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CHAPTER 1
-Chemistry
-Matter
-Elements
-Atoms
-Molecules
-States of Matter (Gas/Liquid/Solid)
-Pure Substance
-Elements
-Compounds
-Mixtures
-Law of Constant Composition (Law of Definite
Proportions)
-Solutions
-Physical Properties
-Chemical Properties
-Intensive Properties
-Extensive Properties
-Physical Change
-Chemical Change
-Changes of State
-Scientific Method
-Hypothesis
-Scientific Law
-Theory
-Metric System
-SI Units
-Mass
-Celsius Scale
-Kelvin Scale
-Density
-Accuracy
-Precision
-Significant Figures
-Conversion Factor
-Dimensional Analysis
3) A weighing tray has a mass of 9.11 g. 3.2 g of NaOH
is added to this tray. The total mass of the tray and the
NaOH should be recorded as:
(a) 4.1 g
(b) 4.11 g
(c) 4.111 g
(d) 4.100 g
(e) 4.10 g
1) Which of the following represents a chemical
change?
(a) water boiling
(b) iodine subliming
(c) sugar dissolving in water
(d) natural gas burning
(e) ice melting
2) A solid white substance A is heated strongly in the
absence of air. It decomposes to form a new white
substance B and a gas C. The gas has exactly the same
properties as the product obtained when carbon is
burned in an excess of oxygen. Based on these
observations, can you determine whether A and B and
the gas C are elements or compounds? Explain your
conclusions for each substance.
2) How many carbon atoms are contained in 4.2 g of
C3H6?
(a) 1.8 x 10^23
(b) 9.0 x 10^23
(c) 6.0 x 10^23
(d) 1.8 x 10^24
(e) 6.0 x 10^24
4) A combination of sand, salt, and water is an
example of a________.
(a) homogeneous mixture
(b) heterogeneous mixture
(c) compound
(d) pure substance
(c) solid
5) The law of constant composition says_________.
(a) that the composition of a compound is always the
same
(b) that all substances have the same composition
(c) that the composition of an element is always the
same
(d) that the composition of a homogeneous mixture is
always the same
(e) that the composition of a heterogeneous mixture
is always the same
1) Show all work and use the correct set up, units, and
number of sig figs.
Convert 32 mph to km/sec (1.6 km = 1 mile)
1. Which of the following is an example of a chemical
change?
a) Mixing Iron and Sulfur Powder
b) Breaking Glass
c) Formation of Clouds
d) Rotting of Fruit
2. If the high temperature for the day is predicted to
reach 88 ˚F, what is the predicted temperature in ˚C?
a) 304 ˚C
b) 31 ˚C
c) 159 ˚C
d) 67 ˚C
3. 1 Megameter is equal to how many meters?
a) 10
b) 1000
c) 1000000
d) 1000000000
4. Which is an example of a compound?
a) Bronze
b) Silver
c) Gold
d) Platinum
5. Which is an example of an element?
a) Earth
b) Mars
c) Mercury
d) Venus
Free Response Questions
1. A package of aluminum foil contains 50ft2 of foil,
which weighs approximately 8.0 oz. Aluminum has a
density of 2.70 g/cm3. What is the approximate
thickness of the foil in millimeters?
2. (a) What is the difference between a hypothesis and
a theory? (b) Explain the difference between a theory
and a scientific law. Which addresses how matter
behaves, and which addresses why it behaves that
way?
Chapter 2:
Atoms
Subatomic Particles
Dalton’s Atomic Theory of Matter
Law of Conservation of Mass/Matter
Law of Constant Composition
Law of Multiple Proportions
Cathode Rays
Radioactivity
Nucleus
Protons, Neutrons, Electrons
Electronic Charge
Atomic Mass Unit
Angstrom
Atomic Number
Mass Number
Atomic Symbol
Isotopes
Atomic Weight
Periodic Table
Periods
Groups
Metals/Metallic Elements
Nonmetals/Nonmetallic Elements
Metalloids
Molecule
Diatomic Molecule
Chemical Formula
Molecular Compounds
Molecular Formulas
Empirical Formulas
Structural Formula
Ion
Cation
Anion
Polyatomic Ions
Ionic Compound
Chemical Nomenclature
Oxyanions
Acids
Binary Molecular Compounds
Organic Chemistry
Hydrocarbons
Alcohol
Alkanes
1) The symbols 11H, 21H, 31H represent three different
(a) Homologs
(b) Isotopes
(c) Isomers
(d) Allotropes
(e) Conformations
2)An isotope has the atomic number of 8 and a mass
number of 17. This element
(a) Is an isotopic form of oxygen
(b) Has 8 neutrons
(c) Has 9 protons if it is in ionic form
(d) Is an isotopic form of fluorine
(e) Has 17 electrons
3) Which of the following ions has the same number of
electrons as Br(a) Ca+2
(b) K+
(c) Sr+2
(d) I(e) Cl4) For which of the following pairs are the atoms most
likely to form an ionic compound?
