Download Honors Chemistry – Mr

Honors Chemistry – Mr. Scott
Semester 1 Final Exam Study Guide (Chapters 1 – 10)
Name: _______________
Date: __________
Period: ____
Atomic and Molecular Structure (Chapters 2, 7 & 8)
1. The periodic table displays the elements in increasing atomic number and shows how periodicity of
the physical and chemical properties of the elements relates to atomic structure.
a. Learning objective: know how to relate the position of an element in the periodic table to its
atomic number and atomic mass.
The nucleus of an atom contains protons and neutrons. The number of protons (atomic number)
determines the unique properties of an element. Elements are arranged in the Periodic Table in order of
increasing atomic number. Each element has a unique atomic mass, which is the combined mass of its’
protons and neutrons.
Q1 – How many protons, electrons, and neutrons are in a neutral atom of beryllium?
a) 4,4,2
b) 4,4,4
c) 4,2,3
d) 4,4,5
Q2 – The mass number of an atom is the number of ________________ in the atom.
a) protons
b) neutrons
c) protons plus the number of neutrons
c) protons plus the number of electrons
Q3 – The atomic number and mass number of cobalt is _____ and _____, respectively?
a) 27, 58.93
b) 27, 45
c) 26, 55.85
d) 27, 58
Q4 – Using the periodic table, determine the number of neutrons in the
a) 27
b) 30
c) 33
60Co
isotope.
d) 59
Q5 – The number of protons in a silicon atom is _______.
a) 28
b) 14
c) 32
d) 16
b. Learning objective: know how to use the periodic table to identify metals, semimetals, nonmetals, and halogens.
Most periodic tables include a dark stepped line running from Boron to Astatine. Elements to the
immediate right and left of this line, excluding Al, Po, and At are semimetals (metalloids) that have
properties intermediate between metals and nonmetals. Elements further to the left are metals.
Elements further to the right are nonmetals.
Q1 – Which group of the following elements contains all metals?
a) Be, Cr, Se, Ag
b) K, Ca, Fe, Si
c) Mg, Ti, Fe, Zn
d) Li, C, Fe, Ni
Q2 – Silicon and germanium are ______________________?
a) metals
b) non-metals
1
c) metalloids (semimetals)
d) transition metals
Q3 – Fluorine and chlorine are __________________?
a) metals
b) halogens
c) noble gases
d) metalloids (semimetals)
Q4 – Which group of the following elements contains all transition metals?
a) Ti, Mo, Pd, Ag
b) Zr, Cr, Ni, As
c) Sc, V, Ba, Fe
d) Ti, Mn, Cu, Se
c. Learning objective: know how to 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.
A group is a vertical column and a period is a row in the periodic table. Alkali metals (Group 1) like Na
and K are soft, white, and extremely reactive chemically. Alkaline earth metals (Group 2) like Mg and
Ca are more solid and less reactive than Group 1 metals. The transition metals (Groups 3-12) are
represented by some of the most common metals such as Cu, Hg, Fe, Ag, and Zn. These elements all
have electrons in their outer D orbitals.
Electronegativity is a measure of the ability of an atom to attract electrons toward itself in a chemical
bond. The values of electronegativity vary from one or less for the Group 1 metals to about 4 for fluorine.
Ionization energy is the energy it takes to remove an electron from an atom to form an ion. An element
often has multiple ionization energies, which corresponds to the energy required to remove the 1 st, 2nd, or
3rd electrons (in order) from an atom. Electronegativity and ionization energy increase from left to right
across a period because of increasing numbers of protons, and decrease down a group because of the
increasing distance between the outer electrons and the nucleus. Atomic size and ionic sizes
decrease from left to right across a period and increase down a group.
Q1 – Which group of the following elements contains all alkaline earth metals?
a) Li, Be, B, C
b) Li, Na, Rb, Cs
c) Be, Mg, Ca, Sr
d) Mg, Ca, Sr, Y
Q2 – Of the three elements fluorine, bromine, and calcium, which has the highest and which has the
lowest ionization energy?
a) Bromine has the highest and calcium has the lowest.
b) Calcium has the highest and fluorine has the lowest.
c) Fluorine has the highest and bromine has the lowest.
d) Fluorine has the highest and calcium has the lowest.
