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Chm.1.1 Analyze the structure of atoms and ions.
Chm.1.1.1 Analyze the structure of atoms, isotopes, and ions.
Chm.1.1.2 Analyze an atom in terms of the location of electrons.
Chm.1.1.3 Explain the emission of electromagnetic radiation in spectral form in terms of the Bohr model.
Chm.1.1.4 Explain the process of radioactive decay using nuclear equations and half-life.
name
symbol
charge
relative
mass
location
proton
p+
+1
1
in nucleus
neutron
no
0
1
in nucleus
electron
e-
-1
~1/2000
out of
nucleus
The atomic number (Z) of an element is the
number of protons in the nucleus of an atom of
that element. The number of protons
determines identity of an element, as well as
many of its chemical and physical properties.
The sum of the protons and neutrons in the
A
nucleus is the mass number (A) of that
particular atom.
Z
X
DALTON’S ATOMIC THEORY
BOHR MODEL OF THE ATOM
1) All matter is made of tiny indivisible particles
called atoms.
2) Atoms of the same element are identical; those of
different atoms are different.
3) Atoms of different elements combine in whole
number ratios to form compounds
4) Chemical reactions involve the rearrangement of
atoms. No new atoms are created or destroyed.
1. An electron circles the nucleus only in fixed energy ranges called orbits.
2. An electron can neither gain nor lose energy inside this orbit, but could
move up or down to another orbit.
3. The lowest energy orbit is closest to the nucleus.
CURRENT MODEL OF THE ATOM
Atoms are composed of electrons in a cloud around a positive nucleus.
Balanced nuclear equations require that both the atomic number and the mass number must be
balanced. The atomic number (the number on the bottom) determines the identity of the element.
235
92 U

4
2
231
He + 90 Th
FUSION – 2 or more small nuclei combine to one larger nucleus
FISSION – 1 large nucleus splits to form two or more smaller nuclei
A half-life is the time required for one-half
of a radioisotope’s nuclei to decay into its
products. The half-life of any particular
radioisotope is constant and therefore
cannot be sped up.
Frequency and wavelength are inversely
related. When one is high, the other is low.
Energy and frequency are directly related.
Bohr’s model of the atom helped to
explain spectral lines.
There are 2 electrons
in the first energy
level (orbit). There
are a maximum of 8
e- in the 2nd orbit, and
18 in the 3rd orbit.
time
amount
0
Initial amount
Half-life
Initial amount ÷ 2
Half-life x 2
Initial amount ÷ 4
Half-life x 3
Initial amount ÷ 8
Atoms of an element that are chemically alike but differ in mass are called isotopes of
the element. Isotopes of an element have different mass numbers because they have
different numbers of neutrons, but they all have the same atomic number.
Electron configurations represent the way electrons are arranged in atoms. Electrons
enter the lowest energy first. At most there can be only 2 electrons per orbital, and they
must have opposite “spins.”
Electron configuration for iron (Fe): 1s2 2s2 2p6 3s2 3p6 4s2 3d6
2
6
NobleExam
gas configuration
Chemistry 1 Final
Review for iron (Fe): [Ar] 4s 3d
Page 1
Chm.1.1 Analyze the structure of atoms and ions.
1. Which idea of John Dalton is no longer considered part of the modern view of atoms?
A. Atoms are extremely small.
B. Atoms of the same element have identical masses.
C. Atoms combine in simple whole number ratios to form compounds.
D. Atoms of different elements can combine in different ratios to form different compounds.
2. Which best describes the current atomic theory?
A. Atoms consist of electrons circling in definite orbits around a positive nucleus.
B. Atoms are composed of electrons in a cloud around a positive nucleus.
C. Atoms can easily be split, at which time they become radioactive.
D. An atom’s mass is determined by the mass of its neutrons.
3. What is the nuclear composition of uranium-235?
A. 92 electrons + 143 protons
C. 143 protons + 92 neutrons
B. 92 protons + 143 electrons
D. 92 protons + 143 neutrons
4. Which best describes the relationship between subatomic particles in any neutral atom?
A. The number of protons equals the number of electrons.
B. The number of protons equals the number of neutrons.
C. The number of neutrons equals the number of electrons.
D. The number of neutrons is greater than the number of protons.
5. Consider the spectrum for the hydrogen atom. In which situation will light be produced?
A. Electrons absorb energy as they move to an excited state.
B. Electrons release energy as they move to an excited state.
C. Electrons absorb energy as they return to the ground state.
D. Electrons release energy as they return to the ground state.
6. Which statement regarding red and green visible light is correct?
A. The speed of green light is greater than that of red light.
B. The wavelength of green light is longer than that of red light.
C. The energy of green light is lower than that of red light.
D. The frequency of green light is higher than that of red light.
7. Which color of light would a hydrogen atom emit when an electron changes from the n = 5 level to the
n = 2 level?
A. red
B. yellow
C. green
D. blue
8. What energy level transition is indicated when the light emitted by a hydrogen atom has a wavelength of 103 nm?
A. n = 2 to n = 1
B. n = 3 to n = 1
C. n = 4 to n = 2
D. n = 5 to n = 2
9. Consider this diagram:
Which of the three types of radiation will penetrate the paper and wood?
A. alpha, beta, gamma
B. alpha and beta only
C. gamma only
D. beta only
10. Given the electronic configuration of 1s2 2s2 2p4 , how many electrons does this element have in its outer level?
A. 2
B. 4
C. 6
D. 8
11. The half-life of phosphorus-32 is 14.30 days. How many milligrams of a 20.00 mg sample of phosphorus-32 will remain
after 85.80 days?
