Honors Chemistry 1
Chapter 3 Study Guide
Converting between grams, moles, atoms, and molecules
WANT
GIVEN
Moles (mol)
Grams (g)
CONVERSION FACTOR TO
USE TO OBTAIN ANSWER
FINAL
ANSWER
UNIT
mols
𝑚𝑜𝑙
𝑔
Grams (g)
Moles (mol)
Atoms or
molecules
Moles (mol)
Moles (mol)
Atoms or
Molecules
Grams (g)
Atoms or
molecules
𝑔
𝑚𝑜𝑙
6.02 𝑥 10
Atoms or
molecules
grams
23
𝑎𝑡𝑜𝑚𝑠 𝑜𝑟 𝑚𝑜𝑙𝑒𝑐𝑢𝑙𝑒𝑠
1 𝑚𝑜𝑙
mols
1 𝑚𝑜𝑙
6.02 𝑥 1023 𝑎𝑡𝑜𝑚𝑠 𝑜𝑟 𝑚𝑜𝑙𝑒𝑐𝑢𝑙𝑒𝑠
grams
1 𝑚𝑜𝑙
6.02 𝑥 1023 𝑎𝑡𝑜𝑚𝑠 𝑜𝑟
𝑚𝑜𝑙𝑒𝑐𝑢𝑙𝑒𝑠
x
𝑔
𝑚𝑜𝑙
Grams (g)
𝑚𝑜𝑙
𝑔
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Atoms or
molecules
x
6.02 𝑥 1023 𝑎𝑡𝑜𝑚𝑠 𝑜𝑟 𝑚𝑜𝑙𝑒𝑐𝑢𝑙𝑒𝑠
Atoms or
molecules
𝑚𝑜𝑙
Page 1
Section 3.1: Substances Are Made Of ATOMS!
I.
Atomic Theory
a. The idea of an __________________________ is more than 2000 years old.
b. Until recently, scientists had never seen evidence of atoms.
c. The law of definite proportions, the law of conservation of mass and the law of multiple
proportions support the current atomic theory.
d. The figure on the right is a more accurate representation of an atom than the figure on the left
e. The Law of Definite Proportions
i. A chemical compound always contains the same elements in exactly the same
proportions by weight or mass. Additionally, every molecule of a substance is made of
the same number and types of atoms.
f. The Law of Conservation of Mass
i. The law of conservation of mass states that mass cannot be
_________________________________in ordinary chemical and physical changes.
ii. The mass of the reactants is equal to the ____________________________________.
1. This conservation law is similar to the “law of conservation of energy” that we
learned last chapter! Both of these laws are EXREMELY IMPORTANT TO
EVERY BRANCH OF SCIENCE!!
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II.
Dalton’s Atomic Theory
a. In 1808, John Dalton developed an atomic theory.
b. Dalton believed that a few kinds of atoms made up all matter.
c. According to Dalton, elements are composed of only one kind of atom and compounds are made
from two or more kinds of atoms.
d. Dalton’s Theory Contains Five Principles:
i. All matter is composed of extremely small particles called _____________, which cannot
be subdivided, created, or destroyed.
ii. Atoms of a given element are identical in their physical and chemical _______________.
iii. Atoms of different elements differ in their physical and chemical properties.
iv. Atoms of different elements combine in simple, whole-number ratios to form
compounds.
v. In _______________________________, atoms are combined, separated, or rearranged
but never created, destroyed, or changed.
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Section 3.2: Structure of Atoms
I.
Subatomic Particles
a. The smaller parts that make up atoms are called subatomic particles.
b. The three subatomic particles that are most important for chemistry are
_________________________________________
c. Please read your book section about Thomson’s discovery of electrons! It WILL be on your
quizzes and tests! See section 3.2!!
II.
Electrons have NEGATIVE CHARGE
a. Thomson’s experiments showed that a cathode ray consists of particles that have mass and a
negative charge.
b. These particles are called electrons.
c. An electron is a subatomic particle that has a negative electric charge.
d. Electrons are negatively charged, but atoms have ______________________.
i. Atoms contain some positive charges that balance the negative charges of the electrons.