(a) Carbon and Oxygen
(b) Calcium and Chlorine
(c) Chlorine and Oxygen
(d) Sodium and Magnesium
(e) Chlorine and Neon
5) The discovery of the electron came from:
(a) The gold foil experiment
(b) Dalton’s Atomic Theory
(c) The plum pudding model
(d) Cathode ray experiments
(e) The calculation of the mass/charge relationship
1) The element boron consists of two isotopes, 10B and
11B. Their masses, based on the carbon scale, are 10.01
and 11.01, respectively. The abundance of 105B is
20.0% and the abundance of 115B is 80.0%. What is the
atomic mass of boron?
2) What ionic compound is formed with the ions of
Strontium (Sr) and Phosphorus (P) combined? Label
which ion is the cation and which ion is the anion.
1.
Nitrogen Gas is bubble through an aqueous
solution of calcium Bromate. Besides water,
this statement refers to which chemical
formulas?
A) N and CaBrO3
B) N2 and CaBrO4
C) N2 and CaBrO3
D) N and CaBrO3
E) N and Ca (BrO)3
2.
Magnesium Nitride reacts with water to form
ammonia and magnesium hydroxide. Besides
water, this statement refers to which chemical
formulas?
A) Mg3N2 , NH3 , and Mg(OH)2
B) Mg(NO3)2 , NH4+ , and MgH2
C) Mg(NO3)2 , NH3 , and MgH2
D) Mg(NO2)2 , NH4+ , and Mg(OH)2
E) Mg3N2 , NH3 , and MgOH
3.
Which of the following compounds contains
four carbon atoms?
A) Propane
B) Butanoic acid
C) Ethyl methyl ether
D) 2-pentanol
E) Heptanal
4.
Which is true of the 243Am3+ ion?
Protons Electrons Neutrons
A) 95
92
243
B) 95
98
243
C) 95
95
148
D) 95
92
148
E) 92
95
148
5.
The correct name for the compound Cu3N is:
A) Copper (III) Nitride
B) Copper Nitrogen
C) Copper (II) Nitrate
D) Copper (I) Nitride
E) Copper (III) Nitrate
1) Using the periodic table to guide you, predict
the chemical formula and name of the
compound formed by the following elements.
(A) Ga and F
(B) Li and H
(C) AL and I
(D) K and S
2) Predict wheather each of the following
compounds is ionic or molecular.
A) B2H6
B) CH3OH
C) LiNO3
D) Sc2O3
E) CsBr
F) NOCL
G) NF3
H) Ag2SO4
Chapter 3
StoichiometryMole Ratio
ReactantsProductsCombination (Synthesis)ReactionSingle Replacement Reaction
Double Replacement reaction
Decomposition ReactionCombustion ReactionAvagadro's NumberMoleLimiting ReactantTheoretical YieldPrecent YieldEmpirical FormulaMolecular Formula2. The percentage of oxygen in C8H12O2 is:
a.(16/140)(100)
b. (32/140)(100)
c. (16/124)(100)
d. (140/32)(100)
e. (32/124)(100)
3. A compound contains 48% O, 40.0 % Ca, and the
remainder is C. What is the empirical formula?
A) O3C2Ca2
B) O3CCa2
C) O3CCa
D) O3CCa2
E) O2CCa
4. Given the unbalanced equation: Al + O2 = Al2O3
When this equation is completely balanced using the
smallest whole numbers, what is the sum of the
coefficients?
A) 9
B) 7
C) 5
D) 4
5. What is the empirical formula of the compound
whose molecular formula is P4O10?