Q3 – Rank the following five elements in order of increasing atomic radius (smallest to largest): Si, Na K,
F, and O
a) K < Na < Si < O < F
b) F < O < Si < Na < K
c) O < F < Si < Na < K
d) O < F < Na < Si < K
Q4 – Arrange the set of elements in order of increasing ionization energy. [Rb, Cl, K , Al]
2
d. Learning objective: know how to use the periodic table to determine the number of electrons
available for bonding.
Only electrons in the highest occupied energy level (valence electrons) are available for bonding. All
elements in the same group have the same number of valence electrons, and hence similar chemical
properties. The number of valence electrons corresponds to the group number, e.g. Group 1 metals
have one valence electron, Group 2 metals have two, etc. All atoms prefer to achieve a low energy
noble gas configuration of 8 valence when bonding with another atom (octet rule). Metals usually lose
electrons and nonmetals gain electrons when forming chemical bonds.
Q1 – The number of valence electrons in the Group that includes nitrogen, phosphorus, and arsenic is:
a) 5
b) 4
c) 3
d) 7
Q2 – Which element has the same number of electrons available for bonding as oxygen?
a) nitrogen
b) sulfur
c) chlorine
d) arsenic
e. Learning objective: know the nucleus of the atom is much smaller than the atom yet contains
most of its mass.
Most of the mass of an atom consists of protons and neutrons, which are densely packed, in the nucleus.
The electrons occupy a large region of space around the nucleus, which defines the volume of an atom.
The mass of an electron is approximately 1/1800 that of a proton or neutron.
Q1 – Consider the following three statements:
I. The mass of a proton and a neutron are virtually identical.
II. The charge of an electron and a proton are equal but opposite.
III. The mass of the electrons is a small fraction of the total mass of any atom.
a) all are correct.
b) I and II are correct but III is false.
c) II and III are correct but I is false.
d) I and III are correct but II is false.
g.* Learning objective: know how to 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.
Q1 – Draw the orbital diagram showing valence electrons, and write the condensed ground-state
electron configuration for chlorine.
j.* Learning objective: know that spectral lines are the result of transitions of electrons between
energy levels and that these lines correspond to photons with a frequency related to the energy
spacing between levels by using Planck’s relationship (E = hv).
Q1 – A radio wave has a frequency of 3.6 x 1010 Hz. What is the energy (in J) of one photon of this
radiation?
Q2 – An FM station broadcasts music at 93.5 MHz (megahertz or 10 6 Hz). Find the wavelength in
meters of these radiowaves.
Chemical Bonds (Chapters 9 & 10)
2. Biological, chemical, and physical properties of matter result from the ability of atoms to form bonds
from electrostatic forces between electrons and protons and between atoms and molecules.
a. Learning Objective: know atoms combine to form molecules by sharing electrons to form covalent or
metallic bonds or by exchanging electrons to form ionic bonds.
3
Atoms (usually nonmetals) of similar electronegativities can form covalent bonds to become molecules.
In a covalent bond, therefore, bonding electron pairs are localized in the region between the bonded
atoms. In metals valence electrons are not localized to individual atoms but are free to move to
temporarily occupy vacant orbitals on adjacent metal atoms. For this reason metals conduct electricity
well.
When an electron from an atom with low electronegativity (e.g., a metal) is removed by another atom
with high electronegativity (e.g., a nonmetal), the two atoms become oppositely charged ions that attract
each other, resulting in an ionic bond. Chemical bonds between atoms can be almost entirely covalent,
almost entirely ionic, or in between these two extremes. The triple bond in nitrogen molecules (N2) is
nearly 100 percent covalent. A salt such as sodium chloride (NaCl) has bonds that are nearly completely
ionic. However, the electrons in gaseous hydrogen chloride are shared somewhat unevenly between the
two atoms. This kind of bond is called polar covalent.
b. Learning Objective: know chemical bonds between atoms in molecules such as H2, CH4, NH3,
H2CCH2, N2, Cl2 and many large biological molecules are covalent.