III.
Rutherford discovered the NUCLEUS
a. Thomson proposed that the electrons of an atom were embedded in a positively charged ball of
matter. His model of an atom was named the _____________________________.
b. Please read your textbook to learn about Rutherford’s “Gold Foil Experiment” and his model of
the atom.
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c. The nucleus is the dense, central portion of the atom.
i. The nucleus is made up of ____________________________________.
ii. The nucleus has all of the positive charge, nearly all of the mass, but only a very small
fraction of the volume of the atom.
d. Protons and neutrons compose the nucleus
i. Protons are the subatomic particles that have a _____________________and that is
found in the nucleus of an atom.
1. The number of protons of the nucleus is the atomic number, which determines the
identity of an element.
2. Because protons and electrons have equal but opposite charges, a neutral atom
must contain equal numbers of protons and electrons.
ii. Neutrons are the subatomic particles that have ____________________ and that are
found in the nucleus of an atom.
e. Protons and neutrons can form a stable nucleus
i. Coulomb’s law states that the closer two charges are, the greater the force between them.
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ii. The repulsive force between two protons is large when two protons are close together.
iii. Protons form stable nuclei despite the ___________________ force between them.
iv. A strong attractive force between these protons overcomes the repulsive force at small
distances.
v. Because neutrons also add attractive forces, some neutrons can help stabilize a nucleus.
1. All atoms that have more than one proton also have neutrons.
f. Atomic Number is the number of__________________ in the nucleus
i. The number of protons that an atom has is known as the atom’s atomic number.
1. The atomic number is the same for all atoms of an element.
2. Because each element has a unique number of protons in its atoms, no two
elements have the same atomic number.
a. Example: the atomic number of hydrogen is 1 because the nucleus of each
hydrogen atom has one proton.
i. The atomic number of oxygen is _____.
ii. Atomic numbers are always whole numbers.
iii. The atomic number also reveals the number of electrons in an atom of an element.
1. For atoms to be neutral, the number of negatively charged electrons must
_____________ the number of positively charged protons.
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iv. The atomic number for oxygen tells you that the oxygen atom has 8 protons and 8
electrons.
g. Mass Number is the number of particles (protons + neutrons) in the nucleus
i. The ________________________ is the sum of the number of protons and neutrons in
the nucleus of an atom.
ii. You can calculate the number of neutrons in an atom by subtracting the atomic number
(the number of protons) from the mass number (the number of protons and neutrons).
iii. Unlike the atomic number, the mass number can vary among atoms of a single element.
1. Example: a particular atom of neon has a mass number of 20.
a. Because the atomic number for an atom of neon is 10, neon has 10
protons.
b. The neon atom has 10 protons, 10 electrons, and
10 neutrons. The mass number is_______.
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iv. Sample Problem A
1. How many protons, electrons, and neutrons are present in an atom of copper
whose atomic number is 29 and whose mass number is 64?
h. Atomic structure can be represented by symbols
i. Each element has a name, and the same name is given to all atoms of an element.
1. Example: sulfur is composed of sulfur atoms.
ii. Each element has a symbol, and the same symbol is used to represent one of the
element’s atoms.
iii. Atomic number and mass number are sometimes written with an element’s ___________.
iv. The atomic number always appears on the lower left side of the symbol.
v. Mass numbers are written on the upper left side of the symbol.
vi. Both numbers may be written with the symbol.
vii. An element may be represented by more than one notation.
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i. Isotopes of an element have the same ________________________
i. All atoms of an element have the same atomic number and the same number of protons.
Atoms do not necessarily have the same number of neutrons.
ii. Atoms of the same element that have different numbers of neutrons are called isotopes.
iii. One standard method of identifying isotopes is to write the mass number with a hyphen
after the name of an element.
iv. The second method of identifying isotopes shows the composition of a nucleus as the
isotope’s nuclear symbol.
v. All isotopes of an element have the same atomic number. However, their atomic masses
are _____________________________ because the number of neutrons of the atomic
nucleus of each isotope varies.
