Chemistry Review
Matter and Change
• Chm.1.1 – Analyze the structure of atoms,
isotopes, and ions.
• Chm.1.2 – Understand the bonding that occurs
in simple compounds in terms of bond type,
strength, and properties.
• Chm.1.3 – Understand the physical and
chemical properties of atoms based on their
position on the periodic table.
Chm.1.1.1 – Analyze the structure of
atoms, isotopes, and ions.
• Characterize protons, neutrons, electrons by
location, relative charge, relative mass (p=1, n=1,
e=1/2000).
• Use symbols: A = mass number, Z = atomic number
• Use notation for writing isotope symbols: 23592U or
U-235.
• Identify isotope using mass number and atomic
number and relate to number of protons, neutrons
and electrons.
• Differentiate average atomic mass of an element
from the actual isotopic mass and mass number of
specific isotopes.
Energetic particles that move in all directions
around the nucleus of an atom are called:
A. Neutrons.
B. Protons.
C. Elements.
D. Electrons.
• An ion with a net charge of +2 loses 5 electrons.
What is its charge?
A. 7+
B. 3+
C. 3D. 7-
• What is the net charge of an ion with 4 protons
and 6 electrons?
A. 24+
B. 10+
C. 2D. Neutral
What is the total number of subatomic particles in
an atom of potassium-39?
A. 19
B. 20
C. 39
D. 78
The number of protons in the nucleus of an atom
of a specific element is the same as that
element’s:
A. Atomic mass
B. Energy levels
C. Atomic number
D. Neutrons
• The number 80 in the name of bromine-80
represents
A. The atomic number
B. The mass number
C. The sum of protons and electrons
D. None of these
• The table gives the natural percent abundance of
the stable isotopes of sulfur. Based on the data,
what is the reasonable estimate for the average
atomic mass of sulfur?
Isotope
Natural %
Mass (amu)
A. 33.7
Abundance
B. 35.0
Sulfur-32
95.022
31.972
C. 32.1
Sulfur-33
0.76
32.971
D. 34.0
4.22
33.967
E. None Sulfur-34
Sulfur-36
0.014
35.967
Chm.1.1.2 – Analyze an atom in terms
of the location of electrons.
• Analyze diagrams related to the Bohr model of
the hydrogen atom in terms of allowed, discrete
energy levels in the emission spectrum.
• Describe the electron cloud of the atom in terms
of a probability model.
• Relate electron configurations of atoms to the
Bohr and electron cloud models.
• The quantum mechanical model of the atom:
A. Is concerned with the probability of finding an
electron in a certain position.
B. Was proposed by Neils Bohr.
C. Defines the exact path of an electron around
the nucleus.
D. Has many analogies in the visible world.
• What best describes how the Bohr model differs
from the quantum mechanical model of the atom?
A. The quantum model does not define an exact
pattern the electrons take around the nucleus.
B. The quantum model is based on fixed energy levels
of electrons.
C. The Bohr model restricts the energy of electrons to
certain values.
D. The Bohr model uses probability to determine the
location of finding an electron around the nucleus.
• Which best describes Bohr’s model of a
hydrogen atom?
A. The electron bound in a circular orbit around
the nucleus
B. A large, dense nucleus surrounded by atoms in
different orbits
C. A nucleus surrounded by electrons in specific
energy levels
D. All the particles mixed together like “plum
pudding”
Chm.1.1.3 – Explain the emission of
electromagnetic radiation in spectral form in
terms of the Bohr model.
• Understand that energy exists in discrete units
called quanta.
• Describe the concepts of excited and ground
state of electrons in the atom:
▫ Gaining energy results in the electron moving
from its ground state to a higher energy level.
▫ When the electron moves to a lower energy level,
it releases the energy difference in the two levels
as electromagnetic radiation (emissions
spectrum).
Chm.1.1.3 – Explain the emission of
electromagnetic radiation in spectral
form
in
terms
of
the
Bohr
model
• Understand that electromagnetic radiation is
given off as photons.
• Use the “Bohr Model for Hydrogen Atom” and
“Electromagnetic Spectrum” diagrams from the
Reference Tables to relate color, frequency, and
wavelength of the light emitted to the energy of
the photon.