Organic and biological molecules consist primarily of carbon, oxygen, hydrogen, and nitrogen. These
elements share valence electrons to form bonds so that the outer electron energy levels of each atom
are filled and have electron configurations like those of the nearest noble gas element. (Noble gases, or
inert gases, are in the last column on the right of the periodic table.) For example, nitrogen has one lone
pair and three unpaired electrons and therefore can form covalent bonds with three hydrogen atoms to
make four electron pairs around the nitrogen. Carbon has four unpaired electrons and combines with
hydrogen, nitrogen, and oxygen to form covalent bonds sharing electron pairs.
c. Learning Objective: know salt crystals, such as NaCl, are repeating patterns of positive and negative
ions held together by electrostatic attraction.
The energy that holds ionic compounds together, called lattice energy, is caused by the electrostatic
attraction of cations, which are positive ions, with anions, which are negative ions. To minimize their
energy state, the ions form repeating patterns that reduce the distance between positive and negative
ions and maximize the distance between ions of like charges.
e. Learning Objective: know how to draw Lewis dot structures.
A Lewis dot structure shows how valence electrons and covalent bonds are arranged between atoms in
a molecule. Lewis dot diagrams provide a method for predicting correct combining ratios between atoms
and for determining aspects of chemical bonds, such as whether they are covalent or consist of single,
double, or triple bonds.
Sample Problems 10.1, 10.2, 10.3 in text (pps. 308 & 309)
f. *Learning Objective: know how to predict the shape of simple molecules and their polarity from Lewis
dot structures.
Sample Problems 10.6, 10.7, 10.8, 10.9 in text (p. 322, 324, & 325)
g.* Learning Objective: know how electronegativity and ionization energy relate to bond formation.
Sample Problem 9.4 in text (p. 298)
Conservation of Matter and Stoichiometry (Chapter 3 & 4)
3. The conservation of atoms in chemical reactions leads to the principle of conservation of matter and
the ability to calculate the mass of products and reactants.
a. Learning objective: know how to describe chemical reactions by writing balanced equations.
Reactions are described by balanced equations because all atoms of the reactants must be included in
the products (law of conservation of mass). The simplest way to balance an equation is by inspection.
4
When balancing the atoms on each side of the equation, only change the coefficient of the molecule and
not the subscript of the atoms making up the molecule.
Q1 – Balance the following chemical equation with the smallest whole number coefficients.
___KOH + ___CO2  ___K2CO3 + ___H2O
a) 2:1:1:4
b) 1:1:1:1
c) 2:1:1:2
d) 2:1:1:1
Q2 – Balance the following chemical equation with the smallest whole number coefficients. What is the
value of the coefficient of the CO2 species:
___C4H10 + ___O2  ___CO2 + ___H2O
a) 5
b) 4
c) 10
d) 8
b. Learning objective: know the quantity one mole is set by defining one mole of carbon 12
atoms to have a mass of exactly 12 grams.
The mole is a number just like a “dozen” eggs. It is defined as the number of atoms in 12 grams of
carbon-12. To assign atomic masses to elements, the mass of 12 grams of carbon-12 was selected as a
standard reference to which the masses of all other elements are compared.
Q1 – How many moles are in 18 grams of carbon-12?
a) 12
b) 1.5
c) 2.0
d) 0.5
c. Learning objective: know one mole equals 6.02 x 1023 particles (atoms or molecules).
The number of atoms in 12.0 grams of carbon-12 is defined as one mole, which is equal to 6.02 x 1023
atoms (Avogadro’s number).
Q1 – How many atoms are in 36.0 grams of carbon-12?
a) 2.01 x 1023
b) 1.81 x 1024
c) 7.22 x 1024
d) 1.12 x 1025
Q2 – How many moles of atoms are in 3.0 moles of CH4 molecules?
a) 3.0
b) 5.0
c) 15.0
d) 17.0
d. Learning objective: know how to determine the molar mass of a molecule from its chemical
formula and a table of atomic masses and how to convert the mass of a molecular substance to
moles, number of particles, or volume of gas at standard temperature and pressure.