1. The two stable helium isotopes are helium-3 and helium-4.
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2. The Stable Isotopes of Lead
vi. Sample Problem B
1. Calculate the numbers of protons, electrons, and neutrons in oxygen-17 and in
oxygen-18.
Section 3.3: Electron Configuration
I.
Bohr’s Model Confines Electrons To “Energy Levels”
a. According to Niels Bohr’s model (the best atomic model we currently have), electrons can be
only certain distances from the nucleus. Each distance corresponds to a certain quantity of
energy that an electron can have.
i. An electron that is as close to the nucleus as it can be is in its
__________________________________
ii. The farther an electron is from the nucleus, the higher the ____________________ that
the electron occupies.
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b. The difference in energy between two energy levels is known as a ________________of energy.
II.
Electrons act like both particles and waves (true statement!)
a. Thomson’s experiments demonstrated that electrons act like particles that have mass.
b. In 1924, Louis de Broglie pointed out that the behavior of electrons according to Bohr’s model
was similar to the behavior of _______________.
c. De Broglie suggested that electrons could be considered waves confined to the space around a
nucleus.
i. As waves, electrons could have only certain frequencies which correspond to the specific
energy levels.
d. The present-day model of the atom takes into account both the ___________________________
properties of electrons.
e. In this model, electrons are located in orbitals, regions around a nucleus that correspond to
specific energy levels.
i. Orbitals are regions where electrons are likely to be found.
ii. Orbitals are sometimes called electron clouds because they do not have sharp, welldefined boundaries.
f. According to the current model of an atom, electrons are found in _________________.
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III.
Electrons and Light
a. By 1900, scientists knew that light could be thought of as moving waves that have given
frequencies, speeds, and wavelengths.
b. In empty space, light waves travel at 2.998  108 m/s.
c. The _________________________ is the distance between two consecutive peaks or troughs of
a wave.
d. The distance of a wavelength is usually measured in ________________.
i. THE ELECTROMAGNETIC SPECTRUM!!!
e. The electromagnetic spectrum is all of the frequencies or wavelengths of electromagnetic
radiation.
i. The wavelength of light can vary from 105 m to less than 10–13 m.
f. In 1905, Albert Einstein proposed that light also has some properties of ___________________.
g. His theory would explain a phenomenon known as the photoelectric effect.
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i. This effect happens when light strikes a metal and electrons are released.
h. Einstein proposed that light has the properties of both waves and particles.
i. Light can be described as a stream of particles, the energy of which is determined by the
light’s frequency.
ii. Light is an electromagnetic wave.
1. Red light has a ______________________ and a long wavelength.
2. Violet light has a ______________________ and a short wavelength.
3. The frequency and wavelength of a wave are inversely related.
a. The frequency and wavelength of a wave are inversely related.
i. As frequency increases, wavelength decreases.
1. High frequency = short wavelength = HIGH ENERGY
2. Low frequency = long wavelength = LOW ENERGY
i. Light emission
i. When a high-voltage current is passed through a tube of hydrogen gas at low pressure,
lavender-colored light is seen. When this light passes through a prism, you can see that
the light is made of only a few colors. This spectrum of a few colors is called a
_________________________________
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ii. Experiments with other gaseous elements show that each element has a line-emission
spectrum that is made of a different pattern of colors.
iii. In 1913, Bohr showed that hydrogen’s line-emission spectrum could be explained by
assuming that the hydrogen atom’s _______________ can be in any one of a number of
distinct energy levels.
1. An electron can move from a low energy level to a high energy level by absorbing
energy.
2. Electrons at a higher energy level are _______________ and can move to a lower
energy level by releasing energy. This energy is released as light that has a
specific wavelength.