A) PO
B) PO2
C) P2O5
D) P8O20
Balance the following equation
___C0 +___ H2 > ____ C8H18 + ______H2O
Combustion of 8.652 grams of a compound containing
C, H, O, and N yields 11.088g CO2, 3.780 grams of H20
and 3.864 grams of NO2.
a. How many moles of C, H, and N are contained in the
sample?
b. How many grams of Oxygen are contained in the
sample?
c. What is the simplest formula of the compound?
d. If the molar mass of the compound lies between 200
and 300, what is its molecular formula?
e. Write and balance a chemical equation for the
combustion of the compound.
How many liters of hydrogen are needed to react with
10 liters of nitrogen gas in the reaction forming
ammonia?
2.)What is the total number of atoms in 0.260 mol of
glucose?
1.)How many moles of Al are there in 2.16 mol of
Al2O3?
a. 2.16
b. 1.08
c. 4.32
d. 10.8
2.)What is the empirical formula of a compound whose
molecular formula is P4O10?
a. PO
b. PO2
3. P2O5
4. P8O20
3. What is the Percent by mass of oxygen in
magnesium oxide, MgO?
a. 20%
b. 40%
c. 50%
d. 60%
4.)In the chemical equation _CH4+_Cl2--->_CCl4+_HCl,
what are the correct coefficients?
a. 0,4,0,4
b.4,8,4,2
c.2,5,2,4
d.1,5,2,3
5.)What is the mass in grams of 3.0 x 1023 molecules
of CO2?
a. 22 g
b. 44 g
c. 66 g
d. 88 g
Chapter 4:
Aqueous Solutions
• Solvent
• Solutes
• Dissociate
• Dissolve
• Electrolyte
• Non-electrolyte
• Strong Electrolyte
• Weak Electrolyte
• Solvation
• Chemical Equilibrium
• Precipitation Reactions
• Precipitate
• Solubility
• Exchange/Metathesis
Reactions
• Molecular Equation
• Complete Ionic Equation
• Spectator Ions
• Net Ionic Equation
• Acids
• Bases
• Strong Acids & Bases
• Weak Acids & Bases
• Neutralization Reaction
• Oxidation Reduction (redox)
Reaction (OIL RIG)
• Oxidation
• Reduction
• Oxidation Number
• Displacement Reactions
• Activity Series
• Concentration
• Molarity (M)
• Dilution
• Titration
• Standard Solution
• Equivalence Point
• Acid-Base Indicators
1. Consider the equation: Cl2(g)+2KI(aq )-> I2(s)+KCl(aq), which species was oxidized?
a. Cl2
b. K+
c. Id. I2
e. Cl2. How many grams of NaOH are needed to
prepare 500 mL of a 1.0M solution?
a. 40 grams
b. 20 grams
c. 200 grams
d. 400 grams
e. 12.5 grams
3. How many millimeters of 0.2M FeCl3 solution
would be necessary to precipitate all of the Ag+
from 60mL of a 0.6M AgNO3 solution? (Hint: write
a balanced equation)
a. 10 mL
b. 30 mL
c. 60 mL
d. 90 mL
e. 180 mL
4. The definition of a precipitate reaction is:
a. A reaction that results in a soluble product
b. A reaction that results in an insoluble product
c. A reaction where water is produced
d. A reaction when an acid and a base are mixed
e. A reaction when electrons are transferred as
reactants
5. What mass of KCl is needed to precipitate the
silver ions from 15.0 mL of 0.200M AgNO3 solution?
a. 0.16 g KCl
b. 29 g KCl
c. 18 g KCl
d. 0.4456 g KCl
e. 0.227 g KCl
1.
Label the following reactions with as many labels
as possible (Synthesis, Decomposition, Single
Replacement, Double Replacement, Redox,
neutralization, precipitation)
a. Ca(OH)2(aq)+2HNO3 --> Ca(NO3)2(aq) + 2H20
b. Fe2O3(s)+3CO(g) --> 2Fe(s)+3CO2(g)
c. CuBr2(aq)+2NaOH(aq) -->
Cu(OH)2(s)+NaBr(aq)
d. Cl2(g)+NaI(aq) --> I2(s)+2NaCl(aq)
2. HCl, HBr, and HI are strong acids, yet HF is a
weak acid. What does this mean in terms of the
event to which these substances are ionized in
solution?