The molar mass of a compound is the sum of the atomic masses of the atoms making up the compound,
expressed in grams. For mole to mass conversions use the formula m = MM x n where m = mass (g),
MM = molar mass (g/mole), and n = number of moles. The number of particles (atoms or molecules) in a
sample is calculated by multiplying the number of moles by Avogadro’s number. The volume of an ideal
gas at a fixed temperature and pressure is proportional to the number of moles of gas (1.0 mol = 22.4 L
at STP).
Q1 – Identify the INCORRECT statement below:
a) One mole of water contains 6.022 x 1023 H2O molecules.
b) There are 16.00 g in 1.00 mole of O2 molecules.
c) 12.01 grams of carbon have the same number of atoms as 14.01 grams of nitrogen.
d) A mole is the amount of a substance having the same number of particles as 12 grams of pure
carbon-12.
Q2 – How many hydrogen atoms are present in 27.0 g of H2O?
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a) 1.80 x 1024
b) 2.71 x 1024
c) 9.03 x 1023
d) 6.02 x 1023
Q3 – Calculate the molar mass of each of the following: (a) (NH4)3PO4; (b) CH2Cl2; (c) CuSO4•5H2O; and
(d) BrF5
Q4 – How many moles are there in 37.0 L of H2 gas at STP?
e. Learning objective: know to 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.
Atoms are neither created nor destroyed in a chemical reaction (law of conservation of mass). When a
chemical reaction is written as a balanced equation, the mass of any one product or reactant can be
calculated using stoichiometric ratios if the mass of just one reactant or product is known. The
coefficients in a balanced chemical equation are mole quantities and NOT masses.
Q1 – The rusting process is described by the reaction 4Fe(s) + 3O 2(g)  2Fe2O3(s). Determine the
maximum amount of Fe2O3(s) that can be produced from 56.0 g of Fe and unlimited oxygen.
a) 40.2 g
b) 80.1 g
c) 120.0 g
d) 160.0 g
Q2 – Chromium (III) oxide reacts with hydrogen sulfide gas to form chromium (III) sulfide and water
according to the following reaction:
Cr2O3(s) + 3H2S(g)  Cr2S3(s) + 3H2O(l)
To produce 421 g of Cr2S3 (a) how many moles of Cr2O3 are required? (b) How many grams of Cr2O3 are
required?
Gases and Their Properties (Chapter 5)
4. The kinetic molecular theory describes the motion of atoms and molecules and explains the
properties of gases.
a. Learning objective: know the random motion of molecules and their collisions with a surface
create the observable pressure on that surface.
The kinetic molecular theory states that particles of all substances are in constant motion. The theory
assumes: (a) gases are mostly empty space, i.e. the gas particles have negligible volume, (b) the gas
particles are in constant motion, (c) collisions between gas particles are elastic (no energy is lost), and
(d) gas pressure is caused by collisions of molecules with the walls of the container.
Q1 – If 4 moles of gas are added to a container that already holds 1 mole of gas, how will the pressure
change inside the container?
a) The pressure will be five times as great.
b) The pressure will be two times as great.
c) The pressure will be four times as great.
d) The pressure will not change.
Q2 – Increasing the volume of a given amount of gas at constant temperature causes the pressure to
decrease because ____________.
a) the same number of molecules are striking a larger area
b) there are fewer molecules
c) the molecules are moving more slowly
d) there are more molecules
b. Learning objective: know the random motion of molecules explains the diffusion of gases.
6
Gases diffuse into each other to form a homogeneous mixture. The rate of diffusion of gases is inversely
proportional to their molar mass.
Q1 – At STP which of the following gases would have the lowest average molecular speed?
a) NH3
b) CO2
c) Ar
d) H2
c. Learning objective: know how to apply the gas laws to relations between the pressure,
temperature, and volume of any amount of an ideal gas or any mixture of ideal gases.
