pressure of
3.0 atm is heated to 300. C. What pressure does the gas exert at the higher
temperature?
g) How many grams of N2 are in a flask with a volume of 250. ml at a pressure of
3.0 atm and a temperature of 300 K?
h) What is the density of silicon tetrafluoride gas at 72 C and a pressure of 144.5
kPa?
10. Know how to solve for the pressure of a gas if it is collected over water.
a) A sample of nitrogen gas is collected over water at 23 C with a vapor pressure of
.0278 atm. What is the pressure of the nitrogen gas if the atmospheric pressure is
785 mmHg?
10.Use the law of combining volumes, Avogadro’s law, and molar volume to solve
these problems.
a) 3O2 → 2O3 Both gases are measured at the same temperature and pressure. How
many O3 molecules are formed from the reaction of 24 O2 molecules?
b) How many moles of O2 are required to make 24 moles of O3?
c) How many liters of O3 are formed from 12 L of O2?