A. 3.333 mg
B. 0.6250 mg
C. 0.3125 mg
D. 0.1563 mg
Chemistry 1 Final Exam Review
Page 2
12. Identify the element that has the abbreviated electron configuration given by [Ne] 3s 2.
a) O
b) Mn
c) Ca
d) Mg
13. In the figure below, what type of nuclear activity is represented?
A. fission
B. fusion
C. alpha emission
14. Which particle will complete this reaction?
A. electron
A.
0
1
e
60
Co  ?27
Co
59
27
B. neutron
15. Which will complete this equation?
B.
0
0
D. beta emission
C. nucleus
D. proton
U 234
90Th  ___
238
92

C.
16. Nuclear decay is a random event and
A. is independent of other energy influences
C. can be sped up upon cooling
1
1
H
D.
4
2
He
B. can be sped up upon heating
D. can be sped up at room temperature
Chm.1.2 Understand the bonding that occurs in simple compounds in terms of bond type, strength,
and properties.
Chm.1.2.1 Compare (qualitatively) the relative strengths of ionic, covalent, and metallic bonds.
Chm.1.2.2 Infer the type of bond and chemical formula formed between atoms.
Chm.1.2.3 Compare inter- and intra- particle forces.
Chm.1.2.4 Interpret the name and formula of compounds using IUPAC convention.
Chm.1.2.5 Compare the properties of ionic, covalent, metallic, and network compounds.
Metals lose electrons to attain noble
gas configurations. They make
positive ions, cations. The electron
that is removed comes from the
highest energy level. Nonmetals gain
electrons to attain noble gas
configurations. This means they want
an octet of electrons, 8 electrons. They
make negative ions, anions.
As atoms bond
with each other,
they decrease their
potential energy,
thus creating more
stable
arrangements of
matter. Energy is
released when the
bond forms.
A single bond is formed from the sharing of
two valence electrons. Cl2 has a single bond.
A double bond occurs when atoms share
two pairs of electrons, 4 electrons. O2 has a
double bond. A triple bond forms when
atoms share three pairs of electrons, 6
electrons. N2 has a triple bond. Triple bonds
are the strongest and shortest and have the
highest bond energy. Bond energy is the
energy required to break a bond.
Macromolecules have large numbers of atoms linked by covalent bonds in chains or sheets (such as graphite), or in
3-dimensional structures (such as diamond and quartz). Macromolecules are in your hair and fingernails. Man-made
macromolecules include polymers like PVC and nylon. A network solid is a macromolecule in which the atoms are
bonded covalently in a continuous network. Examples include diamond, graphite, and quartz.
COVALENT BONDING
Covalent compounds occur between two nonmetals or a nonmetal and hydrogen. In a covalent bond, atoms share
electrons and neither atom has an ionic charge. The electronegativity difference of the bonding elements is less
than 1.7. When hydrogen bonds with a nonmetal the bond is covalent. Characteristics: low MP, low BP, poor
electrical conductivity, and polar nature.
Chemistry 1 Final Exam Review
Page 3
IONIC BONDING
Metallic Bonds
Anions and cations are involved in ionic bonding
and are held together by opposite charges,
electrostatic attraction. The bond is formed
through the transfer of electrons. Ionic bonds
occur between metals and nonmetals. The
electronegativity difference of the bonding
elements is greater than 1.7. Characteristics:
high MP, high BP, brittle, and high electrical
conductivity either in molten state or in aqueous
solution.
The bonding in metals is explained by the electron sea model, which
proposes that the atoms in a metallic solid contribute their valence
electrons to form a “sea” of electrons that surrounds metallic cations.
The bond that results from this shared pool of valence electrons is
called a metallic bond.

Metals conduct electricity.

Metals generally have extremely high melting and boiling
points.

Metals are malleable (able to be hammered into sheets).

Metals are also ductile (able to be drawn into wire).

Metals have luster (are shiny).
A mixture of elements that has metallic properties is called an alloy.
ACIDS
Acids will always contain one or more hydrogen ions
next to an anion. The anion determines the name of
the acid.
Binary acids contain hydrogen and an anion whose
name ends in –ide. When naming the acid, use the
prefix hydro- and change -ide to -ic acid.
 Example: HCl = hydrochloric acid
The acid is a ternary acid if the anion has oxygen in
it. The anion ends in -ate or -ite. Change the suffix ate to –ic acid. Change the suffix -ite to -ous acid.
The hydro- prefix is NOT used!
 Example: HNO3 = nitric acid
HNO2 = nitrous acid
Naming Binary Ionic Compounds
Binary ionic compounds are composed of a metal bonded with a
nonmetal.
 Name the metal ion using a Roman numeral in parenthesis if
necessary.
 Follow this name with the name of the nonmetal ion.
Examples: Ca3P2 = calcium phosphide
PbCl4 = lead (IV) chloride
Naming Ternary Ionic Compounds
Ternary ionic compounds are composed of at least 3
elements. Name the metal ion, using a Roman numeral in
parenthesis if necessary.
Follow this name with the name of the polyatomic ion.
 Examples: Na3PO4 = sodium phosphate
Cu2SO4 = copper (I) sulfate
Zn2+
Ag1+
MOLECULAR COMPOUNDS
Molecular compounds are made of 2 nonmetals. A
molecular compound’s name tells you the number of atoms
through the use of prefixes.
1 mono- 2 di3 tri4 tetra5 penta6 hexa- 7 hepta8 octa- 9 nona10 decaThe name will consist of two words. One exception is we
don’t write mono- if there is only one of the first element.
You will not need to criss-cross oxidation numbers.
 Examples: SO2 = sulfur dioxide;
N2O5 = dinitrogen pentoxide
Intermolecular forces are forces of attraction. They
are what make solid and liquid molecular compounds
possible. Solids have higher intermolecular forces than
liquids, and liquids have higher intermolecular forces
than gases. The three intermolecular forces are hydrogen
bonds, dipole–dipole forces and London dispersion
forces.
