Fall 2021_CH1010_Dr. Kreider-Mueller
CH1010 Exam #1 Study Guide
For reference see “Chemistry: An Atoms-focused Approach” by Gilbert, Kirss, and Foster
Chapter 1: Matter and Energy, An Atomic Perspective
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Know the following SI base units of measure and their abbreviations (Table 1.2 in text):
*Mass (kg)
*Temperature (K)
*Time (s)
*Length (m)
*Amount of substance (mol)
You need to know (memorize) the following prefixes for multiples of SI Units
o Mega, kilo, deci, centi, milli, micro, and nano (Table 1.1 in your textbook).
o Be able to do the following:
 Know the abbreviations for each of the prefixes
 Know the conversion factors (ex. 1 g = 103 mg)
 Be able to perform these types of conversions in word problems
Know the equations for converting between Kelvin & Celsius (we will not provide you
with these equations on the exam!)
K = ˚C + 273.15˚
˚C = K – 273.15˚
Know what the difference is between SI base units and derived units and how to
perform conversions with derived units (ex: convert 3 m3 to 3 cm3)
o Know the following conversions for volume:
1 dm3 = 1 L
1 cm3 = 1 mL
o Know the equation for density (g/mL) (we will not provide you with this
equation!) and how to use it in a word problem. Practice solving for density given
a mass and volume. Practice solving for mass given a density and volume. Practice
solving for volume given a density and mass. Recognize how changes in mass
and/or volume affect density.
o Unit for energy is Joules (J). 1 J = (kg•m2)/s2
Know how to work with scientific notation. Given a number, be able to represent it in
scientific notation. Given a number in scientific notation, be able to represent the number
in its standard form. (See Appendix 1 in your textbook)
KNOW THE RULES OF SIGNIFICANT FIGURES:
1) All nonzero digits are significant
2) Zeros in the middle of a # are significant
 Ex: 2.9905 has 5 significant figures
3) Zeros at the beginning of a # are not significant
 Ex: 0.025 has 2 significant figures
4) Zeros at the end of a # and after a decimal point are significant
 Ex: 245.0 has 4 significant figures
5) Zeros at the end of a # and before a decimal point may or may not be significant
 Ex: 34,200 may have 3, 4, or 5 significant figures
 In this case it is best to use scientific notation to indicate the proper number
of significant figures
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Know the rules for significant figures when performing addition & subtraction, as well
as multiplication & division.
What is the difference between accuracy and precision in measurements?
Know how to round numbers:
o If the 1st digit you remove is less than 5 round down by dropping it and all
following digits.
o It the 1st digit you remove is 5 or greater, round up by adding 1 to digit on the left.
PRACTICE CONVERTING BETWEEN ONE UNIT AND ANOTHER. We have done
countless examples in class, on quizzes, in the lecture slides, on HMWK assignments etc.
You need to be comfortable with conversion problems. You may need to perform several
conversions in any one problem.
Be able to distinguish between solids, liquids, and gases at all 3 levels of representation:
symbolic (H2O), particulate (one molecule of H2O), and macroscopic (1 g of H2O)
o Be able to recognize graphical representations of molecules
o The properties of matter are related to its molecular level structure
Understand that chemistry is an experimental science and depends on accurate, precise,
and reproducible measurements
o Scientific Method: an approach to acquiring knowledge based on observation of
phenomena, development of a testable hypothesis, and additional experiments
that test the validity of the hypothesis.
o Hypothesis: a tentative and testable explanation for an observation or a series of
observations
o Scientific Theory: A general explanation of widely observed phenomena that has
been extensively tested.
o Scientific Law: a concise and generally applicable statement of a fundamental
scientific principle
Matter can be divided into two principle classes (Figure 1.2 in textbook):
o Pure substances: Matter that cannot be separated into simpler matter by a
physical process
 Element: a pure substance that cannot be separated into simpler
substances by an chemical process
 Atom: the smallest particle of an element that retains the chemical
characteristics of the element
 Compound: a pure substance that is composed of 2 or more elements
linked together in fixed proportions and that can be broken down into
those elements by some chemical process. (A sample of a given compound
may contain 1 molecule or several molecules of that compound.)
