Download Semester 1 Study Guide – Chemistry

Semester 1 Study Guide – Chemistry
Chapter 2 – Matter
1.
Classify each of the following as a physical change (P) or a chemical change (C)
a. Cooking an egg _______
b. Boiling water _______
c. Ironing a shirt _______
d. Burning gasoline _______
e. Decomposing water _______
f. Evaporating alcohol _______
g. Sanding a table top _______
h. Grinding grain ________
i. Fermenting fruit juice _______
j. Dissolving sugar in water _______
2.
Classify each of the following as an element (E), compound (C), a homogenous mixture (H) or a
heterogeneous mixture (HM)
a. Table salt _______
b. Chlorine gas _________
c. Sand in water ________
d. Petroleum __________
e. Caffeine _________
3.
Classify each of the following as an element (E), a compound (C), or a mixture (M).
a. 14K gold __________
b. Pure silver __________
c. Aluminum __________
d. Distilled water __________
e. Tap water _________
f. Brass __________
g. Tungsten _________
h. Sodium chloride __________
i. Air __________
Chapter 3 – Chemical Foundations: Elements, Atoms and Ions
1.
Give the symbol for the following elements:
a. Lead __________
b. Iron ____________
c. Silver ___________
2.
Give the names that correspond to the following:
a. Hg ______________________
b. K ________________________
c. I _________________________
3.
Four of the most important elements on the planet are hydrogen, oxygen, carbon and nitrogen.
Give the symbol for each of these elements
a. Hydrogen _____________
c. Carbon ______________
b. Oxygen _____________
d. Nitrogen ______________
4.
Classify the following as a metal, nonmetal or metalloid.
a. Arsenic _____________________
b. Argon ______________________
c. Calcium ______________________
d. Plutonium _____________________
5.
For the element Sulfur, S, give the following,
a. Atomic number _____________
b. Mass number ______________
c. Number of neutrons ______________
d. Number of electrons ______________
e. Number of protons _______________
6.
For the ion boron,
Br, give the following
a. Atomic number _________________
b. Mass number ________________
c. Number of neutrons _________________
d. Number of electrons _________________
e. Number of protons ___________________
Chapter 4 – Nomenclature
1.
The name for ClO4- is ____________________________________
2.
The correct name for Sn4+ is _____________________________
3.
The name for the Fe2+ species is ____________________________
4.
The name for CO32- is ______________________________
5.
The name for NO3- is ____________________________
6.
The name for NH4F is _________________________________
7.
The name for the compound Sn(NO3)2 is ______________________________
8.
The name for the compound Fe2O3 is ______________________________
9.
The name for SrCl2 is _________________________________
10. The name for HCl is ________________________________
11. The name for CuCl is _________________________________
12. The name for CF4 is ________________________________
13. The name for N2O5 is __________________________________
14. The name for (NH4)2CO3 is ___________________________________
15. The name for the compound (NH4)2S is ______________________________
16. The name for the compound K2CrO4 is ________________________________
Fill in the blank with the corresponding formula:
17. Phosphorus trichloride _____________________________________
18. hydrobromic acid _______________________________________
19. magnesium sulfite _______________________________________
20. nitrogen trichloride ___________________________________
21. sodium carbonate ____________________________________
22. lead (II) acetate ________________________________
Review for Chapter 5 – 8
1.
Convert these numbers into scientific notation:
125000 =
2.
0.003405 =
Convert these numbers into decimals:
4.26x104 =
3.
3.05x10-3 =
Convert the metric measurements:
2.5 m = _____________cm
0.765 kL = ___________________ mL
650 mg = ______________cg
750 mL = __________________cm3
4.
In the metric system, what are the fundamental units of mass, length, temperature and volume?
5.
How many significant figures are in each of these measurements?
230 cm
6.
0.027 m
0.12458 =
6.022x1023 =
Solve each of these calculations and round to the correct number of significant figures.
2.035 + 7.75 =
8.
3.05x108 kg
Round each of the numbers to 3 significant figures.
357,500,000 =
7.
4.50 mm
(125.0)(12.7) =
Which contains more atoms, 1.0 mol of Na or 2.0 mol of He?
(3.5x103)(2.62x10-4) =
9.
Which weighs more, 1.0 mol of Na or 2.0 mol of He?
10. Write the empirical formula for each compound:
C3H12
H2SO4
H2O4
11. Convert 345 kg into ounces using the factor-label method:
(2.20 lbs = 1.00 kg & 16 oz. = 1 lb)
12. Calculate the density of each object and determine if it will sink or float in water.
a. A block that weighs 4.0g and occupies 3.4cm3
b.
