Download Chemistry Midterm Review 2006

Chemistry Midterm Review 2012
(Thematic)
Scientists Tools
Objective 1.1 – 1.2a
1. State the difference between quantitative and qualitative data and list an example
of each.
2. What is the difference between accuracy and precision?
3. Record the following measurements using the instruments seen below.
Objective 1.2a – 1.2b: HONORS ONLY
1. Define significant figures. What are the 4 major rules used to count sig figs?
2. How many significant figures are in each:
a. 5.730 x 108
b. 3000
c. 0.01552
d. 9009
e. 629.55
f. 1.777 x 10-3?
3. What is the answer expressed in proper significant figures for the following:
a. 6.54 + 3.053
b. 8.95 x .02
Objective 1.3
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 = ______ mL
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
a. 157 cg into g
b. 8.6 kg into cg
c. 100m into cm
d. 7068 mm into m
Objective 1.4
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
of each.
Objectives 2.2
1. What kind of elements form ionic compounds? What kind of elements form
covalent(molecular) compounds?
2. Which of the following compounds are ionic? H2O, Na2O, CO2, CaS2, SO2,
CaCO3.
3. Know the difference between a formula unit and a molecule.
4. Define ion. What is the difference between a cation and an anion?
5. What are the names of the following ions; Ba2+, Al3+, O2-, and Sn4+?
6. Metals form _______ions and nonmetals form ________ ions.
7. What are binary compounds? Ternary compounds?
8. _________ are used to indicate the charge of transition elements in a ternary ionic
compound.
9. Name the following chemical formulas:
a. KClO3
b. BaSO4
c. MgBr2
d. Li2CO3
e. CoF2
f. Ag3N
g. Na3PO4
10. What is a polyatomic ion? What are the names of the following ions: PO4-3, SO3-2,
CO3-2, NO2-1, NH4+1
11. Prefixes are used to indicate the number of atoms in _________ compounds.
12. Name the following covalent compounds:
a. CH4
b. N2O5
c. SO3
d. CO
e. H2O
f. Cl2P3
Objectives 2.3
1. What is a chemical formula?
2. _________ are used to indicate the number of atoms in a chemical formula of a
compound.
3. What is the chemical formula for the following:
a. sulfide ion
b. sodium ion
c. fluoride ion d. mercury (II) ion
4. Write formulas for the following:
a. potassium nitrate
b. lithium oxide
d. ammonium carbonate
e. barium chloride
c. calcium phosphate
f. copper(I) chloride
5. Write the formulas for the following compounds:
a. diphosphorus pentoxide
b. hexasulfur trichloride
c. sulfur triiodide
Objectives 2.4
1. 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
d. HNO3
e. Fe(OH)2
f. H2CO3
g. HNO2
h. H3PO4
Objectives 2.5a
1.
2.
3.
List 3-4 characteristics of both an acid and base.
What is the characteristic ion of an acid? For a Base?
What is the ph range for an acid? For a base?
a. Which pH is more acidic?, pH of 5 or pH of 1? By factor is the hydrogen ion concentration
greater in pH =1
b. Which ph is more basic? pH of 11 or pH of 8? By factor is the hydroxide ion concentration
greater in pH =11
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 acid-base neutralization?
7. Know how to predict the products and balance neutralization (double replacement) reactions.
a) H2CO3 + Fe(OH)2 →
b) HClO4 + NaOH →
c) HBr + Ba(OH)2 →
8. What is the difference between a strong and weak acid and base?
9. What is the difference between a concentrated and dilute acid & base?
10. Using the diagram below,
a. Select the strong acids in the diagram
b. Select the dilute bases in the diagram
A
B
X
H
H
X
H
H