 Molecule: a collection of atoms chemically bonded together
o Mixtures: a combination of pure substances in variable proportions in which the
individual substances retain their chemical identities and can be separated from
one another by a physical process
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Homogeneous Mixture: a mixture in which the components are
distributed uniformly throughout and have no visible boundaries or regions
Heterogeneous Mixture: a mixture in which the components are not
distributed uniformly, so that the mixture contains distinct regions of
different compositions
Introduction to Energy
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Matter exists in 3 phases or physical states: solid, liquid gas
Matter can be transformed from one physical state to another as its temperature is raised
or lowered:
o solid  liquid (melting; endothermic)
o liquid  gas (vaporization; endothermic)
o solid  gas (sublimation; endothermic)
o gas  liquid (condensation; exothermic)
o liquid  solid (freezing; exothermic)
o gas  solid (deposition; exothermic)
Understand the difference between an exothermic and endothermic process
o Exothermic process: energy flows from a system into its surroundings
o Endothermic process: energy flows from the surroundings into the system
Be able to identify the system and the surroundings for phase changes. Where does the
energy (heat) flow? From the system to the surroundings? Or from the surroundings to
the system?
Understand the difference between heat and temperature:
o Heat: a flow of energy from one object or place to another due to differences in the
temperatures of the objects or places
o Thermal energy: the portion of the total internal energy of a system that is
proportional to its absolute temperature
o Temperature: a measure of thermal energy
What is kinetic energy?
KE = ½mv2
(where m = mass, v = velocity)
What is potential energy?
Why is thermal energy classified as kinetic energy rather than potential energy?
Chapter 2: Atom, Ions, and Molecules
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Know the following Laws:
o Law of Mass Conservation: Mass is neither created nor destroyed in chemical
reactions
o Law of Definite Proportions: Different samples of a pure chemical compound
always contain the same proportion of elements by mass
o Law of Multiple Proportions: Elements can combine in different ways to form
different chemical compounds, whose mass ratios are simple whole-number
multiples of each other.
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Know the various points in Dalton’s Atomic Theory:
1) Matter is composed of small particles called atoms.
2) Atoms of the same element are identical in shape and mass, but differ from the
atoms of other elements.
3) Atoms of an element cannot be changed into atoms of a different element by
chemical reactions. Atoms cannot be created or destroyed in chemical
reactions.
4) Atoms of different elements may combine with other atoms in fixed, simple,
whole number ratios to form compounds. A given compound always has the
same relative number and kind of atoms.
 Note: we now know that atoms can be further subdivided into protons,
neutrons, and electrons (but not by chemical processes). We also know that
atoms of some elements vary in their masses and densities (isotopes). Dalton’s
2nd postulate stated above is not correct!
What were the main conclusions from these experiments (What did we learn from
them? Section 2.1 in textbook)
o Thomson’s cathode ray experiment
o Millikan’s Oil drop experiment
o Rutherford’s Gold Foil experiment
Understand our current model of the atom:
o Nucleus: the positively charged center of an atom that contains nearly all the
atom’s mass
o Proton: a subatomic particle, present in the nucleus of an atom, that has a
relative charge of 1+ and a mass number of 1
o Neutron: An electrically neutral (uncharged) subatomic particle with a mass
number of 1
o Electron: a subatomic particle that has a relative charge of 1− and essentially
zero mass
Be able to relate the symbols and names for the elements in the first four rows of the
periodic table (elements hydrogen through krypton)
Dimitri Mendeleev published a table that is considered the forerunner of the modern
periodic table of elements.
The modern periodic table arranges the elements in order of their atomic numbers
What is a period of elements? All of the elements in a row of the periodic table.
What is a group or family of elements? All of the elements in a column of the periodic
table. Elements in a group have similar properties!
Know how to use the periodic table to look up the atomic number (Z) and Mass
Number (A) of different elements.
Know how to read/write a chemical symbol for a given element. Where to we put the
mass number? Where do we put the atomic number?
o Atomic number (Z): the number of protons in the nucleus of an atom
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o Mass Number (A): the number of nucleons in an atom
 Nucleon: a proton or neutron in a nucleus
Know how to calculate the mass number (A)
A = Number of protons (Z) + number of neutrons
Understand the difference between Mass Number and Average Atomic Mass.
o Average Atomic Mass: the weighted average of masses of all isotopes of an
element, calculated by multiplying the natural abundance of each isotope by its
mass in atomic mass units and then summing the products. The average atomic
mass is not a whole number!
o Isotope: atoms of an element containing the same number of protons but
different numbers of neutrons.
o Be able to identify the information given by the symbol for the isotope of an
element
o Be able to calculate atomic masses from relative abundances of isotopes
and vice versa.
