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Honors Chemistry
End of Course Exam
Study Guide
Chapter 1 Keys to the Study of Chemistry (5.7%)
Learning Objectives: Students should be able to:
1.
Calculate density from mass and volume measurements.
2.
Use conversion factors in calculations.
3.
Use significant figures in calculations.
Chapter 2 Components of Matter (13.2%)
Learning Objectives: Students should be able to:
4.
Describe results of the key experiments by Thomson, Millikan, and Rutherford concerning atomic
structure.
5.
Relate the nuclear symbol to the number of protons and neutrons in the nucleus.
6.
Relate the position of an element in the periodic table to its atomic number and atomic mass.
7.
Use the periodic table to determine the number of electrons available for bonding.
8.
Predict formulas of ionic compounds from charges of ions.
9.
Name and write the formula of ionic and binary covalent compounds.
Chapter 3 Stoichiometry of Formulas and Equations
Learning Objectives: Students should be able to:
10.
Understand that the quantity one mole is set by defining one mole of carbon 12 atoms to have a
mass of exactly 12 grams.
11.
Describe chemical reactions by writing balanced equations.
12.
Determine the molar mass of an ionic or molecular compound from its chemical formula and a
table of atomic masses and convert the mass of a substance to moles.
13.
Calculate the masses of reactants and products in a chemical reaction from the mass of one of the
reactants or products and the relevant atomic masses.
Chapter 4 Three Major Classes of Chemical Reactions
Learning Objectives: Students should be able to:
14.
Write net ionic equations for precipitation and acid-base reactions.
15.
Identify reactions that involve oxidation and reduction and balance oxidation-reduction reactions.
16.
Describe how water dissolves an ionic compound.
Chapter 5 Gases and the Kinetic Molecular Theory
Learning Objectives: Students should be able to:
17.
Understand that the random motion of molecules and their collisions with a surface create the
observable pressure on that surface.
18.
Understand that the kinetic theory of gases relates the absolute temperature of a gas to the
average kinetic energy of its molecules or atoms.
Chapter 6 Thermochemistry” Energy Flow and Chemical Change
Learning Objectives: Students should be able to:
19.
Describe temperature and heat flow in terms of the motion of molecules (or atoms).
20.
Identify processes that can either release (exothermic) or absorb (endothermic) thermal energy.
21.
Solve problems involving heat flow and temperature changes, using known values of specific heat
and latent heat of phase change.
22.
Apply Hess’s law to calculate enthalpy change in a reaction.
Chapter 8 Electron Configuration and Chemical Periodicity
Learning Objectives: Students should be able to:
23.
Relate the position of an element in the periodic table to its quantum electron configuration and to
its reactivity with other elements in the table.
24.
Use the periodic table to identify alkali metals, alkaline earth metals and transition metals, trends
in ionization energy, electronegativity, and the relative sizes of ions and atoms.
25.
Explain how the ionization energy is affected by the shielding effect.
Chapter 9 Models of Chemical Bonding
Learning Objectives: Students should be able to:
26.
Describe why chemical bonds between atoms in molecules such as H2, CH4, NH3, H2CCH2, N2,
Cl2 and many large biological molecules are covalent.
27.
Describe why salt crystals, such as NaCl, are repeating patterns of positive and negative ions held
together by electrostatic attraction.
Chapter 10 The Shapes of Molecules
Learning Objectives: Students should be able to:
28.
Predict the shape of simple molecules and their polarity from Lewis dot structures.
Chapter 12 Intermolecular Forces: Liquids, Solids, and Phase Changes
Learning Objectives: Students should be able to:
29.
Identify solids and liquids held together by ionic bonds, covalent bonds, Van der Waals forces or
hydrogen bonding and relate these forces to physical properties such as electrical conductivity,
volatility and boiling/melting point temperatures.
Chapter 16 Kinetics: Rates and Mechanisms of Chemical Reactions
Learning Objectives: Students should be able to:
30.
Explain that the rate of reaction is the decrease in concentration of reactants or the increase in
concentration of products with time.
31.
Describe role a catalyst plays in increasing the reaction rate.
Chapter 17 Equilibrium: The Extent of Chemical Reactions
Learning Objectives: Students should be able to:
32.
Use LeChatelier’s principle to predict the effect of changes in concentration, temperature, and
pressure.