X
H
H
E
H
M
OH 
OH 
H
X

X
X
X
H
C
X
X

MOH
X
H
M

M
M
OH 
OH 
M
F
H
H
MOH
OH 
X
MOH
MOH
OH 
OH 
X
MOH
M
MOH

D
MOH
G
X
X
MOH
M
H

H
X
H
H
OH 
M
OH 
M
OH

M

M
M

OH

OH 
OH
H
OH
M
OH 

MOH
X
MOH
X
HX
HX
OH 
HX
H
HX

HX
M
HX
MOH
HX
OH 
OH 
M
HX
M

HX
X
MOH
M
M
HX
H
HX
HX
Objective 2.5b
1. Know how to calculate the pH from hydrogen and hydroxide ion concentrations
a) What is the pH of a [OH-] = 1.0 x 10-5 M?
b) What is the pH of a [H+] = 1.0 x 10-5 M?
c) What is the pOH of a [H+] = 1.0 x 10-1 M?
d) What is the pOH of a [OH-] = 1.0 x 10-12 M?
e) What is the [H+] when the pH = 8.0
f) What is the [OH-] when the pH = 8.0
2. a. What is the hydrogen ion concentration of 0.001 M HNO3? b. What is the
[OH-]?
Honors Only:
3. What is the hydrogen ion concentration of [OH-] = 3.0 x 10-2 M? What is the pH?
4. What is the pH of a solution if the [H+] = 3.4 x 10-5 M? What is the hydroxide
concentration?
5. Determine the pH of a 2.0 x 10-2 M Sr(OH)2?
6. The pH of a solution is measured and determined to be 7.52? What is the
hydrogen ion concentration? Is the solution acidic or basic?
Objective 2.5c
1. What piece of glassware holds the analyte? What piece of glassware holds the
titrant?
2. What is the difference between end point and equivalence point in a titration?
3. Why is an indicator added to the analyte?
4. In a titration, how much of .15 M NaOH is needed to neutralize 20.0 ml of .500M
HCl solution?
5. In a titration, what is the molarity of HNO3 if 25.0 ml of it neutralized 15.0 ml of
.60M Ca(OH)2?
Objective 2.6 & 2.7
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:
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
5. What are 5 indicators of a chemical reaction?
6. List the chemical formulas for the 7 diatomic molecules.
7. 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
8. What is the law of conservation of mass?
9. Differentiate between a subscript and a coefficient? Which one is used to balance a
chemical equation?
10. 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. 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) 
2. 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?
3. What is the precipitate for the following reaction?
a. CaS + FeCl2 → CaCl2 + FeS
b. AgNO3 + PbCl2  AgCl + Pb(NO3)2
Objective 2.9
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?
4. 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 – 3.3
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. (Honors Only)
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:
a. food spoils
b. water boils
c. nail rusting d. baking bread
e. sugar dissolving in water
f. tarnishing silver
g. acid neutralizing a base
h. 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 3.4
1. What is the sign of an endothermic reaction and exothermic reaction?
2. Using the specific heat values for water and iron from reference sheet, which one
would have the largest temperature change if they have the same mass?
3. Calculate the heat released or absorbed during a physical change.
a. Calculate the heat absorbed when 15.0 g of ice melts to liquid. See
reference sheet for Hfus
b. Calculate the heat released when 75.4 g of vapor condenses into liquid.
See reference sheet for Hvap
4. Calculate the heat released or absorbed in a chemical reaction?
a) What is the specific heat of a metal that releases 2500 J of energy. The
metal has a mass of 25.0 g and had a temperature change of 5.0°C.
b) How much heat is released when iron is dropped in a beaker of water.
The mass of the metal was 43.0 g and the initial temperature of the
metal was 78.0°C. The water temperature changed from 25 C to 32
°C. The specific heat of the metal is .45 J/g°C.
c) What is the amount of heat absorbed by water if 23.4 g of water is
heated from 34.0°C to 78.0° C. See reference sheet for specific heat of
water.
Objective 3.5
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?
Objective 3.6
1. 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.
2. Which of the following is a molecular formula of XY3? X2Y3, XY4, X2Y5, X2Y6
3. What is the empirical formula of the following:
a) C4H16
b) P4O10
4. 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?
Honors Only:
c) What is the molecular formula if the empirical formula is NaO and the gram
formula mass is 78 g?
Objective 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 ____________.
Objective 3.8-3.9
Define pressure. What are some common pressure units?
1. 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
2. What is standard temperature and standard pressure?
3. 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?
4. 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?