Solve all gas law problems with the Combined Gas Law that is on your Chemistry Reference Sheet.
Q1 – A 2.0 L sample of O2 gas is initially at STP. The volume is doubled while holding the temperature
constant. What is the final pressure in torr units?
a) 380 torr
b) 760 torr
c) 0.5 torr
d) 2 torr
Q2 – The combined gas law relates which of these?
a) pressure and volume only
b) temperature and pressure only
c) volume and temperature only
d) temperature, pressure and volume
Q3 – A 2.0 liter sample of gas initially at 25oC is heated to 100oC at fixed pressure. What is its final
volume?
a) 8.0 L
B) 0.5 L
c) 2.5 L
d) 4.7 L
d. Learning objective: know values and meanings of standard temperature and pressure (STP).
In order to compare volume of gases, the temperature and pressure must be specified. Standard
Temperature and Pressure (STP) are defined as 0oC and 1 atm.
Q1 – The values of standard temperature and pressure are
a) 0 oC, 1 atm
b) 0 K, 1 atm
c) 0 oC, 760 atm
d) 0 K, 760 atm
e. Learning objective: know how to convert between the Celsius and Kelvin temperature scales.
Use the formula K = oC + 273 to convert between Celsius and Kelvin temperature scales.
Q1 – Convert 574 K to Celsius.
a) 273 oC
b) 847 oC
c) 301 oC
d) 574 oC
f. Learning objective: know there is no temperature lower than 0 Kelvin.
In theory, the lowest temperature than anything can be cooled to is 0 K or -273 oC. The Kelvin scale
starts at absolute zero (0 K). Note: All gas law calculations must be performed with temperature in units
of Kelvin (K).
g.* Learning Objective: the kinetic theory of gases relates the absolute temperature of a gas to
the average kinetic energy of its molecules or atoms.
Q1 – Use the kinetic molecular theory to explain the change in pressure that results from warming a
sample of gas.
Q2 – How does the kinetic molecular theory explain why 1 mole of krypton and 1 mole of helium have
the same volume at STP?
7
h.* Learning objective: know how to solve problems using the ideal gas law in the form PV=nRT.
Q1 – A 75.0-g sample of dinitrogen monoxide is confined in a 3.1-L vessel. What is the pressure (in atm)
at 115oC?
i.* Learning Objective: know how to apply Dalton’s law of partial pressures to describe the composition
of gases and Graham’s law to predict diffusion of gases.
Sample Problem 5.10 in text (p. 164)
Chemical Thermodynamics (Chapter 6)
7. Energy is exchanged or transformed in all chemical reactions and physical changes of matter.
a. Learning objective: know how to describe temperature and heat flow in terms of the motion of
molecules (or atoms).
Temperature is a measure of the average kinetic energy of molecular motion in a sample. Heat is the
transfer of thermal energy from a point of higher temperature to one at a lower temperature. Heat is
described as flowing from the system to the surroundings.
Q1 - Which transfer of energy occurs when ice cubes are placed in water that has a temperature of
45oC?
a) Chemical energy is transferred from the ice to the water.
b) Chemical energy is transferred from the water to the ice.
c) Thermal energy is transferred from the ice to the water.
d) Thermal energy is transferred from the water to the ice.
b. Learning objective: know chemical processes can either release (exothermic) or absorb
(endothermic) thermal energy.
Endothermic processes absorb heat from the surroundings (∆H is positive) and the thermochemical
equation is written with heat as a reactant. Exothermic processes release heat (∆H is negative) to the
surroundings and the thermochemical equation is written with heat as a product. The net heat that is
released or absorbed comes from the breaking of chemical bonds in the reactants (absorbs energy) and
making of chemical bonds in the products (releases energy).