33.
Write and calculate an equilibrium constant expression for a reaction.
Chapter 18 Acid-Base Equilibria
Learning Objectives: Students should be able to:
34.
Use the pH scale to characterize acid and base solutions.
35.
Calculate pH from the hydrogen-ion concentration.
Chapter 19 Ionic Equilibria in Aqueous Systems
Learning Objectives: Students should be able to:
36.
Describe how buffers stabilize pH in acid-base reactions.
37.
Describe the essential feature of a good buffer system.
38.
Describe how strong and weak acid-base titration curves differ.
39.
Predict the pH range of salt solutions for acid-base neutralization reactions.
Chapter 21 Electrochemistry: Chemical Change and Electrical Work
Learning Objectives: Students should be able to:
40.
Write oxidation and reduction half reactions and identify the oxidizing and reducing agents.
41.
Draw a diagram for a voltaic cell, labeling electrodes and direction of current flow.
42.
Calculate the cell potential using standard potential reduction half-reactions.
Chapter 23 Nuclear Reactions and Their Applications
Learning Objectives: Students should be able to:
43.
Explain that protons and neutrons in the nucleus are held together by nuclear forces that
overcome the electromagnetic repulsion between the protons.
44.
Describe the three most common forms of radioactive decay (alpha, beta, and gamma) and know
how the nucleus changes in each type of decay.
45.
Explain that alpha, beta, and gamma radiation produce different amounts and kinds of damage in
matter and have different penetrations.
46.
Describe the differences between fission and fusion processes.
Practice Problems
Chapter 1 Keys to the Study of Chemistry
1.
To determine the density of an irregularly shaped object, a student immersed the object in 21.2
milliliters of H2O in a graduated cylinder, causing the level of the H2O to rise to 27.8 milliliters. If
the object had a mass of 22.4 grams, what was the object’s density
2.
The measurement 0.5689 gram, rounded to three significant figures is _________.
3.
The measurement 8792 meters rounded to two significant figures is __________.
4.
Perform the following operation and give the answer to the correct number of significant figures
____________.
5.
The density of manganese, a metallic element, is 7.21 g/cm 3. What is the density of manganese
expressed in units of kg/m 3?
Chapter 2 Components of Matter
6.
How does the Rutherford model describe the structure of atoms?
7.
What is the charge on the ion having 12 protons, 11 neutrons, and 11 electrons?
8.
In an atom symbolized as 35X17, there are how many electrons, protons, and neutrons?
9.
The formula for iron (III) phosphate is _________.
10.
The name for PCl5 is ____________.
11.
What ion has atomic number of 31 and only 28 electrons?
12.
Germanium has _______ valence electrons.
13.
What element forms ions having the same electronic configuration as Argon and a charge of -2?
14. What name is given to the compound formed from Cu+2 and NO3- ions?
15. Which ion is not possible: Na+, Ca+2, or Br-2?
Chapter 3 Stoichiometry of Formulas and Equations
Learning Objectives: Students should be able to:
16.
What has the greatest mass: 1 g of magnesium or 1 mole of sodium?
17.
How many molecules are in 237 grams (about a cup) of water?
18. What are the coefficients when the following chemical equation is balanced? AlCl 3 + H2SO4 ---->
Al2(SO4)3 + HCl
19. In the reaction CH4 + 2O2 ---> CO2 + 2H2O, how many grams of oxygen are required to burn 8.0
grams of methane?
20. Aspirin has a formula C9H8O4. The molar mass of aspirin is _______.
21. Write the net ionic equation for the following reaction: KCl(aq) + AgNO3(aq) AgCl(s) + KNO3(aq).
22. Write the net ionic equation for the reaction between sodium carbonate and hydrochloric acid.
Chapter 4 Three Major Classes of Chemical Reactions
Learning Objectives: Students should be able to:
23.
Write the net ionic equation for the reaction that occurs when solutions of lead nitrate and
potassium chloride are mixed.
24.
Is the following reaction an oxidation-reduction reaction: 2Ca(s) + O2(g)  2CaO(s)? Explain
your reasoning.
25.
Write the net ionic equation for the neutralization reaction between HCl(aq) and NaOH(aq).
Chapter 5 Gases and the Kinetic Molecular Theory
Learning Objectives: Students should be able to:
26.