3. Each different move from a particular energy level to a lower energy level will
release light of a different wavelength.
j. Light provides information about electrons
i. An electron in a state of its lowest possible energy, is in a ______________________.
1. The ground state is the lowest energy state of a quantized system
ii. If an electron gains energy, it moves to an excited state.
1. An excited state is a state in which an atom has more energy than it does at its
ground state
iii. An electron in an excited state will release a specific quantity of energy as it quickly
“falls” back to its ground state.
iv. An electron in a hydrogen atom can move between only certain ____________________,
shown as n = 1 to n = 7.
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v. In dropping from a higher energy state to a lower energy state, an electron emits a
characteristic wavelength of light.
IV.
Quantum Numbers
a. The present-day model of the atom is also known as the ____________________________.
i. According to this model, electrons within an energy level are located in orbitals, regions
of high probability for finding a particular electron.
b. To define the region in which electrons can be found, scientists have assigned four quantum
numbers that specify the properties of the electrons.
i. A quantum number is a number that specifies the ___________________ of electrons.
c. The principal quantum number, symbolized by n, indicates the main energy level occupied by
the electron.
i. Values of n are positive integers, such as ____________________________.
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ii. As n increases, the electron’s distance from the nucleus and the electron’s energy
increases.
d. The main energy levels can be divided into sublevels. These sublevels are represented by the
angular momentum quantum number, l.
i. This quantum number indicates the shape or type of orbital that corresponds to a
particular sublevel.
ii. A __________________________ is used for this quantum number.
1. l = 0 corresponds to an s orbital
2. l = 1 to a ___ orbital
3. l = 2 to a ___ orbital
4. l = 3 to an ___ orbital
e. The spin quantum number, indicates the orientation of an electron’s magnetic field relative to an
outside magnetic field.
i. The spin quantum number is represented by:
ii. A single orbital can hold a maximum of two electrons, which must have opposite spins.
f. Quantum numbers for the first 30 atomic orbitals
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V.
Electron Configurations
a. In 1925 the German chemist Wolfgang Pauli established a rule is known as the Pauli exclusion
principle.
b. The __________________________________ states that two particles of a certain class cannot
be in the exact same energy state.
i. This means that that no two electrons in the same atom can have the same ___________
quantum numbers.
c. Two electrons can have the same value of n by being in the same main energy level.
d. These two electrons can also have the same value of l by being in orbitals that have the same
shape.
e. These two electrons may also have the same value of m by being in the same orbital.
f. But these two electrons cannot have the same _______________quantum number.
g. If one electron has the value of 1/2, then the other electron must have the value of –1/2.
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h. The arrangement of electrons in an atom is usually shown by writing an electron configuration.
i. Like all systems in nature, electrons in atoms tend to assume arrangements that have the lowest
possible energies.
j. An _________________________________ of an atom shows the lowest-energy arrangement of
the electrons for the element.
k. A electron occupies the lowest energy level available
i. The aufbau principle states that electrons fill orbitals that have the lowest energy first.
1. Aufbau is the German word for “building up.”
ii. The smaller the principal quantum number, the ______________ the energy. Within an
energy level, the smaller the l quantum number, the lower the energy.
iii. So, the order in which the orbitals are filled matches the order of energies.
iv. The energy of the 3d orbitals is slightly higher than the energy of the 4s orbitals.
1. As a result, the order in which the orbitals are filled is as follows:
v. Additional irregularities occur at higher energy levels.
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vi. This diagrams shows how the energy of the orbitals can overlap.
l. An Electron Configuration is a Shorthand Notation
i. Based on the quantum model of the atom, the arrangement of the electrons around the
nucleus can be shown by the nucleus’s electron configuration.
1. Example: sulfur has sixteen electrons.
a. Its electron configuration is written as __________________________
2. Two electrons are in the 1s orbital, two electrons are in the 2s orbital, six
electrons are in the 2p orbitals, two electrons are in the 3s orbital, and four
electrons are in the 3p orbitals.
ii. Each element’s __________________________ builds on the previous elements’
configurations.