1. What is the concentration of a solution of HCL if 400
mL of
distilled water is added to 10 mL of stock solution?
a. 400. M
b. 410 M
c. 1.25 M
d. 0.250 M
e. 0.244 M
2. The difference between a strong acid and a weak
acid is:
a. A strong acid is more concentrated than a weak acid
b. A weak acid is more soluble in water than a strong
acid
c. Strong acids have more hydrogens per molecule
than weak
acids
d. Weak acids cannot conduct electrical currents in
solution
e. Strong acids are more completely dissociated
compared to
weak acids
3. Which of the following species is reduced in the
following netionic
reaction:
2HNO3(aq) + 3H3AsO3(aq) 2NO(g) + 3H3AsO4(aq) +
H20(l)
a. HNO3
b. H3AsO3
c. NO
d. H3AsO4
e. H2O
4. Which substance will not form a gas upon mixing
with an
aqueous acid?
a. NaHCO3(s)
b. CaS(s)
c. Ca(s)
d. K2SO3(s)
e. Al2O3(s)
5. Which of the following is not a strong electrolyte:
a. CH3CO2H
b. KCl
c. KNO3
d. NaClO4
6. One commercial method used to peel potatoes is to
soak them in a solution of NaOH for a short time,
remove them from the NaOH and spray off the peal.
The concentration of the NaOH is normally in the
range of 3 to 6 M. The NaOH is analyzed periodically.
In one such analysis, 45.7 mL of 0.500 M H2SO4 is
required to neutralize a 20.0 mL sample of the NaOH
solution. What is the concentration of the NaOH
solution?
7. Write the net ionic equation for the precipitation
reaction that occurs when solutions of calcium
chloride and sodium carbonate are mixed.
The net ionic equation for the reaction between iron
and copper (II) chloride is:
A) Fe+CuCl2àCu+FeCl2
B) Fe+Cu2++2Cl-àCu+Fe2++2ClC) Fe+Cu2+àCu+Fe2+
D) Fe2++CuCl2àCu2++FeCl2
E) Cu2++2Cl-àCuCl2
If 53 grams of Na2Cl3 (Molar Mass 106g) are added to a
500 mL solution of 0.50 M NaCl solution, what will the
resulting concentration of the Na+ ion be?
A) 0.10 M
B) 0.50 M
C) 0.75 M
D) 1.5 M
E) 2.5 M
50.0 mL of 1.0 M CaCl2 were required to precipitate
out all of the silver in a 100 mL solution of AgNO3 to
create AgCl.
A) 0.10 M
B) 0.50 M
C) 1.0 M
D) 2.5 M
E) 5.0 M
Which of the following species is reduced in the
following net-ionic reaction?
2HNO3(aq)+3H3AsO3(aq)à2NO(g)+3H3AsO4(aq)+H2O(l)
A) HNO3
B) H3AsO3
C) NO
D) H3AsO4
E) H2O
When copper chloride reacts with sodium sulfide the
spectator ions are:
A) Copper and Sulfide
B) Sodium and Chloride
C) Chloride and Sulfide
D) No ions are spectators
E) All 4 ions are spectators
1. 2M potassium hydroxide solution is titrated with a
0.1M nitric acid solution.
Balanced Equation:
What would be observed if the solution was titrated
well past the equivalence point using bromthymol
blue as the indicator? (Bromthymol blue is yellow in
acidic solution and blue in basic solution.)
2. Propane is burned completely in excess oxygen gas.
Balanced Equation:
When the products of there action are bubbled
through distilled water, is the resulting solution
neutral, acidic, or basic? Explain.
Chapter 5
Thermodynamics:
Thermochemistry:
Energy:
Work:
Force:
Heat:
Kinetic Energy:
Potential Energy:
Joule:
Calorie:
System:
Surroundings:
First Law of Thermodynamics:
Inertial Energy:
Endothermic:
Exothermic:
State Function:
Enthalpy:
Enthalpy of Reaction:
Calorimeter:
Calorimetry:
Heat Capacity:
Specific Heat:
Molar Heat Capacity:
Bomb Calorimeter:
Hess’s Law
Enthalpy of Formation:
Standard Enthalpy Change:
Standard Enthalpy of Formation:
3) Which of the following is endothermic?
a) rusting iron
b) making ice cubes
c) cooking an egg
d) a candle flame
1) Which change will result in an increase in enthalpy
of the system?
a) Burning a candle
b) Freezing water
c) Evaporating alcohol
d) Dropping a ball
e) Condensing steam
3.