Q1 – Which of the following processes are exothermic: condensing steam, evaporating alcohol, burning
alcohol, or baking a potato.
a) condensing steam and burning alcohol
b) burning alcohol and baking a potato.
c) evaporating alcohol
d) baking a potato
Q2 – Consider the following balanced thermochemical equation for the decomposition of the mineral
magnesite:
MgCO3(s)  MgO(s) + CO2(g)
∆Hrxn = 117.3 kJ
a) Is heat absorbed or released in the reaction?
b) What is ∆Hrxn for the reverse reaction?
c) What is ∆H when 5.35 mol of CO2 reacts with excess MgO?
d) What is ∆H when 35.5 g of CO2 reacts with excess MgO?
c. Learning objective: know energy is released when a material condenses or freezes and is
absorbed when a material evaporates or melts.
The three physical states are gas, liquid, and solid. Changes between states either absorb or release
8
heat (thermal energy). Evaporation and melting require energy to overcome the intermolecular forces of
attraction holding the substance together. Condensation and freezing release heat to the surroundings
as the internal energy is reduced, allowing the intermolecular forces to bond the particles together.
Q1 – Which phase change is an exothermic process?
a) H2O(s)  H2O (g)
b) H2O(g)  H2O (l)
c) H2O(s)  H2O (l)
d) H2O (l)  H2O (g)
d. Learning objective: know how to solve problems involving heat flow and temperature
changes, using known values of specific heat and latent heat of phase change.
Specific heat is defined as the energy needed to change the temperature of one gram of a substance by
one degree Celsius. Use the formula Q = m(∆T)Cp to calculate any one of the four variables. During
phase changes, energy is added or removed from the system without a corresponding change is
temperature. This is called latent (hidden) heat. Use the formula Q = m∆Hfus. You should be able to
draw a heating/cooling curve for a substance.
Q1 – Copper has a specific heat of 0.38 J/g oC. If 2.5 g of copper absorbs 2.8 J of heat, what is the
change in temperature?
a) 0.42 oC
b) 1.6 oC
c) 2.8 oC
d) 4.1 oC
Q2 – How many grams of ice at 0 oC and 1 atm could be melted by the addition of 2.2 kJ of heat? The
∆Hfus of ice is 6.0 kJ/mol.
a) 2.7 g
b) 4.7 g
c) 5.2 g
d) 6.7 g
Q3 – Calculate the heat (q) released when 0/10 g of ice is cooled from 10. oC to – 75.oC (cice = 2.087
J/goC).
e.* Learning objective: know how to apply Hess's Law to calculate enthalpy change in a reaction.
Q1 – Calculate ∆Hrxn for the reaction: 2 C2H6(g) + 7 O2(g)  4 CO2(g) + 6 H2O(g). [ H of CO2(g) = - 393.5
kJ/mol; H of H2O(g) = - 241.862 kJ/mol; H of C2H6(g) = - 84.667 kJ/mol]
9
Study Guide Answer Key
Atomic and Molecular Structure
a) Q1
a) Q2
a) Q3
a) Q4
a) Q5
d) 4,4,5
c) protons plus the number of neutrons
a) 27, 58.93
c) 33
b) 14
b) Q1
b) Q2
b) Q3
b) Q4
c) Mg, Ti, Fe, Zn
c) metalloids (semimetals)
b) halogens
a) Ti, Mo, Pd, Ag
c) Q1
c) Q2
c) Q3
c) Q4
c) Be, Mg, Ca, Sr
d) Fluorine has the highest and calcium has the lowest.
b) F < O < Si < Na < K
[Rb < K < Al < Cl]
d) Q1
d) Q2
a) 5
b) sulfur
e) Q1
a) all are correct.
g) Q1
Cl (Z = 17); [Ne]3s23p5
Ne
3s
3p
j*) Q1
E = h
E = (6.626 x 10–34 J•s) (3.6 x 1010 s–1) = 2.385 x 10–23 = 2.4 x 10–23 J
j*) Q2
(m) =
c
=

3.00 x 108 m/s
= 3.208556 = 3.21 m
 106 Hz  s 1 
 93.5 MHz  



 1 MHz  Hz 
Conservation of Matter and Stoichiometry
a) Q1
a) Q2
d) 2:1:1:1
d) 8
b) Q1
b) 1.5
c) Q1
c) Q2
b) 1.81 x 1024
c) 15.0
d) Q1
d) Q2
b) There are 16.00 g in 1.00 mole of O2 molecules.