• Understand the inverse relationship between
wavelength and frequency, and the direct
relationship between energy and frequency.
Chm.1.1.3 – Explain the emission of
electromagnetic radiation in spectral
form in terms of the Bohr model
• Explain that Niles Bohr produced a model of the
hydrogen atom based on experimental
observations. This model indicated that:
▫ An electron circles the nucleus only in fixed energy
ranges called orbit
▫ An electron can neither gain or lose energy inside
this orbit, but could move up or down to another
orbit
▫ That the lowest energy orbit is closest to the
nucleus
• The frequency and wavelength of all waves are
A. Directly related.
B. Inversely related.
C. Unrelated
D. Equal
• Once the electron in a hydrogen atom absorbs a
quantum of energy, it
A. Is now in its ground state
B. Is now in its excited state
C. Has released a photon
D. Follows an exact path around the nucleus
• Visible light is part of the electromagnetic
spectrum. If a light wave had a wavelength of
5.86 x 10-7 nm, what color would it appear?
A. Violet
B. Blue
C. Green
D. Yellow
• According to the Bohr Model for a hydrogen
atom, which change in energy level would emit
visible light?
A. n = 2 to n = 1
B. n = 2 to n = 3
C. n = 4 to n = 2
D. n = 4 to n = 6
• According to the Bohr Model for a hydrogen
atom, what wavelength of light would be emitted
if an electrons moved from n=3 energy level to
its ground state?
A. 103 nm
B. 434 nm
C. 656 nm
D. 1094 nm
FLAME TEST
Metal
Flame Color
Stronium
Scarlett
Calcium
Orange
Copper
Blue-Green
Potassium
Purple
Sodium
Yellow
A student performs a
flame test on an
unknown solution.
The unknown solution
causes the flame to
turn orange. Which
element is most likely
contained in the
solution?
A.
B.
C.
D.
Sodium
Copper
Potassium
Calcium
• According to the Bohr model for a hydrogen
atom, what wavelength of light would be emitted
when an electron jumps from n=2 to its ground
state?
A. 486 nm
B. 122 nm
C. 102 nm
D. 97 nm
Chm.1.1.4 – Explain the process of
radioactive decay using nuclear equations
and half-life.
• Use the symbols for and distinguish between
alpha and beta nuclear particles, and gamma
radiation , including relative mass.
• Use shorthand notation for particles involved in
nuclear equation to balance and solve for
unknowns.
• Compare the penetrating ability of alpha, beta,
and gamma radiation.
Chm.1.1.4 – Explain the process of
radioactive decay using nuclear
equations and half-life.
• Describe nuclear decay, including:
▫ Decay as a random event, independent of other
energy influences
▫ Using symbols to represent simple balanced decay
equations
▫ Simple half-life calculations
• Compare radioactive decay with fission and
fusion.
• Which type of ionizing radiation can be blocked
by clothing?
A. Alpha particle
B. Gamma radiation
C. X-rays
D. Beta particle
• If an isotope undergoes beta emission
A. The mass number changes.
B. The atomic number changes.
C. The atomic number remains the same.
D. The number of neutrons remains the same.
• Which of the following types of radiation has no
mass and no charge?
A. Alpha
B. Beta
C. Gamma
D. positron
• When Rn-222 undergoes decay to become Po218, it emits
A. An alpha particle
B. A beta particle
C. Gamma radiation
D. X-rays
• Which nuclear equation shows the radioactive
process for beta emission by argon-37?
• After 252 days, a 48-g sample of scandium-42
contains only 6.0-g of the isotope. What is the
half-life of scandium?
A. 84 days
B. 42 days
C. 32 days
D. 28 days
• When nuclear fission occurs
A. Two nuclei combine to produce a heavier
nucleus
B. The chain reaction that results cannot be
controlled
C. It is a spontaneous reaction
D. It must be initiated by bombardment with
neutrons
• Which of the following has a +2 charge?
A. Alpha particle
B. Beta particle
C. Neutrino
D. Gamma ray
Chm.1.2.1 – Compare (qualitatively) the
relative strengths of ionic, covalent, and
metallic
bonds.