2) Given the following data, what is the heat of
formation of methane gas?
I. CH4+2O2CO2+2H2O
H=-803kJ
II. H2+1/2O2H2O
H=-242kJ
III. C+O2CO2
H=-394kJ
IV. C+1/2O2CO
H=-111kJ
a)-803kJ/mol
b)-75kJ/mol
c)+167kJ/mol
d)+208kJ/mol
e)-1439kJ/mol
4) A 2.200g sample of quinone (C6H4O2) is burned in a
bomb calorimeter whose total heat capacity is
7.854kJo/C. The temperature of the calorimeter
increases from 23.44oC to 30.57oC. What is the heat of
combustion per gram of quintone.
a)-25.5 kJ/mol
b)-330 kJ/mol
c)44 kJ/mol
d)990 kJ/mol
5) Calculate the Kinetic energy in Joules of a 45g golf
ball moving at 61 m/s.
a)6966J
b)95J
c)84J
d)-990J
Free Response
1) What are the three laws of thermodynamics?
1.
2.
2) How much heat is released when 4.50g of methane
gas is burned in a constant pressure system?
1) Energy is Measured in
a) Degrees
b) Joules
c) grams
d) liters
3) Calculate the work done on the system when 1.00
mol of gas held behind a piston expands irreversibly
from a volume of 1.00 dm3 to a volume of 10.0 dm3
against an external pressure of 1.00 bar.
a) -5710 J
B) -900 J
c) +1000 J
d) +9120 J
4) 1.00 mol of gas in a cylinder is compressed
reversibly by increasing the pressure from 1.00 bar to
10.0 bar at a constant temperature of 500 K. Calculate
the work done on the gas by the compression.
A) +9570 J
b) +3740 J
c) +37.4 kJ
d) +5710 J
5) The temperature of a copper block of mass 423g
rises by 10.1°C. Calculate the heat transferred, given
that the specific heat capacity of copper is 385 J K-1
kg-1.
a) 385 J
B) 1.63 kJ
c) 1.10 kJ
d) 25.6 kJ
Free Response
1) Imagine a book falling from a shelf. At a particular
moment during its fall, the book has a kinetic energy of
13 J and a potential energy with respect to the floor of
72 J. How does the books kinetic and potential energy
change as it continues to fall?
2) What is its total kinetic energy at the instant just
before it strikes the floor?
1) What is the change in internal energy for the
following? A system absorbs 45 kJ of heat and does 5
kJ of work to the surroundings.
a. 45 kJ
b. -45 kJ
c. 40 kJ
d. 50 kJ
e. -50 kJ
2) The value of the change in heat for the reaction is 70 kJ.
kJ of heat are released when 14 grams of
hydrogen gas are completely reacted.
a. 5 kJ
b. 10 kJ
c. 140 kJ
d. 490 kJ
e. 980 kJ
3) The temperature of a 10g sample of metal increase
from 20°C to 30°C when 30 J of heat are added to it.
The specific heat capacity of the metal is:
a. 3.3 J/°C
b. 0.3 J/°C
c. 30 J/°C
d. 3.3 J/(g°C)
e. 0.3 J/(g°C)
4) What is the molar heat of fusion of a substance that
required 45 kJ to melt 38 grams of the substance? (the
molar mass of the substance is 152 grams/mol)
a. 11 kJ
b. 180 kJ
c. 90 kJ
d. 3 kJ
e. 6 kJ
5) Which of the following statements are true for the
statment: "When a substance was dissolved in water,
the temperature of the solution increased."
I.
The process was endothermic
II.
The process is exothermic
III.
The enthalpy of the process can be
calculated by multiplying the temperature
change and heat capacity of the system.
IV.
The heat that caused the temperature
change came from the surroundings.
a. I
b. II
c. I and III
d. II and III
e. I,III, and IV
1) If a system absorbs 53 kJ of heat and does 3 kJ
of work, what is the change in internal
energy?
2) How much heat is released when 4.50 g of
methane gas is burned in a constant pressure
system?