A sphere that has a volume of 3.14 mL and a mass of 2.78 g
13. How much does 1.00 mol of NaOH weigh?
14. What is the molar mass of Al(NO3)3?
15. Calculate the mass of 0.35 moles of BaF2:
16. How many atoms are there in 2.50 moles of argon?
17. A 45.0g sample of water contains how many water molecules?
18. What is the mass percent of lithium in Li2O?
19. Add coefficients to balance the following equations:
_____CaCO3  ____CaO + _____CO2
_____Al(OH)3 + _____ HCl  ___AlCl3 + ___H2O
_____Sb4O6 + _____C  _____Sb + _____CO
_____Fe2O3 + _____C  _____Fe + _____CO2
_____ KO2 + _____H2O + _____CO2  _____KHCO3 + _____O2
20. Write a balanced equation for each reaction described below. Please include the states of matter
for each:
a. Solid silver oxide decomposes into silver metal and oxygen gas
b.
Solutions of potassium chromate and barium chloride are mixed to produce a bright
yellow precipitate (barium chromate) and potassium chloride which stays in the solution:
21. Complete and balance the following reactions:
a.
AgNO3 (aq) + KOH (aq) 
b.
Na2S (aq) + CoCl2 (aq) 
22. There are several classifications listed below. Identify which classifications apply to each
equation. More than one classification may apply to each equation:
Oxidation-reduction
Decomposition
Acid-base
Synthesis
Precipitation
Combustion
a.
HCl + NaOH  H2O + NaCl
b.
CH4 + 2O2  CO2 + 2H2O
c.
2H2O2  2H2O + O2
d. NaCl + AgNO3  NaNO3 + AgCl
Chapter 9 – Chemical Quantities
1.
The balanced equation P4 (s) + 6H2 (g)  4PH3 (g) tells us that 3 mol H2:
a. Reacts with 1 mol P4
b. Produces 4 mol PH3
c. Cannot react with phosphorus
d. Produces 2 mol PH3
e. Reacts with 2 mol P4
2.
A balanced chemical equation is one that has the same number of moles of molecules on each
side of the equation. True or False? ____________________
3.
A mole ratio is used to convert the moles of starting substance to the moles of the desired
substance. True or False? _____________________
4.
The equation for a reaction should be balanced before doing stoichiometric calculations. True or
False? ____________________
5.
The limiting reactant is the reactant
a. For which you have the lowest mass in grams
b. That has the lowest coefficient in the balanced equation
c. That has the lowest molar mass
d. That is left over after the reaction has gone to completion
e. None of the above.
6.
How many moles of ammonia will be required to produce 13.0 mol of water?
4NH3 (g) + 7O2 (g)  4NO2 (g) + 6H2O (g)
7.
Calculate the moles of oxygen need to react with 21.2g of octane, C 8H18, in the following reaction:
2C8H18 (g) + 25O2 (g)  16CO2 (g) + 18H2O (g)
8.
Nitrogen and hydrogen gases are combined at high temperatures and pressures to produce
ammonia, NH3. If 80.0g of N2 is reacted with excess H2, what number of moles of NH3 will be
formed?
9.
Methane, CH4, the major component of natural gas burns in air to form CO2 and H2O. What
mass of water is formed in the complete combustion of 4.20 x 10 3g of CH4?
10. For the reaction of C2H4 (g) with O2 (g) to form CO2 (g) and H2O (g), what number of moles CO2 can
be produced by the reaction of 0.520 mol of C2H4 and 0.98 mol of O2?
11. Calculate the theoretical yield of iron in the reaction of 258g of iron (III) oxide with 22.5g of
carbon according to the equation:
2Fe2O3 (s) + 3C (s)  4Fe (s) + 3CO2 (g)
12. If the theoretical yield in problem #11 is 71.3g, what is the percentage yield?
Chapter 10 – Energy
1.
Energy due to position is called
a. Potential energy
b. Kinetic energy
c. Bond energy
d. Heat energy
e. None of these.
2.
Energy due to motion is called
a. Potential energy
b. Kinetic energy
c. Bond energy
d. Heat energy
e. None of these.
3.
The law of conservation of energy states
a. Kinetic energy is neither created nor destroyed
b. Potential energy is neither created nor destroyed
c. Energy cannot change forms
d. The energy of the universe never changes
4.
Kinetic energy depends on
a. The mass of the object
b. The velocity of the object
c. The mass and velocity of the object
d. The mass, velocity, and position of the object
e. The velocity and position of the object
5.
Which of the following is a measure of the random motions of the components of a substance?
a. Heat
b. Energy
c. Temperature
d. Work
e. Frictional heating
6.
Thermal energy is transferred from a hot object to a colder object by means of
a. Heat
b. Energy
c. Temperature
d. Work
e. None of these.
7.
What are the metric units for heat?
8.
If a 100.0g sample of water at 60.0oC is added to a 100.0g sample of water at 100.0oC, determine
the final temperature of the water. Assume no heat is lost to the surroundings.
a. 60.0oC
b. 70.0oC
c. 80.0oC
d. 90.0oC
9.