• Describe metallic bonds :”metal ions plus ‘sea’ of mobile
•
•
•
•
•
electrons”.
Describe how ions are formed and which arrangements
are stable (filled sub-levels).
Appropriately use the term cation as a positively charged
ion and anion as negatively charged ion.
Predict ionic charges for representative elements based
on valence electrons.
Apply the concept that sharing electrons form a covalent
compound that is a stable arrangement.
Draw Lewis structure for simple compounds and
diatomic elements indicating single, double or triple
bonds.
• Why does a cation have a positive charge?
A. It has lost valence electrons.
B. It has gained valence electrons.
C. It has an ionic bond.
D. It has a metallic bond.
• How many electrons must an atom of strontium
lose to have to gain a noble gas electrons
configuration?
A. None
B. 1
C. 2
D. 3
• Which group of elements from the periodic table
gain one electron when they become ions?
A. Group 1
B. Group 2
C. Group 16
D. Group 17
• How many valence electrons do the elements in
Group 6A have?
A. 8
B. 6
C. 2
D. 0
• In general, metals react by
A. Losing valence electrons
B. Gaining valence electrons
C. Sharing valence electrons
D. Sometimes gaining and sometimes losing
valence electrons
• Which of the following is a diatomic molecule
containing a triple covalent bond?
A. N2
B. N2O
C. Br2
D. N2O2
Chm.1.2.2 – Infer the type of bond and
chemical formula formed between atoms.
• Determine that a bond is predominantly ionic by
the location of the atoms on the Periodic Table
(metals combined with nonmetals) or when
∆EN>1.7.
• Determine that a bond is predominantly
covalent by the location of the atoms on the
Periodic Table (nonmetals combined with
nonmetals) or when ∆EN<1.7.
• Predict chemical formulas of compounds using
Lewis structures.
• A covelent bond forms
A. When an element becomes a gas
B. When atoms share electrons
C. Between metals and nonmetals
D. When electrons are transferred from one atom
to another
• Which of the following compounds contain ionic
bonds?
A. CO2
B. NO
C. FeO
D. CH4
• Which of the following is a true statement?
A. When the electronegativelity difference between
two atoms is greater than 1.7, the atoms form a
covalent bond.
B. The size of th electronegativity difference between
two atoms has no bearing on the type of bond that
the two atoms may form.
C. As the electronegativity difference between two
atoms increases, the polarity of the bond
increases.
D. If the electronegativity difference between two
atoms is 0.3, the atoms have a weak polar bond.
Which pair of atoms is most likely to form an ionic
bond?
A. Mg and S
B. C and O
C. Br and S
D. Ca and O
Chm.1.2.3 – Compare inter- and intraparticle forces.
• Explain why intermolecular forces are weaker
than ionic, covalent or metallic bonds.
• Explain why hydrogen bonds are stronger than
dipole-dipole forces which are stronger than
dispersion forces.
Chm.1.2.3 – Compare inter- and intraparticle forces.
• Describe intermolecular forces for molecular
compounds.
▫ H-bond as attraction between molecules when H
is bonded to O, N, or F.
▫ Dipole-dipole attractions between polar
molecules.
▫ London dispersion forces (electrons of one
molecules attracted to nucleus of anotehr
molecule – i.e. liquefied inert gases
▫ Relative strengths – (H>dipole>London/van der
Waals)
• Which molecules has the strongest forces of
attraction between molecules?
A. HI
B. HBr
C. HF
D. HCl
Chm.1.2.4 – Interpret the name and
formula of compounds using IUPAC
convetion.
• Write binary compounds of two nonmetals (use
Greek prefixes)
• Write binary compounds of metal/nonmetal.
• Write ternary compounds (polyatomic ions) using
the polyatomic ions on the reference table.
• Write, with charges, these polyatomic ions: nitrate,
sulfate, carbonate, acetate, and ammonium.
• Know the names and formulas for these common
laboratory acids: HCl, HNO3, H2SO4, HC2H3O2 (or
CH3COOH)
Chm.1.2.5 – Compare the properties of
ionic, covalent, metallic, and network
compounds.
• Apply VSEPR for electron pair geometries.