Chapter 6
Electrons of an atom represent the density of that
atom and their place on the periodic table
The Pauli Exclusion Principle:
Orbitals (S,P,D.F)
Electron configuration:
Paramagnetic:
Diamagnetic:
Ground state electron configuration:
An excited state configuration:
Electron shell:
Subshell:
Degenerate orbitals:
Hund’s Rule (Bus seat rule):
Valence electrons:
Core electrons:
electronic structure
Wavelength λ
frequency ν
Amplitutde
Electromagnetic radiation:
Visible spectrum
Line spectrum:
Photon:
A quantum
The speed of light =
Planck’s constant =
The uncertainty principle:
Quantum numbers:
Quantum Number
Principle
Azimuthal
Magnetic
Spin
Symbol
n
l
ml
ms
Possible values
1, 2, 3, 4, 5…
0, 1, 2, 3…n
s, p, d, f
-l, …0… +l
+ 1/2
Electron/ Orbital
characteristic
size
shape
orientation
Magnetic spin
1.what is the maximum number of electrons that can
fill the 5f subshell?
A.2 B.5 C.10. D.14. E.18
2.a blue line in the atomic emission spectrum of
hydrogen has a wavelength of 434nm. What is the
energy of this light per mole of photon?
A. (10^6)(6.63)(3.00)(6.02)/(434) kJ/mol
B. (10^3)(6.63)(3.00)(6.02)/(434) kJ/mol
C. (10^6)(6.63)(3.00)(6.02)/(434) J/mol
D. (10^3)(6.63)(3.00)(6.02)/(434) J/mol
E. (10^3)(434)(6.02)/(6.63)(3.00) kJ/mol
3.which set of quantum numbers is not allowed?
A.(2,2,1,1/2). B.(3,2,0,-1/2). C.(4,3,-3,1/2).
D.(5,3,-2,1/2). E.(6,2,-1,1/2)
4.According to Heisenberg's uncertainty prunciple, if
you know the position of the electron you won't know
the...
A.speed. B. shape of the orbital it's in. C. Energy.
D. Momentum. E. principle quantum number.
5. Wavelength and frequency are
A. Directly related. B. proportional. C. Inversely
related. D second cousins. E.the same
1. An electromagnetic wave has a wavelength of
656nm. How much energy is there in a mole of
photons?
2.. Write the Noble gas configuration of cobalt.
1.) What is the wavelength of light that has a
frequency of 6.0×10∆14 Hz?
a. 2000nm
b. 500nm
c. 200nm
d. 2.0×106 nm
e. 5.0×10-7 nm
2.) The outer most electron in a ground state
potassium atom can be described by which of
the following sets of four quantum numbers?
a. 4, 0, 0, ½
b. 4, 1, 0, ½
c. 4, 1, 1, ½
d. 5, 0, 0, ½
e. 5, 1, 0, ½
3.) One type of sunburn occurs on exposure to UV
light of wavelength in the vicinity of 325nm.
What is the energy of a photon of this
wavelength?
a. 6.11×10∆-19 J/photon
b. 3.2×10∆-8 J/photon
c. 680 KJ/photon
d. 12×6∆-19 KJ/photon
4.) Arrange the following types of
electromagnetic radiation in order of
increasing wavelength:
a. Radio waves < ultra violet< green
light< red light<infrared< X-rays
b. Green light< red light< infrared< Xrays< radio waves< ultra violet
c. X-rays< ultra violet< green light< red
light< infrared< radio waves
d. Red light< radio waves< infrared<
green light< ultra violet< X-rays
5.) Gaseous atoms of which of the following
elements are paramagnetic in their ground
states?
I.
Na
II.
Mg
III.
Al
IV.
P
a. I, II, III, IV
b. I, II, III only
c. I, III, IV only
d. II only
e. III, IV only
Which orbital is the biggest?
a)1s
b)3p
c)3s
d)4d
e)2f
The outermost electron in a ground state K atom can
be described by
which of the following sets of four quantum numbers?
a)4,0,0,1/2
b)4,1,0,1/2
c)4,1,1,1/2
d)5,0,0,1/2
e)5,1,0,1/2
What visible color has the longest wavelength?
a)purple
d)green
b)yellow
e)blue
c)red
Which wave would cause the most damage?
a)microwaves
b)infrared
c)ultra-violet
d)X rays
e)gamma rays
How many orientations does a d orbital have?
a)1
b)2
c)3
d)4
e)5
1) In a condensed electron configuration, what does
the symbol [Xe] represent?