Know where the following groups are located in the periodic table, and be familiar
with a few of their properties (recognize that there are multiple ways to label groups
in the periodic table):
o Transition Metals
o Inner Transition Metals
 Lanthanides
 Actinides
o Main Group Elements
 Main Group nonmetals
 Halogens (Group 7A or Group 17)
 Noble gases (Group 8A or Group 18)
 Main Group Metals
 Alkali Metals (Group 1A or Group 1)
 Alkaline-Earth Metals (Group 2A or Group 2)
How do metals, semimetals, and nonmetals differ? Where are they located in the
periodic table? What are the basic properties of each? Which ones are good
conductors of electricity?
Be able to recognize the difference between ionic solids and discrete molecules
o Ion: An atom or molecule that has a positive or negative charge
 Cation: A positively charged ion
 Anion: A negatively charge ion
o Ionic solid: a solid consisting of monatomic or polyatomic ions held together
by ionic bonds
o Be able to assign charges to the ion derived from MAIN GROUP elements
(Figure 2.10 in textbook)
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Interpret chemical formulas & calculate the formula mass of a compound.
Understand that in our macroscopic world, quantities of substances contain extremely
large numbers of atoms, ions, or molecules! We use the mole (the chemist’s dozen) to
express quantities of substances. The mole relates macroscopic quantities of
substances, such as their masses expressed in grams, to the # of particles they contain.
o Mole: An amount of a substance that contains Avogadro’s number of particles
(atoms, ions, molecules, or formula units)
o Avogadro’s number (NA): 6.0221 × 1023. The number of carbon atoms in
exactly 12 grams of the carbon-12 isotope
o Molar Mass: the mass of 1 mole of a substance.
o Be able to convert between the mass of a substance and the amount (# of
moles) and vice versa.
Chapter 3: Atomic Structure, Explaining the Properties of Elements
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What is radiant energy? What is the electromagnetic spectrum?
o Electromagnetic Radiation: any form of radiant energy in the electromagnetic
spectrum
Know the order of the different regions in the electromagnetic spectrum (Figure
3.1 in textbook). In order of increasing wavelength:
Gamma rays < X-rays < UV < Visible < IR < Microwaves < Radio waves
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Know the different parts of a wave: amplitude, peak, trough, wavelength, frequency
How are frequency and wavelength related?
c = 
c = speed of light (2.998 × 108 m/s)
= frequency (Hz or s−1)
 = wavelength (m)
Understand the development of the current model of the electronic structure of
atoms
o Understand what the Balmer-Rydberg equation tells us (YOU DO NOT NEED
TO USE/MEMORIZE THIS EQUATION FOR THE EXAM!)
o What is a continuous spectrum? What does it look like?
o What is an emission (line) spectrum? What does it look like?
 Atomic emission spectra: characteristic patterns of bright lines
produced when atoms are vaporized in high-temperature flames or
electrical discharges.
o What is an absorption spectrum? What does it look like?
 Atomic Absorption spectra: characteristic patterns of dark lines
produced when an external source of radiation passes through free,
gaseous atoms.
o Use the concepts of energy levels and orbitals to explain the occurrence of
emission and absorption spectra
o Understand the development of the dual wave/particle nature of the
electrons.
 The photoelectric effect: the release of electrons from a material as a
result of electromagnetic radiation striking it.
 Understand the experimental setup
 Be able to solve word problems involving the photoelectric effect
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o Work function (Φ): the amount of energy needed to
dislodge an electron from the surface of a material
o Threshold Frequency (0): the minimum frequency of
light required to produce the photoelectric effect
 Does the intensity of a beam of light affect the number of
electrons ejected from a metal? If so, how?
 What does it mean when we say energy is quantized?
 Quantized: having values restricted to whole-number multiples
of a specific base value
 Quantum: the smallest discrete quantity of a particular form of
energy
 Photon: a quantum of electromagnetic radiation
 Know De Broglie’s Hypothesis and how to use it in a word problem
 = h/(mv)
where h = Planck constant (6.626 × 10−34 J•s)
m = mass
v = velocity
 Know what Heisenberg’s Principle tells us (YOU DO NOT NEED TO
USE/MEMORIZE THIS EQUATION FOR THE EXAM!)
 The principle that one cannot simultaneously know the exact
position and the exact momentum of an electron
Know how to use Quantum Numbers:
o Quantum Number: one of four related numbers that specify the energy, shape,
and orientation of orbitals in an atom and the spin orientation of electrons in
the orbitals.
o What is an orbital?