 H-bond is an attraction between molecules when H
is bonded to N, O, or F.
 Dipole-dipole attractions occur between polar
molecules.
 London dispersion forces (also called van der Waals)
occurs when electrons of one molecule attracted to
nucleus of another molecule, i.e. liquefied inert
gases.
Relative strengths: H > dipole > London dispersion.
A polar molecule has uneven charge distribution; it has a
region of partial positive charge and a region of partial
negative charge.
Chemistry 1 Final Exam Review
Page 4
VSEPR SHAPES AND ANGLES
LINEAR
180°
EITHER only 2 elements OR central element surrounded by 2
elements and NO dots
TRIGONAL PLANAR
120°
central element surrounded by 3 elements and NO dots
BENT
<120°
central element surrounded by 2 elements and ONE pair of dots
TETRAHEDRAL
109.5°
central element surrounded by 4 elements and NO dots
TRIGONAL PYRAMIDAL
107°
BENT
104.5°
central element surrounded by 3 elements and ONE pair of dots
central element surrounded by 2 elements and TWO PAIRS of dots
Polar and Nonpolar Covalent Bonds and Molecules
VSEPR
LINEAR
When the atoms in a bond are the same, the electrons are shared equally. This
results in a nonpolar covalent bond. Diatomic elements (Br2, I2 N2, Cl2, H2,
O2 and F2) have pure nonpolar covalent bonds. Nonpolar molecules are
symmetrical because there are no unshared electrons around the central atom.
180°
VSEPR stands for Valence Shell Electron Pair
Repulsion.
It predicts the three120°
dimensional
TRIGONAL
geometry of molecules. The electron pairs try to
get asPLANAR
far away as possible to minimize repulsion.
VSEPR is based on the number of pairs of valence
electrons,
both bonded and unbonded.
BENT
<120° An
unbonded pair of electrons is referred to as a
TETRAHEDRAL
109.5°
lone pair.
TRIGONAL
107°
BENT
104.5°
The bond that forms when electrons are shared unequally is called a polar
covalent bond. There are either unshared electrons around the central atom
or there are different elements bonded with the central element. Polar
molecules are asymmetrical.
17. What is the name of the compound PbO2?
PYRAMIDAL
A. lead oxide
B. lead (II) oxide
C. lead oxide (II)
D. lead (IV) oxide
18. What is the name of HCl (aq)?
LINEAR
180°B. hydrochloric acid
A. chloric acid
C. hydrogen chloride
D. perchloric acid
19. What is the chemical formula for calcium nitrate?
A. CaNO3
B. Ca(NO2)2
C. Ca(NO3)2
D. Ca3N2
20. Which is the correct formula for dinitrogen pentoxide?
A. N4O
B. NO2
C. N2O5
D. NO4
21. What type of bonding is associated with compounds that have the following characteristics?
• high melting points
• conduct electricity in the molten state
• solutions conduct electricity
• normally crystalline solids at room temperature.
A. covalent
B. ionic
C. hydrogen
D. metallic
22. Which pair of elements would most likely bond to form a covalently bonded compound?
A. sodium and fluorine
B. barium and chlorine
C. phosphorus and oxygen
D. magnesium and sulfur
23. Based on the VSEPR theory, what is the molecular geometry of a molecule of PI 3?
A. linear
B. tetrahedral
C. trigonal planar
24. Which is a unique characteristic of the bonding between metal atoms?
A. Atoms require additional electrons to reach a stable octet.
B. Atoms must give away electrons to reach a stable octet.
C. Atoms share valence electrons only with neighboring atoms to reach a stable octet.
D. Delocalized electrons move among many atoms creating a sea of electrons.
Chemistry 1 Final Exam Review
Page 5
D. trigonal pyramidal
25. What type of intermolecular force is found in carbon dioxide, CO 2?
A. hydrogen
B. dipole
C. London dispersion
D. Metallic
26. A triple bond has ________ bond energy and is _______ in length than a double bond.
A. greater, shorter
B. greater, longer
C. smaller, shorter
D. smaller, longer
27. How many lone pair electrons should be added to one nitrogen atom in the following dash diagram for hydrazine?
A. 0
B. 1
C. 2
D. 3
Chm.1.3 Understand the physical and chemical properties of atoms based on their position on the
Periodic Table.
Chm.1.3.1 Classify the components of a periodic table (period, group, metal, metalloid, nonmetal, transition).
Chm.1.3.2 Infer the physical properties (atomic radius, metallic and nonmetallic characteristics) of an element based
on its position on the Periodic Table.
Chm.1.3.3 Infer the atomic size, reactivity, electronegativity, and ionization energy of an element from its position on
the Periodic Table.
PERIODIC TREND
Top to
Bottom
Left to
Right
Atomic Radius (Size)
increase
decrease
Ionization Energy
decrease
increase
Electronegativity
decrease
increase
Metallic Character
increase
decrease
Nonmetallic Character
decrease
increase
PERIODICITY
Metals, Nonmetals and Metalloids
Metals are elements that have luster, conduct heat and
electricity, usually bend without breaking (malleable) and
are ductile. Most have extremely high melting points.
Reactivity increases as you go down within a group for
metals. With metals the greater the tendency to lose
electrons, the more reactive the metal is. Reactive metals
have low ionization energies and low electronegativities.
Most nonmetals don’t conduct electricity, are much
poorer conductors of heat than metals, and are brittle
when solid. Many are gases at room temperature; those
that are solids lack the luster of metals. Their melting
points tend to be lower than those of metals. Reactivity
decreases as you go down within a group for nonmetals.
Smaller nonmetals have greater nuclear charge because
the outer electrons are closer to the nucleus. Thus the
tendency to attract electrons will increase.