• Describe bond polarity. Polar/nonpolar
molecules and solubility.
• What is the correct formula for sulfur trioxide?
A. S3O3
B. SO3
C. SO
D. S3O
• If you see an –ide at the end of a chemical name,
what can you assume?
A. It is a binary compound.
B. It is an acid.
C. It has a neutral charge.
D. It is a polyatomic ion.
• Which of the following represents the compound
formula from these two ions: silver and
hydroxide?
A. Ag2(OH)2
B. AgOH
C. Ag(OH)2
D. 2Ag(OH)2
• What is the formula for aluminum oxide?
A. AlO3
B. Al2O3
C. Al3O2
D. Al2O
• What is the name for the compound CuCl2?
A. Copper (I) chloride
B. Copper (II) chloride
C. Copper (I) chlorine
D. Copper (II) chlorine
• What is the formula for the compound FeCO3?
A. Iron carbon oxide
B. Iron (I) carbide
C. Iron (I) carbonate
D. Iron (II) carbonate
• What is the name for the compound N2O3?
A. Nitrogen oxide
B. Nitrous oxide
C. Nitrogen (II) oxide
D. Dinitrogen trioxide
Chm.1.2.5 – Compare the properties of
ionic, covalent, metallic, and network
compounds.
• Explain how ionic bonding in compounds
determines their characteristics: high MP, high
BP, brittle, high electrical conductivity either in
molten state or in aqueous solution.
• Explain how covalent bonding in compounds
determines their characterics: low MP, low BP,
poor electrical conductivity, polar nature.
• Explain how metallic bonding determines the
characteristics of metals: high MP, high BP, high
conductivity, malleability, ductillity, luster
• Using VSEPR theory, predict the geometry
around the central atom of methane, CH4.
A. Planar triangular
B. Tetrahedral
C. Linear
D. Bent
Test Results
Melting
Point
Sample A
Sample B
148
946
Electrical
None
conductivity
when
dissolved
Good
A student performs the
same test on two white
crysalline solids, A and
B. The two results are
shown below. Based on
the results which
statement is true?
A. Solid A contains only covalent
bonds and solid B contains
only ionic bonds.
B. Solid A contains only ionic
bonds and solid B contains
only covalent bonds.
C. Both solids contain only ionic
bonds.
D. Both solids contain only
covalent bonds.
• What is the shape of a molecule which has three
shared pairs of electrons and no unshared pairs?
A. Trigonal planar
B. Tetrahedral
C. Linear
D. Bent
• Which set of characteristics correctly describes
an ionic compound?
A. High melting point, high boiling point, brittle
crystals
B. Low melting point, low boiling point, soft
crystals
C. High melting point, low boiling point, soft
crystals
D. Low melting point, high boiling point, brittle
crystals
Chm.1.3.1 – Classify the components of a
periodic table (period, group, metal,
metalloid, nonmetal, transition)
• Identify groups as vertical columns on the
periodic table.
• Know that main group elements in the same
group have similar properties, the same number
of valence electrons, and the same oxidation
number.
• Summarize that reactivity increases as you go
down within a group for metals and decreases
for nonmetals.
Chm.1.3.1 – Classify the components of a
periodic table (period, group, metal,
metalloid, nonmetal, transition)
• Identify periods as horizontal rows on the periodic
table.
• Identify regions of the periodic table where metals,
nonmetals, and metalliods are located.
• Classify elements as metals/nonmetals.metallids
based on location.
• Identify representative (main group) elements as A
groups or as groups 1, 2, 13 – 18.
• Identify alkali metals, alkaline metals, halogens, and
noble gases based on location on the periodic table.
• Identify transition elements as B groups or as
groups 3 – 12.
• Which group of metals is so reactive that is
members are never found uncombined in
nature?
A. Alkali metals
B. Halogens
C. Alkaline earth metals
D. Noble gases
• In the periodic table, elements with similar
properties are found in the same
A. Group
B. Period
C. Row
D. Series
• Which of the following statements about a
column of the periodic table is true?
A. The elements have similar properties.
B. The elements have a wide range of properties.
C. The elements have the same atomic number.
D. The elements have the same atomic mass.