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atomic radius
nonmetals
alkali metals
halogens
valence orbitals
effective nuclear charge
bonding atomic radius
isoelectronic series
ionization energy
electron affinity
metallic character
1) In general, as you go across a period in the periodic
table from left to right the atomic radius:
a) Decreases
b) Increases
c) Stays the same
2) In general, as you go across a period in the periodic
table from left to right the electron affinity becomes:
a) Increasingly negative
b) Decreasingly negative
c) Stays the same
3) The ion with the largest diameter is:
a) Brb) Clc) Fd) I-
2) What is the ground state configuration of V?
3) If human height were quantized in one-foot
increments, what would happen to the height of the
child as she grows?
4) Explain how the existence of line spectra is
consistent with Bohr’s theory of quantized energies
for the electron in the hydrogen atom.
Chapter 7
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noble gases
force of attraction
core electrons
shielding effect
4) In general, as you go across a period in the periodic
table from left to right the first ionization energy:
a) Stays the same
b) Increases
c) Decreases
5) How many unpaired electrons does Ni2+ contain
when the electron configuration is written out?
a) 1
b) 2
c) 4
d) None
1) What is the general relationship between the
size of an atom and its first ionization energy?
Which element in the periodic table has the
largest ionization energy? Which has the
smallest?
E) Ne, Na, Mg
1)
2) Which will experience the greater effective
nuclear charge, the electrons in the n = 3 shell
in Ar or in the n = 3 shell in Kr? Which will be
closer to the nucleus? Explain
1) Which list of elements is arranged in order of
increasing atomic size (largest last)?
A) Be, Mg, Ca, Sr, Ba
B) Ba, Sr, Ca, Mg, Bc
C) Be, Ca, Ba, Mg, Sr
D) Be, Ba, Ca, Mg, Sr
E) Ba, Sr, Ca, Be, Mg
2) Which sequence is arranged in order of increasing
ionization energies, lowest to highest?
A) Be, B, C, N, O
B) B, Be, C, O, N
C) Be, B, C, O, N
D) B, Be, C, N, O
E) O, N, C, B, Be
3) Electrons in the 1s subshell are much closer to the
nucleus in Ar than in He due to the larger ______ in Ar.
A) nuclear charge
B) paramagnetism
C) diamagnetism
D) Hund’s rule
E) azimuthal quantum number
4) ______ is isoelectronic with argon and _____ is
isoelectronic with neon.
A) Cl-, Cl+
B) F+, FC) Ne-, Kr+
D) Ne-, Ar+
E) Cl-,F5) In which set of elements would all members be
expected to have very similar chemical properties?
A) N, O, F
B) Na, Mg, K
C) O, S, Se
D) S, Se, Si
A) What is meant by the term effective
nuclear charge?
B) How does the effective nuclear charge
experienced by the valence electrons of an
atom vary going from left to right across a
period of the periodic table?
2)
A) What is an isoelectronic series?
B) Which neutral atom is isoelectronic with
each of the following ions: Al3+, Ti4+, Br-, Sn2+.
Chapter 8
Chemical bonds
ionic bonds
covalent bonds
metallic bonds
Lewis symbol (Lewis dot structure)
octet rule
lattice energy
electron configuration
single bond
double bond
triple bond
bond length
bond polarity
nonpolar covalent bond
polar covalent bond
electronegativity
polar molecule
dipole
dipole moment
formal charge
resonance structures
valence electrons
bond enthalpy
Multiple Choice
1. As the number of bonds increases, the bond
length
a. decreases
b. increases
c. is not affected (stays the same)
d. there is not enough information to
know
2.
Of the molecules, which bond is most polar?
a. HI
b. HF
c. H2
d. HBr
e. none are polar
3.
Which of the following elements will tolerate
less than a full octet of electrons?
a. I
b. Na
c. B
d. O
e. none of the above
4.
Which pair of ions has the highest lattice
energy?
a. Na+ and Brb. Li+ and Fc. Cs+ and Fd. Li+ and O2e. K+ and F-
5.
Resonance structures differ by:
a. number and placement of electrons
b. number and placement of atoms
c. number of electrons only
d. placement of atoms only
e. number of atoms only
f. placement of electrons only
Free Response
1. Draw the Lewis structures for IF6+ and Sulfur
Trioxide
2.