Metalloids have some chemical and physical properties
of metals and other properties of nonmetals. In the
periodic table, the metalloids lie along the border between
metals and nonmetals. Metalloids include B, Si, Ge, As,
Sb, Te, and At.
Electronegativity is the tendency for an atom to attract electrons to
itself when it is chemically combined with another element. Nonmetals
attract electrons to achieve an octet (8 valence electrons), so fluorine,
the most active nonmetal, has the highest electronegativity.
Ionization energy (IE) is the amount of energy required to completely
remove an electron from a gaseous atom. Nonmetals do not want their
electrons removed because they gain to achieve an octet, so nonmetals
have high ionization energies.
Atomic size is influenced by two factors.
 Energy Level – A higher energy level is further away.
 Charge on nucleus - More charge (protons) pulls electrons in
closer.
Ionic Size
As you might expect, the loss of electrons produces a
positive ion (a cation) with a radius that is smaller than
that of the parent atom. Metals form cations because they
tend to lose electrons in order to achieve the stability of a
filled or even half-filled octet. Conversely, when an atom
gains electrons, the resulting negative ion (an anion) is
larger than the parent atom. Nonmetals form anions in
order to achieve the stability of a filled octet.
Chemistry 1 Final Exam Review
Page 6
Groups are vertical columns on the periodic table, while periods are horizontal rows. Main group elements (Group
A elements) in the same column have similar properties, the same number of valence electrons, and the same
oxidation number.
Representative Elements/
Main Group/ Group A
Elements are shaded.
Transition Elements are
white.
28. Which electron configuration represents a transition element?
A. 1s22s22p3
B. 1s22s22p63s2
C. 1s22s22p63s23p64s23d7
29. Which correctly lists four atoms from smallest to largest radii?
A. I, Br, Cl, F
B. F, I, Br, Cl
C. Si, P, S, Cl
Chemistry 1 Final Exam Review
Page 7
D. 1s22s22p63s23p64s23d104p4
D. Cl, S, P, Si
30. Which best explains why cations are smaller than the atoms from which they are formed?
A. The metallic atom gains electrons, causing a larger effective nuclear pull.
B. The metallic atom loses electrons, resulting in loss of an entire energy level.
C. The nonmetallic atom gains electrons, causing a larger effective nuclear pull.
D. The nonmetallic atom loses electrons, resulting in loss of an entire energy level.
31. Which have the lowest electronegativities?
A. alkali metals
B. halogens
C. rare earth elements
D. transition metals
32. The compound formed between element X and oxygen has the chemical formula X2O. Which element would X most likely
represent?
A. Fe
B. Zn
C. Ag
D. Sn
33. An element X has the outer ground state electronic configuration ns 2np5. Which statement about X is incorrect?
A. X forms diatomic molecules.
B. X forms an X– ion.
C. Compared to other elements in its period, X has a relatively low first ionization energy.
D. X is in group 17.
34. Which of the following groups of elements contains metals and non-metals?
A. Group 3.
B. Group 2.
C. Group 14.
35. Which element is the odd one out because it does not have a high melting point?
A. Carbon.
B. Molybdenum.
C. Cesium.
D. Group 18.
D. Rhenium.
36. Which list below correctly gives the relative values of ionization energy (IE1) for the elements stated?
A. Ne > F > O < N > C > B
B. Ne < F < O < N < C < B
C. Ne > F > O > N > C > B
D. Ne < F < O > N < C < B
37. Which statement about metallic radii (rmetal) is incorrect?
A. Values of rmetal increase across the first row of the f-block.
C. Values of rmetal increase down group 13 from Al onwards.
B. Values of rmetal increase down group 2.
D. Values of rmetal increase down group 1.
38. Which of the following statements about ionic radii is incorrect?
A. The ionic radius of Fe3+ is less than that of Fe2+.
B. The ionic radius of F– is less than that of Cl–.
2+
C. Values of the ionic radii of M increase down group 2. D. The ionic radius of Br– is less than the covalent radius of Br.
39. Match the highest oxidation state available to one or more elements in a given group to the group number. Which pairing is
incorrect?
A. Oxidation state +4; group 14.
B. Oxidation state +1; group 1.
C. Oxidation state +9; group 9.
D. Oxidation state +8; group 18.
Chm.2.1 Understand the relationship among pressure, temperature, volume, and phase.
Chm.2.1.1 Explain the energetic nature of phase changes.
Chm.2.1.2 Explain heating and cooling curves (heat of fusion, heat of vaporization, specific heat, melting point, and
boiling point).
Chm.2.1.3 Interpret the data presented in phase diagrams.
Chm.2.1.4 Infer simple calorimetric calculations based on the concepts of heat lost equals heat gained and specific
heat.
Chm.2.1.5 Explain the relationships among pressure, temperature, volume, and quantity of gas, both qualitative and
quantitative.
A phase diagram typically has three regions,
each representing a different phase and three
curves that separate each phase. The critical
point indicates the critical pressure and the
critical temperature above which a substance
cannot exist as a liquid. The triple point is
the point on a phase diagram that represents
the temperature and pressure at which three
phases of a substance can coexist. “Normal”
means 1 atm.
Chemistry 1 Final Exam Review
Kinetic energy is energy of motion.
Page 8 Kinetic energy increases as the temperature of
molecules goes up. Chemical systems contain both kinetic energy and potential energy.
40. Which substance listed in the table is a liquid at 27°C?
Melting Point (°C)
Boiling Point (°C)
28
140
I
-10
25
II
20
140
III
-90
14
IV
A. I
B. II
C. III
D. IV
41. What happens to the pressure of a constant mass of gas at constant temperature when the volume is doubled?
A. The pressure is doubled.
B. The pressure remains the same.
C. The pressure is reduced by ½.
D. The pressure is reduced by ¼.
42. The total pressure in a closed vessel containing N2, O2 and CO2 is 30 atm. If the partial pressure of N2 is 4 atm, and the partial
pressure of O2 is 6 atm, what is the partial pressure of CO2?