Why does a balloon filled with helium gas decrease in size when exposed to cold CO 2 vapor?
Chapter 6 Thermochemistry Energy Flow and Chemical Change
Learning Objectives: Students should be able to:
27.
When a 120.0-g sample of aluminum (Al) absorbs 9612 J of energy, its temperature increases
from 25.00oC to 115.0oC. Find the specific heat of aluminum.
28.
Review Samples Problem 6.7 (p. 202 in text)
29.
An ice cube melts in a glass of iced tea. Is this an endothermic or exothermic process?
Chapter 8 Electron Configuration and Chemical Periodicity
Learning Objectives: Students should be able to:
30. What element has the electron configuration 1s22s22p5?
31. If an atom’s electron configuration ends in s1 the element is a(n) ___________.
32.
33.
34.
What is the formula of the ion having atomic number 7 and electronic configuration 1s 22s22p6?
What is the symbol of the element having electronic configuration [Ar] 4s 23d2?
Which element exhibits a greater shielding effect, S or Se?
Chapter 9 Models of Chemical Bonding
Learning Objectives: Students should be able to:
35.
When Na2CO3 dissolves in water, the aqueous solution contains ______ and ______ ions.
36.
Calcium and oxygen combine to form an ___________ compound.
37.
Silicon dioxide (SiO2) is a network ______________ solid.
Chapter 10 The Shapes of Molecules
Learning Objectives: Students should be able to:
38.
Boron trifluoride has a ____________ ____________ molecular shape.
39.
Why is the bond angle in SnCl2 only 95o, which is less than the ideal bond angle of 120o for a
molecule with three electron groups?
Chapter 12 Intermolecular Forces: Liquids, Solids, and Phase Changes
Learning Objectives: Students should be able to:
40.
Water (H2O) is a liquid at room temperature, however CO2 and SO2 are gases at room
temperature. The is because ____________ ______________ in water is a stronger
intermolecular force than dipole-dipole or dispersion forces.
41.
If a compound has a high melting point and conducts electricity when dissolved in water, its bonds
are __________.
42.
What type of force exists between oxygen molecules in liquid oxygen?
Chapter 16 Kinetics: Rates and Mechanisms of Chemical Reactions
Learning Objectives: Students should be able to:
43.
Express the rate of the following equation in terms of the change in concentration with time for
each substance: 4NO(g) + O2(g)  2N2O3(g).
44.
A compound that speeds up a reaction, but is not consumed is a ____________.
Chapter 17 Equilibrium: The Extent of Chemical Reactions
Learning Objectives: Students should be able to:
45.
Write the equilibrium constant expression for the reaction SF6(g) + 2 SO3(g)  3 SO2F2(g).
46.
Write the equilibrium constant expression for the reaction CaCO3(s)
CaO(s) + CO2(g)
47.
What effect does the addition of Cl2(g) have on the equilibrium position for the reaction PCl5(g) +
heat  PCl3(g) + Cl2(g)?
48.
What effect does the addition of heat have on the equilibrium position for the reaction PCl 5(g) +
heat  PCl3(g) + Cl2(g)?
Chapter 18 Acid-Base Equilibria
Learning Objectives: Students should be able to:
49.
A sample of rainwater has a pH of 3.5. What ion is sure to be present in relatively large
concentration in this rain sample?
50.
In the reaction C6H5OH + H2O ---> C6H5O- + H3O+ the compound C6H5OH is a(n) _______.
51.
What happens to the pH of an aqueous system if [OH-] is increased.
52.
Is a solution of ammonium chloride (NH4Cl) acidic or basic?
Chapter 19 Ionic Equilibria in Aqueous Systems
Learning Objectives: Students should be able to:
53.
Calculate the pH of a buffer solution that contains 0.100 M HC 2H3O2 and 0.100 M NaC2H3O2. The
acid dissociation constant (Ka) of HC2H3O2 is 1.7 x 10-5.
54.
Which of the following solution systems would make a good buffer, H2CO3/HCO3- or HCl/NaCl?
Chapter 21 Electrochemistry: Chemical Change and Electrical Work
Learning Objectives: Students should be able to:
55.
What are the oxidation states of the hydrogen atom, chlorine atom and oxygen atoms,
respectively, in perchloric acid, HClO4?
56.