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1. To save space, one can write this configuration by using a configuration of a
noble gas.
a. neon, argon, krypton, and xenon
b. The neon atom’s configuration is ___________________, so the electron
configuration of sulfur is
iii. Electron orbitals are filled according to Hund’s Rule.
1. Hund’s rule states that orbitals of the same n and l quantum numbers are each
occupied by one electron before any pairing occurs.
2. Orbital diagram for sulfur
iv. Sample Problem C
1. Write the electron configuration for an atom whose atomic number is 20.
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Section 3.4: Counting Atoms
I.
Atomic Mass
a. You would not expect something as small as an atom to have much mass.
b. Example: copper atoms have an average mass of only 1.0552  10-25 kg.
i. Each penny has an average mass of 3.13  10-3 kg and contains copper.
ii. How many copper atoms are there in one penny?
iii. Assuming that a penny is pure copper, you can find the number of copper atoms using the
following conversion factor: __________________________________________
c. Masses of atoms are expressed in atomic mass units
i. A special mass unit is used to express atomic mass.
1. We will use the unit known as the “atomic mass unit,” or “amu”
ii. Mass number and atomic mass units would be the same because a proton and a neutron
each have a mass of about 1.0 amu.
1. Example: A copper-63 atom has an atomic mass of 62.940.
A copper-65 atom has an atomic mass of 64.928.
iii. In the periodic table, the mass shown is an average of the atomic masses of the naturally
occurring isotopes.
1. Example: copper is listed as 63.546 instead of 62.940 or 64.928.
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II.
Introduction to the mole
a. Most samples of elements have great numbers of atoms.
b. A mole is defined as the number of atoms in exactly 12 grams of carbon-12.The mole is the SI
unit for the amount of a substance.
c. The molar mass of an element is the mass in grams of one mole of the element. Molar mass has
the unit _________________________ (g/mol).
d. The mass in grams of 1 mol of an element is numerically equal to the element’s atomic mass
from the periodic table in atomic mass units.
e. Scientists have also determined the number of particles present in 1 mol of a substance, called
Avogadro’s number.
f. One mole of pure substance contains _____________________________ particles.
g. Avogadro’s number may be used to count any kind of particle, including atoms and molecules.
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III.
Lots of Example Problems
i. Sample Problem D: Determine the mass in grams of 3.50 mol of copper.
ii. Sample Problem E: Determine the mass in grams 1.74 mol of lead
iii. Sample Problem F: Determine the mass in grams 3.8 mol of magnesium
iv. Sample Problem G: Determine the number of mols in 85.92 grams of iron.
v. Sample Problem H: Determine the number of mols in 195.0 grams of gold.
vi. Sample Problem I: Determine the number of atoms in 195.0 grams of gold.
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vii. Sample Problem J: Determine the number of atoms in 0.30 mol of fluorine atoms.
viii. Sample Problem K: Determine the number of atoms in 0.1580 mol of phosphorus atoms.
ix. Sample Problem L: Determine the number of mols in 1.89 x 1024 lithium atoms.
x. Sample Problem M: Determine the mass, in grams, of 4.67 x 1022 atoms of titanium.
xi. Sample Problem N: Determine the mass, in grams, of 1.084 x 1025 atoms of sodium.
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IV.
Chemists and Physicists agree on a standard
a. Originally, atomic masses expressed the ratio of the mass of an atom to the mass of a
_____________________ atom.
i. Using hydrogen as the standard turned out to be inconvenient because hydrogen does not
react with many elements.
b. Because oxygen combines with almost all other elements, oxygen became the standard of
comparison.
i. This choice also led to difficulties because oxygen has as three isotopes.
c. In 1962, a conference of chemists and physicists agreed on a scale based on an isotope of carbon.
d. Used by all scientists today, this scale defines the atomic mass unit as exactly
______________________ of the mass of one carbon-12 atom.
i. As a result, one atomic mass unit is equal to 1.600 5402  10-27 kg.
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