A. 20 atm
B. 30 atm
C. 40 atm
D. 50 atm
43. What is the pressure, in atmospheres, exerted by a 0.100-mol sample of oxygen in a 2.00-L container at 273°C?
A. 4.48 × 10−1 atm
B. 2.24 × 100 atm
C. 1.12 × 103 atm
D. 2.24 × 103 atm
44. Consider this phase diagram:
45. At what temperature does the normal boiling point occur?
A. 45°C
B. 60°C
C. 100°C
D. 110°C
46. What happens when energy is removed from liquid water?
A. Molecules slow down, and more hydrogen bonds are formed.
B. Molecules slow down, and more hydrogen bonds are broken.
C. Molecules move faster, and more hydrogen bonds are formed.
D. Molecules move faster, and more hydrogen bonds are broken.
47. The gases helium, neon, and argon are in separate containers at 55°C. Which is true about the kinetic energy of the gases?
A. Helium has the lowest mass and therefore greatest kinetic energy.
B. They each have a different kinetic energy.
C. Argon has the greatest mass and therefore the greatest kinetic energy.
D. They all have the same average kinetic energy.
48. A piece of metal is heated in a Bunsen burner flame and then immersed in a beaker of cool water.
Which statement best describes the effect of the temperature changes on the kinetic energy of the particles?
A. Kinetic energy of metal atoms decreases in the flame.
B. Kinetic energy of water molecules increases when the heated metal is immersed.
C. Kinetic energy of water molecules decreases when the heated metal is immersed.
D. Kinetic energy of metal atoms increases when immersed in the cooler water.
Chemistry 1 Final Exam Review
Page 9
49. This is a heating curve for a substance.
Between points X and Y, which would be observed?
A. Solid and liquid will be present.
C. Liquid and vapor will be present.
B. Only vapor will be present.
D. Only liquid will be present.
50. An open container of water is brought to a boil and heated until all of the water is converted to water vapor.
Which describes the changes in the water molecules?
A. The molecules speed up and move farther apart.
B. The molecules speed up and move closer together.
C. The molecules slow down and move farther apart.
D. The molecules slow down and move closer together.
51) 6.00 g of gold was heated from 20.0°C to 22.0°C. How much heat was applied to the gold?
A. 1.55 J
B. 15.5 J
C. 17.0 J
D. 32.5 J
52. A student has a beaker containing 55 g of water at 100°C. How much heat is needed to convert the water to steam?
A. 120,000 J
B. 18,000 J
C. 2,200 J
D. 330 J
53. How many grams of ice will melt at 0°C if the ice absorbs 420. J of energy?
A. 0.186 g
B. 0.795 g
C. 1.26 g
D. 5.38 × 104 g
54. An 18.0-g piece of an unidentified metal was heated from 21.5°C to 89.0°C. If 292 J of heat energy was absorbed by the
metal in the heating process, what was the identity of the metal? Use the specific heat table below.
A. calcium
B. copper
C. iron
D. silver
55. Consider this phrase diagram above. What process is occurring when a substance changes from point X (−130°C and 50 kPa)
to point Y (30°C and 100 kPa)?
A. boiling
B. freezing
C. melting
D. sublimation
56. When a chemical cold pack is activated, it becomes cool to the touch. What is happening in terms of energy?
A. An exothermic reaction is occurring, absorbing cold from its surroundings.
B. An exothermic reaction is occurring, releasing heat to its surroundings.
C. An endothermic reaction is occurring, releasing cold to its surroundings.
D. An endothermic reaction is occurring, absorbing heat from its surroundings.
57. What is the volume of two moles of hydrogen gas at STP?
A. 44.8 L
B. 22.4 L
C. 11.2 L
Chemistry 1 Final Exam Review
Page 10
D. 2.00 L
Chm.2.2 Analyze chemical reactions in terms of quantities, product formation, and energy.
Chm.2.2.1 Explain the energy content of a chemical reaction.
Chm.2.2.2 Analyze the evidence of chemical change.
Chm.2.2.3 Analyze the law of conservation of matter and how it applies to various types of chemical equations
(synthesis, decomposition, single replacement, double replacement, and combustion).
Chm.2.2.4 Analyze the stoichiometric relationships inherent in a chemical reaction.
Chm.2.2.5 Analyze quantitatively the composition of a substance (empirical formula, molecular formula, percent
composition, and hydrates).
Being lower in potential energy
makes reactants or products
more stable.
 Activated complex = top of
peak
 Activation energy = top of
peak minus reactants
 ΔH = products minus
reactants
 + ΔH is endothermic (heat
absorbed) and −ΔH is
exothermic (heat released).
A catalyst lowers the activation energy.
The Collision Theory
Signs a chemical reaction has occurred are
based on the following criteria:
 Precipitate formation
 Gas formation
 Color Change (not just a change in color
intensity as a result of dilution)
 Temperature change
Molecules must collide in order to react, and they
must collide in the correct or appropriate
orientation and with sufficient energy to equal or
exceed the activation energy.
TYPES OF REACTIO
Reactions fall into 5 categories:
1. Synthesis – only 1 product
2. Decomposition – only 1 reactant
3. Single replacement – one element switches
Places with another element in a compound
Product testing - A test for oxygen gas: Take a glowing splint and place it in a
sample of gas, if it re-ignites the gas is oxygen. A test for hydrogen gas: Use a
flaming wooden splint. H2 gas will make a squeaky pop when lit in air. Care
needs to be taken with this test as large amounts of hydrogen are very
explosive in air. A test for carbon dioxide gas: CO2 will extinguish a flame. A
second test for carbon dioxide gas: CO2 is bubbled through limewater, and
calcium carbonate solid is formed causing the limewater to become cloudy.