 Regions in an atom where the probability of finding an electron is high
 Defined by the square of the wave function (ψ2)
o How do you determine possible values of n, l, ml, ms? (section 3.9 of textbook)
o What do each of the different quantum numbers tell us about an orbital?
o What letters correspond to each value of l?
o Know what a shell is, and which quantum number it refers to
 Shell: orbitals with the same value of n are in the same shell
o Know what a subshell is, and which quantum number it refers to
 Subshell: orbitals with the same values of n and l are in the same
subshell
Be able to identify an orbital (s, p, d) based on it’s shape and orientation on a 3D set of
axes
What is a node? How many nodes are in a p orbital? How many are in a d orbital?
o Node: a location in a standing wave that experiences no displacement
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Electron configurations:
o Know how to write and identify electron configurations using the standard
(expanded) form and shorthand versions.
o Be able to use/generate an orbital-filling diagram
o Know the following Principles & Rules:
 Pauli Exclusion Principle: No 2 electrons in an atom can have the same
values of their 4 quantum numbers
 Hund’s Rule: If 2 or more orbitals with the same energy are available,
one electron goes in each until all are half full. The electrons in the halffilled orbitals all have the same value of their spin quantum number.
 Aufbau Principle
 Lower energy orbitals fill before higher energy orbitals.
 An orbital can hold only 2 electrons, which must have opposite
spins (Pauli Exclusion principle).
 If 2 or more degenerate orbitals are available, 1 electron goes
into each until all are half-full (Hund’s Rule)
o Know that with increasing energy, the gap between one shell and the
next shell decreases. (Ex. there is a large energy gap between 1s and 2s
orbitals, but a smaller energy gap between 2s and 3s orbitals)
o Be able to explain anomalous electron configurations for Cr and Cu
 Cr: [Ar]3d54s1
 Cu: [Ar]3d104s1
o Recognize the relationship between an element’s electron configuration, the
number of valence electrons, and its Group Number. Elemental families have
similar electron configurations
 Core Electrons: Electrons in the filled, inner shells in an atom or ion
that are not involved in chemical reactions
 Valence Electrons: electrons in the outermost occupied shell of an
atom having the most influence on the atom’s chemical behavior
 Group 1A atom: [Noble Gas] ns1
 Group 2A atom: [Noble Gas] ns2
 Group 6A atom: [Noble Gas] ns2 np4
 Group 7A atom: [Noble Gas] ns2 np5
 How do electron configurations help us explain why metals tend to form
cations and nonmetals tend to form anions?
o Be able to write/identify electron configurations for ions
 s block elements: form monatomic cations by losing all the outer-shell
electrons leaving their ions with the electron configuration of the noble
gas immediately preceding them in the periodic table.
 Ex: Na = [Ne]3s1 Na+ = [Ne]
 p block elements: form monatomic anions by gaining enough electrons
to completely fill its valence-shell p orbitals. The ion formed has the
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electron configuration of the noble gas at the end of its row in the
periodic table
 Ex: Cl = [Ne]3s23p5
Cl− = [Ne]3s23p6 or [Ar]
 d block elements: the electrons in orbitals with the highest n value
ionize first . When forming cations, transition metals lose outer-shell s
electrons first, followed by d electrons
 Ex: Ni = [Ar]3d84s2
Ni2+ = [Ar]3d8
6 2
 Ex: Fe: [Ar]3d 4s
Fe2+: [Ar]3d6
Fe3+: [Ar]3d5
Trends to know (and be able to explain the trend…think about Zeff):
o Effective Nuclear Charge (Zeff): the attraction toward the nucleus experienced
by an electron in an atom; the positive charge on the nucleus reduced by the
extent to which other electron in the atom shield the electron from the nucleus
o Trend for atomic radius of neutral atoms (Figure 3.35 in textbook)
o Ionic radius (Figure 3.36 in textbook)
 When you form a cation, the radius shrinks (ex: the ionic radius of Na+ is
smaller than the atomic radius of Na). Why?
 When you form an anion, the radius expands (ex: the ionic radius of Cl−
is larger than the atomic radius of Cl). Why?
o Ionization Energy
 M + Energy  M+ + e−
M+ + Energy  M2+ + e−
M2+ + Energy  M3+ + e−
…and so forth
 Look at Table 3.2 in your book (Think about when/why there are large
jumps between successive ionization energies)
 Be able to make predictions about ionization energies for various
elements
o Electron Affinity
 Which group in the periodic table has the highest (most negative)
electron affinity? Which group(s) in the periodic table have the lowest
(values close to zero) electron affinity
 Look at Figure 3.38 in your book
 Be able to make predictions about electron affinities for various
elements
 Group 1A elements have low electron affinities so they tend to lose their
ns1 electron
 Group 2A elements have low electron affinities so they tend to lose their
ns2 electrons
 Group 7A elements have large electron affinities so they tend to gain one
electron, adopting the electron configuration of the neighboring noble
gas
 Group 8A elements are inert, and don’t generally gain or lose electrons
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