Find the enthalpy change for the following
reaction:
C3H8 (g) + 5 O2 (g)  3 CO2 (g) + 4
H2O (g)
H-C  413
H-O  463
C-C 348
C=C  614
O-O  146
C-O 358
C=O 799
O=O  495
Chapter 9
Bond angles
VSEPR
Bonding pair
Electron domain
Nonbonding pair (or lone pair)
Electron-domain geometry
Linear
Trigonal planar
Tetrahedral
Trigonal Bipyramidal
Octahedral
Molecular geometry
Linear
Trigonal planar
Bent
Tetrahedral
Trigonal pyramidal
Trigonal Bipyramidal
See saw
T-Shaped
Octahedral
Square pyramidal
Square planar
Bond dipole
Polar vs. nonpolar molecules
Valence-bond theory
Hybrid orbitals
Hybridization
Sigma bonds
Pi bond
Delocalized
The electron domain and molecular geometry of
BrO2- is __________.
A) tetrahedral, trigonal planar
B) trigonal planar, trigonal planar
C) trigonal pyramidal, linear
D) tetrahedral, bent
E) trigonal pyramidal, seesaw
The hybridizations of nitrogen in NF3 and NH3 are
_________ and _________, respectively.
A) sp2, sp2
B) sp, sp3
C) sp3, sp
D) sp3,sp3
E) sp2, sp3
PCl5 has _________ electron domains and a ________
molecular arrangement.
A) 6, trigonal bipyramidal
B) 6, tetrahedral
C) 5, square pyramidal
D) 5, trigonal bipyramidal
E) 6, seesaw
Which set does not contain a linear species?
A) CO2, SO2, NO2
B) H2O, HCN, BeI2
C) OCN-, C2H2, OF2
D) I3-, BrF3, SCNE) H2S, ClO2-, NH2What is the best estimate of the H-O-H bond angle
in H3O+?
A) 109.5º
B) 107º
C) 104.5º
D) 116º
E) 120º
Free Response Questions
An AB3 molecule is described as having a trigonal
bipyramidal electron-domain geometry. How
many nonbonding domains are on atom A?
Explain.
Why are there no sp4 or sp5 hybrid orbitals?
Chapter 10
atmosphere
torr
mm Hg
Boyle's Law
Charles's Law
Unnamed Relationship
Combined Gas Law
Avogadro's Hypothesis
Avogadro's Law
Ideal Gas Equation
Ideal Gas
Standard Temp. and Pressure
Dalton's Law of Partial Pressures
Mole Fraction
Kinetic-molecular Theory
Root mean square speed
Effusion
Diffusion
Graham's Law
mean fee path
van der Waal's equation
Deviation from ideal behavior
STP
Collected over water (Partial Pressures)
1.Suppose you are given 2 1-L flasks and told that one
contains a gas of molar mass 30, the other a gas of
molar mass 60, both at the same temperature. The
pressure in flask A is X atm, and the mass of the gas is
1.2g. The pressure in flask B is .5X atm and the mass is
1.2g. Which flask contains the gas of molar mass 60?
1. Under which of the following conditions will a
1-L sample of Kr gas act the most ideal?
a. .1 atm and 273 K
b. 1.0 atm and 273 K
c. 10 atm and 400 K
d. .1 atm and 400K
e. 1.0 atm and 400 K
2. Which of the following gases is most likely to act
ideally under STP?
a. Kr
b. CCl4
c. He
d. O2
e. N2
3. Which increases when a gas is heated at
constant volume?
I. Pressure
II. Kinetic energy of molecules
III. Attractive forces between molecules
a. I only
b. II only
c. III only
d. I and II only
e. I and III only
4. Effusion is:
a. The spread of a substance from a high
concentration to low concentration
b. The spread of a substance through a tiny hole
c. The macroscopic behavior of gases at the atomic
and molecular level
d. The ease at which an atom gains an electron
5. What is the conversion factor between atm and
torr? mmHg and torr?
a. 1 atm = 765 torr, 1mmHg =1 torr
b. 1 mmHg =700 torr
c. 2 atm = 760 torr, 1 mmHg = 760 torr
d. 1 atm = 760 torr, 760 mmHg = 760 torr
What properties of gases can you point to that support
the assumption that most of the volume in a gas is
empty space?