4. Double replacement – elements in 2 compounds
switch places
5. Combustion – products are carbon dioxide (CO2)
and water
In molecular equations, the formulas of the compounds are written as though
all species existed as molecules or whole units.
3NaOH + FeCl3 → Fe(OH)3 + 3NaCl
Chemistry
Final Exam
Review
An ionic1 equation
shows
dissolved ionic compounds in terms of their free
ions.
Page 11
3Na+ + 3OH− + Fe3+ + 3Cl−  Fe(OH)3 + 3Na+ + 3Cl−
The net ionic equation indicates only the species that actually take part in the
%X 
(mass X )(# X atoms)
% H 2O 
compound mass
x 100
(mass H 2O)(# H 2O' s )
compound mass
x 100
58. How many molecules are contained in 55.0 g of H2SO4?
A. 0.561 molecule
B. 3.93 molecules
C. 3.38 x 1023 molecules
D. 2.37 x 1024 molecules
59. How many moles are in 59.6 grams of BaSO4?
A. 0.256 mole
B. 3.91 moles
C. 13.9 moles
D. 59.6 moles
60. Analysis shows a compound to be, by mass, 43.8% N, 6.2% H, and 50.0% O. Which is a possible molecular formula for the
substance?
A. NH4NO2
B. NH4NO3
C. NH3OH
D. N2OH
61. If a sample of magnesium has a mass of 60. g, how many moles of magnesium does the sample contain?
A. 1.1 moles
B. 1.2 moles
C. 2.0 moles
D. 2.5 moles
62. What is the percent by mass of iron in the compound Fe2O3?
A. 70%
B. 56%
C. 48%
D. 30%
63. Metallic sodium reacts violently with water to form hydrogen and sodium hydroxide according to the balanced equation:
2Na + 2H2O→2NaOH + H2
How many moles of hydrogen gas are generated when 4.0 moles of sodium react with excess water?
A. 1.0 mole
B. 2.0 moles
C. 3.0 moles
D. 4.0 moles
64. A compound has an empirical formula of CH2O and a molecular mass of 180 g. What is the compound’s molecular formula?
A. C3H6O3
B. C6H12O6
C. C6H11O7
D. C12H22O11
65. Consider this reaction:
3Ca (s) 2H3PO4 (aq) →Ca3(PO4)2 (s) 3H2 (g)
How many moles of calcium are required to produce 60.0 g of calcium phosphate?
A. 0.145 mole
B. 0.194 mole
C. 0.387 mole
D. 0.581 mole
66. According to the equation 2H2O (l)→2H2 (g) + O2 (g), what mass of H2O is required to yield 22.4 L of O2 at STP?
A. 12 g
B. 18 g
C. 24 g
D. 36 g
Consider this reaction:
3Mg (s) + 2H3PO4 (aq) → Mg3(PO4)2 (s) + 3H2 (g)
67. How many grams of magnesium phosphate should be produced if 5.40 grams of magnesium react with excess phosphoric
acid?
A. 1.80 grams
B. 19.5 grams
C. 58.4 grams
D. 175 grams
68. Methane (CH4) is burned in oxygen according to this balanced chemical equation:
CH4 (g) + 2O2 (g) → CO2 (g) + 2H2O (g)
What volume of carbon dioxide is formed when 9.36 liters of methane are burned in excess oxygen at STP?
A. 9.36 L
B. 15.0 L
C. 18.7 L
D. 22.4 L
69. Consider this reaction: NH3 (g) HCl (g) →NH4Cl (s)
A. combustion
B. decomposition
Which type of reaction does this equation represent?
C. single replacement
D. synthesis
70. Which equation represents a single replacement reaction that can occur?
A. F2 2NaCl→2NaF Cl2
B. Cl2 2NaF → 2NaCl F2
C. Cu 2NaCl→ CuCl2 2Na
D. Zn 2NaF→ZnF2 2Na
71. What products are formed when the metal potassium is added to water?
A. K and H2O
B. KOH and H2O
C. K2O and H2
Chemistry 1 Final Exam Review
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D. KOH and H2
72. When Na2O reacts with H2O, what is produced?
A. HNaO2
B. Na H2O
C. NaO H2
73. Which equation is correctly balanced?
A. Cu + H2SO4 → CuSO4 + H2O + SO2
C. 2Fe + 3O2 → Fe2O3
B. 2Na + 2H2O →2NaOH + H2
D. 4Cu + S8 →8Cu2S
74. What coefficients are required to balance this equation?
_Fe2O3 _CO →_Fe _CO2
A. 2, 6, 3, 6
B. 1, 3, 2, 3
C. 1, 1, 2, 2
D. NaOH
D. 1, 1, 2, 1
75. Which example indicates that a chemical change has occurred?
A. When two aqueous solutions are mixed, a precipitate is formed.
B. As ammonium nitrate dissolves in water, it causes the temperature of the water to decrease.
C. Alcohol evaporates when left in an open container.
D. Water is added to blue copper (II) chloride solution. The resulting mixture is lighter blue in color.
76. An aqueous solution of silver nitrate is added to an aqueous solution of iron (II) chloride. Which is the net ionic equation for
the reaction that occurs?
A. AgNO2 (aq) FeCl (aq)→AgCl (s) FeNO2 (aq)
B. 2AgNO3 (aq) FeCl2 (aq)→ 2AgCl (s) Fe(NO3)2 (aq)
C. 2Ag+1 (aq) NO3-1 (aq) Fe+2 (aq) Cl2 (g)→2AgCl (s)
D. 2Ag+ (aq) 2Cl− (aq)→2AgCl (s)
77. This graph represents the change in energy for two laboratory trials of the same reaction.
Which factor could explain the energy difference between the trials?