Is the following reaction an oxidation-reduction reaction: Na (s) + H2O (l) -> NaOH (aq) + H2 (g)?
57.
58.
59.
60.
Which of the following is an example of an application of electrochemistry: a car battery or a Bic
lighter?
In an electrochemical cell, oxidation occurs at which electrode?
In an electrochemical process called "electrolysis", H2 gas and O2 gas can be obtained by passing
an electric current through liquid water, 2H2O(l) --> 2H2(g) + O2(g). Which species is the
OXIDIZING AGENT and which species is the REDUCING AGENT in this reaction, respectively?
Which of the following statements best describes what will happen when magnesium metal is
added to an aqueous solution containing 1.0 M ferric ion (Fe3+) at 25oC?
Half Reaction
Mg2+(aq) + 2e- --> Mg(s)
Fe3+(aq) + 3e- --> Fe(s)
Cell Potential Eo (V)
-2.375
-0.036
(a) Mg(s) will be oxidized; Fe3+(aq) will be reduced; the standard cell potential will be 2.339 V.
(b) Mg(s) will be oxidized; Fe3+(aq) will be reduced; the standard cell potential will be -2.339 V
Chapter 23 Nuclear Reactions and Their Applications
Learning Objectives: Students should be able to:
61. Carbon-14 radioactively decays via the emission of a beta particle. What is the product nuclide of
this decay?
62. When polonium-216 decays by the emission of one alpha particle and two beta particles, the
product is the isotope __________.
63. Ionizing radiation is many times more dangerous to biological systems than non-ionizing radiation.
The radiation that causes the most damage, if ingested, is __________________.
64. The presence of radon-222 in homes poses a serious health risk. Radon-222 decays by alpha
particle emission with a half-life of 3.8 days. What is the product nuclide from this decay?
65. An isotope of an element decays by emitting an alpha particle and forming lead-208. The original
element was ___________.
66. What type of nuclear process requires a temperature of over a million degrees?
67. The mass of a nucleus is always less than the original mass of its constituent protons and
neutrons. This mass defect is responsible for the nuclear ____________ energy, which holds the
nucleus together.
68. Arrange the following four types of radiation in order from lowest to highest health hazard:
neutrons, alpha particles, gamma rays, and beta particles.
Answers to Problems
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3.4 g/mL
0.569 g
8800 meters
2.6 m2
7.21 x 103 kg/m3
A dense, positively charges nucleus surrounded by electrons, which occupy most of the volume.
+1
17 electrons, 17 protons, 18 neutrons
FePO4
Phosphorus pentachloride
Ga+3
4
Sulfur
Copper(II) nitrate
Br-2
1 mole of sodium
7.92 x 1024 molecules
(2, 3, 1, 6)
32 g O2
180.0 g
Cl- (aq) + Ag+(aq) AgCl(s)
CO32-(aq) + 2H+(aq)  CO2(g) + H2O(l)
Pb2+(aq) + 2 Cl-(aq)  PbCl2(s)
Yes; The oxidation state of Ca changes from 0  2+
H+(aq) + OH-(aq)  H2O(l)
The cold CO2 reduces the helium temperature, which reduces the internal gas pressure in the
balloon.
0.89 J/goC
n/a
endothermic
Fluorine
alkali metal
N+3
Ti
Se
Na+2 and CO3-2 ions.
ionic
covalent
trigonal planar
The stronger repulsion of the lone electron pair decreases the bond angle.
hydrogen bonding
ionic
Dispersion forces
rate = - [O2]/t = - [NO]/4t = [N2O3]/2t
catalyst
Q = [SO2F2]3 / [SF6] [SO3]2
Q = [CO2]
It shifts the equilibrium to the left.
It shifts the equilibrium to the right.
H+ or H3O+
acid
Th pH is decreased.
It is acidic because the ammonium ion reacts with water to make H3O+, which makes the solution
acidic.
pH = 4.770
H2CO3/HCO3- because is consists of a weak acid and its salt (or conjugate base).
+1, +7, -2
Yes, because Na is oxidized.
57.
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63.
64.
65.
66.
67.
68.
A car battery
Anode
H2O(l) ; H2O(l)
Statement (a)
nitrogen-14
polonium-212
alpha particles
polonium-218
polonium
fusion reaction
binding
gamma rays < alpha particles < beta particles < neutrons