Which of the following is exothermic?
a. The melting of a solid
b. Absorption of energy
c. A reaction in which the products are higher in potential energy than the reactants
d. The condensation of gas
10. In an endothermic reaction, energy flows
a. into the surroundings
b. out of the universe
c. out of the system
d. into the system
11. For an endothermic process, the reactants are of ________________ potential energy (than/as) the
products.
a. Higher
b. Lower
c. The same
d. Depends on the process
12. The amount of energy needed to heat 2.00g of carbon from 50.0 oC to 80.0oC is 42.6 J. The specific
heat capacity of this sample of carbon is:
a. 2256 J/goC
b. 0.710 J/goC
c. 639 J/goC
d. 0.355 J/goC
e. 1.42 J/goC
13. The quantity of heat required to change the temperature of 1g of a substance by 1 oC is defined as
a. A joule
b. Specific heat capacity
c. A calorie
d. Density
14. The specific heat capacity of iron is 0.45J/goC. How many joules of energy are needed to warm
1.50g of iron from 20.00oC to 29.00oC?
15. What is the device used to determine the heat associated with a chemical reaction called?
16. This fossil fuel ________ was formed from the remains of plants that were buried and exposed to
high pressure and heat over time.
a. Coal
b. Natural gas
c. Diesel fuel
d. Propane
e. Gasoline
17. Natural gas consists mostly of
a. Methane
b. Ethane
c. Propane
d. Butane
18. Water is a greenhouse gas. What are other important greenhouse gases?
19. What do greenhouse gases do in the atmosphere?
Chapter 11 – Modern Atomic Theory
1.
Distinguish between the ground state of an electron and the excite state.
2.
The energy levels of the hydrogen atom (and all atoms) are ___________________________,
meaning that only certain discrete energy levels are allowed.
3.
The probability map for an electron is called a(n) ________________________.
4.
As the principal energy level increases in an atom’s orbitals, the average, distance of an electron
from the nucleus ___________________________________.
5.
The shape of the s orbital is ________________________.
6.
Sodium has _________ electron(s) in its outermost energy level.
7.
The halogens (Group 7) contain _____________ valence electrons.
8.
The alkali metals have ______________ valence electrons.
9.
Why do potassium and sodium have similar chemical properties?
10. The noble gases (except He) contain ________ valence electrons.
11. The yet undiscovered element with atomic number 117 would be a member of group ___
12. Which of the following atoms has the largest atomic radius?
Na, Mg, Si, P, or C
13. Which of the following atoms has the highest ionization energy?
Na, Mg, Si, P, or Cl
14. Which of the following atoms has the smallest atomic radius?
As, Sb, Bi, P or N
15. Metal atoms ten to ______________ electrons, while nonmetal atoms tend to ____________
electrons.
16. Of the metals in Group 1, which one has the highest ionization energy? _______________
17. What are the general trends in ionization energy:
a. Across a period ______________________________
b. Down a group (family) _____________________________
18. How does magnesium compare with sodium in terms of the following properties?
a. Atomic size __________________________________
b. Number of valence shell electrons _____________________________
c. Ionization energy __________________________
Chapter 12 – Chemical Bonding
1.
Define:
a. Covalent bond
b.
Polar covalent
c.
Ionic bond
2.
Metals typically have _____________________ electronegativity values (low/high)
3.
Nonmetal elements typically have __________________ electronegativity values (low/high)
4.
In general, a larger atom has a smaller electronegativity. True or False? ________________
5.
An N – F bond is expected to be more polar than an O – F bond. True or False? _________
Use the following choices to classify bonds in each of the following molecules as:
a. Ionic
b. Covalent
6.
N2O ____________________
7.
Al2O3 ________________________
8.
Na2S _____________________
9.
BF3 ______________________
10. H2O2 _______________________
Use the following choices to classify the bonds in each of the following molecules:
a. Nonpolar
b. Polar
11. S8 _________________________
12. CF4 _________________________
13. H2S ___________________________
14. I2 ____________________________
15. CO ____________________________
16. The number of polar covalent bonds in NH3 is _____________________
17. The water molecule is a polar molecule. True or false? ____________________
18. When a molecule has a center of positive charge and a center of negative charge, it is said to be a
polar molecule. True or false?
19. Which of these is not an ionic compound
a. K2CO3
b. CO2
c. NaSCN
d. NH4I
e. MgCl2
20. A phosphorous atom needs to gain _________________ electrons to achieve a noble gas
configuration.
21. When they react chemically, the alkali metals (Group 1)
a. Gain one electron
b. Gain seven electrons
c. Gain or loses seven electrons
d. Lose one electron
22. Draw the Lewis electron structure for the following molecules/ions:
a.
Cl2
b.
HI
c.
CO
d. N2
e.
NI3
f.
CCl4