A. Heat was added to trial #2.
B. A catalyst was added to trial #2.
C. Trial #1 was stirred.
D. Trial #1 was cooled.
Chm.3.1 Understand the factors affecting rate of reaction and chemical equilibrium.
Chm.3.1.1 Explain the factors that affect the rate of a reaction (temperature, concentration, particle size and presence
of a catalyst).
Chm.3.1.2 Explain the conditions of a system at equilibrium.
Chm.3.1.3 Infer the shift in equilibrium when a stress is applied to a chemical system (Le Chatelier’s Principle).
HOW TO INCREASE REACTION RATE
Le Châtelier’s Principle: if a system at
equilibrium is disturbed, the system will move in
1) Use more reactive reactants
such a way as to counteract the disturbance.
2) Increase concentrations of reacting particles
•
Adding a reactant or product shifts the
3) Increase the surface area of reactants
equilibrium away from the increase. Removing
4) Increase the temperature at which a reaction
a reactant or product shifts the equilibrium
occurs
towards the decrease.
5) Add a catalyst
•
As volume is decreased, pressure increases. An
6) Compress gas reactants
increase in pressure favors the direction that
has fewer moles of gas. In a reaction with the
same number of product and reactant moles of
EQUILIBRIUM CONSTANT
gas, pressure has no effect.
•
Adding heat (i.e. heating the vessel) favors away
Consider the Haber Process:
from the increase: If H > 0 (endothermic),
N2(g) + 3H2(g)
2NH3(g)
Chemistry 1 Final Exam Review
adding heat favors the forward reaction. If
We ignore the concentrations of pure liquids and pure solids in equilibrium
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H < 0 (exothermic), adding heat favors the
constant expressions (K). Pure liquids and solids have a fixed concentration
reverse reaction.
which will always be constant.
•
A catalyst will decrease the time taken to reach
equilibrium. A catalyst does not affect the
[ NH ]2
78. Consider this balanced chemical equation: 2H2O (aq) → 2H2O (l) + O2 (g) Which will increase the rate of the reaction?
A. increasing pressure on the reaction
B. decreasing concentration of the reactants
C. adding a catalyst to the reaction
D. decreasing the temperature of the reaction
79. For the reaction
A+ (aq) + B− (aq) → AB (s)
increasing the temperature increases the rate of the reaction. Which
is the best explanation for this happening?
A. The pressure increases, which in turn increases the production of products.
B. The concentration of reactants increases with an increase in temperature.
C. The average kinetic energy increases, so the likelihood of more effective collisions between ions increases.
D. Systems are more stable at high temperatures.
80. Which statement explains why the speed of some reactions is increased when the surface area of one or all the reactants is
increased?
A. increasing surface area changes the electronegativity of the reactant particles
B. increasing surface area changes the concentration of the reactant particles
C. increasing surface area changes the conductivity of reactant particles
D. increasing surface area enables more reactant particles to collide
81. Consider the equilibrium: CH3CO2H + CH3OH
CH3CO2CH3 + H2O
Which statement is correct?
A. The equilibrium shifts to the left-hand side if CH3CO2H is added to the system.
B. The equilibrium shifts to the right-hand side when extra CH3CO2CH3 is added.
C. Adding additional CH3OH results in the formation of more CH3CO2CH3.
D. The addition of CH3CO2H and CH3OH together has no effect on the position of the equilibrium.
82. Consider the equilibrium: Cr2O72–(aq) + 2OH–(aq)
2CrO42–(aq) + H2O(l) Which statement is correct?
2–
2–
Cr2O7 (aq) is orange, while CrO4 (aq) is yellow.
A. Adding aqueous NaOH shifts the equilibrium to the right-hand side.
B. The color of the aqueous solution is unaffected by the addition of acid.
C. Adding acid shifts the equilibrium to the right-hand side.
D. Adding acid has no effect on the position of the equilibrium.
83. For an irreversible chemical reaction, the rate will be affected by changes in all of these factors except:
A. temperature
B. presence of a catalyst
C. concentration of products D. surface area of a reactant
84. When a catalyst is added to a system at equilibrium, a decrease occurs in the
A. activation energy
B. heat of reaction
C. potential energy of the reactants
D. potential energy of the products
85. Which statement describes characteristics of an endothermic reaction?
A. The sign of H is positive, and the products have less potential energy than the reactants.
B. The sign of H is positive, and the products have more potential energy than the reactants.
C. The sign of H is negative, and the products have less potential energy than the reactants.
D. The sign of H is negative, and the products have more potential energy than the reactants.
86. In which equilibrium is the forward reaction favored by decreasing the external pressure?
A. 2SO2(g) + O2(g)
2SO3(g)
B. 2CO(g) + O2(g)
2CO2(g)
C. 2NO(g) + O2(g)
2NO2(g)
D. 2NH3(g)
N2(g) + 3H2(g)
87. Which conditions will increase the rate of chemical reaction?
A. decreased temperature and decreased concentration of reactants
B. decreased temperature and increased concentration of reactants
C. increased temperature and decreased concentration of reactants
D. increased temperature and increased concentration of reactants
88. Given the reaction at equilibrium: N2(g) + O2(g)
O2(g) will
A. decrease
B. increase
2NO(g) as the concentration of N2(g) increases, the concentration of
C. remains the same
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89. Which is a property of a reaction that has reached equilibrium?
A. The amount of products is greater than the amount of reactants.
B. The amount of products is equal to the amount of reactants.
C. The rate of the forward reaction is greater than the rate of the reverse reaction.
D. The rate of the forward reaction is equal to the rate of the reverse reaction.
90. Given the equilibrium system at 25°C:
NH4Cl(s)
What change will shift the equilibrium to the right?