a) 1.80 x 1024
d) Q3
a) (NH4)3PO4 = 3(14.01) + 12(1.008) + 30.97 + 4(16.00) = 149.10 g/mol
b) CH2Cl2 = 12.01 + 2(1.008) + 2(35.45) = 84.93 g/mol
c) CuSO4•5H2O = 63.55 + 32.07 + 9(16.00) + 10(1.008) = 249.70 g/mol
d) BrF5 = 79.90 + 5(19.00) = 174.90 g/mol
10
d Q4)
e) Q1
1.65 mol H2
b) 80.1 g
e) Q2
 1 mol Cr2S3  1 mol Cr2O3 
a) Moles Cr2O3 =  421 g Cr2S3  

 = 2.10279 = 2.10 mol Cr2O3
 200.21 g Cr2S3  1 mol Cr2S3 
 1 mol Cr2S3  1 mol Cr2 O3  152.00 g Cr2O3 
b) Grams Cr2O3 =  421 g Cr2S3  



 200.21 g Cr2S3  1 mol Cr2S3  1 mol Cr2 O3 
= 319.624 = 3.20 x 102 g Cr2O3
Gases and Their Properties
a) Q1
a) Q2
a) The pressure will be five times as great.
a) the same number of molecules are striking a larger area
b) Q1
b) CO2
c) Q1
c) Q2
c) Q3
a) 380 torr
d) temperature, pressure and volume
c) 2.5 L
d) Q1
a) 0 oC, 1 atm
e) Q1
c) 301 oC
g*) Q1
As the temperature of the gas sample increases, the most probable speed increases.
This will increase both the number of collisions per unit time and the force of each
collision with the sample walls. Thus, the gas pressure increases.
g*) Q2
At STP (or any identical temperature and pressure), the volume occupied by a mole of
any gas will be identical. This is because at the same temperature, all gases have the
same average kinetic energy, resulting in the same pressure.
h*)
PV = nRT
 1 mol N 2 O 
 = 1.70377 mol N2O
n = 75 .0 g N 2 O
 44.02 g N 2 O 
L atm
(1.70377 mol )(0.0821
)(388 K)
nRT
mol
K
P=
= 17.5075 =18atm N2O

V
(3.1 L)
Chemical Thermodynamics
a) Q1
d) Thermal energy is transferred from the water to the ice.
b) Q1
a) condensing steam and burning alcohol
b) Q2
a) Absorbed
b) ∆Hrxn (reverse) = –117.3 kJ
 117.3 kJ 
c) ∆Hrxn =  5.35 mol CO2  
 = –627.555 = –628 kJ
 1 mol CO2 
 1 mol CO2  117.3 kJ 
d) ∆Hrxn =  35.5 g CO2  

 = –94.618 = –94.6 kJ
 44.01 g CO2  1 mol CO 2 
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c) Q1
b) H2O(g)  H2O (l)
d) Q1
d) Q2
c) 2.8 oC
d) 6.7 g
d) Q3
e*) Q1

J 
q(J) = c x mass x ∆T =  2.087
  0.10 g    75  10. C = –17.7 = –18 J
g C 

2 C2H6(g) + 7 O2(g)  4 CO2(g) + 6 H2O(g)
o
H rxn
= m[ H of (products)] – n[ H of (reactants)]
o
H rxn
= {4 mol ( H of , CO2(g)) + 6 mol ( H of , H2O(g))} – {2 mol ( H of , C2H6(g)) + 7 mol
( H of , O2(g))}
kJ
kJ ) – [2 mol (–84.667 kJ ) + 7 mol (0)]
) + 6 mol (–241.826
mol
mol
mol
= –2855.6 kJ (or –1427.8 kJ for reaction of 1 mol of C2H6)
= 4 mol (–393.5
Revised:
January 22, 2014
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