A. decreasing the temperature to 15°C
C. increasing the temperature to 35°C
NH4+(aq) + Cl-(aq)
(ΔH = +3.5 kcal/mol)
B. dissolving NaCl crystals in the equilibrium mixture
D. dissolving NH4NONH3 crystals in the equilibrium mixture
Chm.3.2 Understand solutions and the solution process.
Chm.3.2.1 Classify substances using the hydronium and hydroxide concentrations.
Chm.3.2.2 Summarize the properties of acids and bases.
Chm.3.2.3 Infer the quantitative nature of a solution (molarity, dilution, and titration with a 1:1 molar ratio).
Chm.3.2.4 Summarize the properties of solutions.
Chm.3.2.5 Interpret solubility diagrams.
Chm.3.2.6 Explain the solution process.
Gas solubility decreases with increasing
temperature and decreasing pressure.
“Like dissolves like.”
Polar solvents dissolve
polar solutes.
Solutions are homogeneous mixtures
in a single phase.
ELECTROLYTES
Making Solutions
A salt dissolves faster if
 it is stirred or shaken,
 the particles are made smaller,
 the temperature is increased.
COLLIGATIVE
PROPERTIES
There are 4 colligative properties.
Adding a solute to a solvent does the
following:
Lowers Vapor Pressure
Elevates Boiling Point
Depresses Freezing Point
Increases Osmotic Pressure
Ultimately, colligative properties depend
only on the number of dissolved
particles in solution.
 On a line means saturated (a solution which holds the
maximum amount of solute per amount of the solution
under the given conditions).
 Below a line means unsaturated (the amount of
solute dissolved is less than the maximum that could
be dissolved).
 Above a line means supersaturated (contains more
solute than the usual maximum amount and is
unstable).
An electrolye has free ions
(such as Na+1 and Cl-1) in the
solution and is able to conduct
electricity and make a light bulb
burn brightly. When sugar
dissolves in water, there are no
free ions to conduct electricity.
The resulting solution is a
nonelectrolyte, so the light
bulb does NOT light up.
An Arrhenius acid is a substance that produces hydronium ions (H3O+) when
it dissolves in water. An Arrhenius base is a substance that produces
hydroxide ions, OH–, when it dissolves in water. A Brønsted-Lowry acid is a
proton (H+) donor. A Bronsted-Lowry base is a proton acceptor.
Properties of Acids
Acids taste sour. Lemon juice and vinegar, for example, are both aqueous solutions of acids. Acids conduct
electricity; they are electrolytes. Acids cause litmus paper to turns red and the indicator phenolphthalein to
turn clear and colorless. Acids react with active metals to form hydrogen gas. This property explains why
acids corrode most metals. Acids react with hydroxides (bases) to form water and a salt.
pH = - log [H+]
pH + pOH = 14.00
A compound is soluble due to the attraction between solvent and solute particles. Consider NaCl
in water: Na+ is attracted to O2− in water, while Cl− is attracted to H+ in water.
Chemistry 1 Final Exam Review
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Properties of Bases
Titration curve for the titration of
a strong base with a strong acid
Bases taste bitter and feel slippery. Bases can be strong or weak
electrolytes. They cause litmus paper to turn blue and the
indicator phenolphthalein to turn pink. Bases react with acids to
form water and a salt, but do not commonly react with metals.
mol
M
DILUTION
L
The number of moles of solute doesn’t change if you add more
solvent.
M1 x V1 = M2 x V2
.
M 
pH < 7 means acidic; H+ concentration is high
pH = 7 means neutral
pH > 7 means basic; H+ concentration is low
91. Which will increase the solubility of most solid solutes?
A. decreasing the temperature
C. increasing the temperature
92. Which 0.1 M solution will turn phenolphthalein pink?
A. NaOH
B. HCl
mol
L
Concentration is based on the molarity of a solution. A 12 M HCl
solution is said to be concentrated, while a 001 M HCl solution is said
to be dilute.
The strength of a base is based on the degree of dissociation. The
strength of a base does NOT depend on the molarity. A strong acid is
said to dissociate completely, while a weak acid only dissociates
(ionizes) slightly.
B. decreasing the amount of solvent at constant temperature
D. increasing the amount of solute at constant temperature
C. H2O
D. NaCl
93. A water sample was found to have a pH of 6 at 25°C. What is the hydroxide concentration in the water sample?
A. 1 × 10−8 M
B. 6 × 10−8 M
C. 1 × 10−6 M
D. 6 × 10−6 M
94. What is the pH of a solution of KOH with a hydroxide concentration of [OH−] = 1 × 10−4 M?
A. −10
B. −4
C. 4
D. 10
95. In a titration experiment, if 30.0 mL of an HCl solution reacts with 24.6 mL of a 0.50-M NaOH solution, what is the
concentration of the HCl solution?
A. 0.41 M
B. 0.61 M
C. 1.5 M
D. 370 M
96. A student wishes to prepare approximately 100 milliliters of an aqueous solution of 6 M HCl using 12 M HCl. How many
milliliters of 12 M HCl must be used?
A. 100 mL
B. 200 mL
C. 50 mL
D. 12 mL
97. Which substance can be classified as an Arrhenius acid?
A. HCl
B. NaCl
C. LiOH
D. KOH
98. Which solution will change red litmus to blue?
A. HCl
B. NaCl
C. CH3OH
D. NaOH
99. An acidic solution could have a pH of
A. 7
B. 10
C. 3
D. 14
100. What is the pH of a 0.00001 molar HCl solution?
A. 1
B. 9
C. 5
D. 4
101. As the hydrogen ion concentration of an aqueous solution increases, the hydroxide ion concentration of this solution will
A. increase
B. decrease
C. remain the same
102. What is the total number of moles of H2SO4 needed to prepare 5.0 liters of a 2.0 M solution of H 2SO4?
A. 2.5
B.5.0
C. 10
D. 20
Chemistry 1 Final Exam Review
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