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HOLT CHEMISTRY
CHAPTER 03
ATOMS AND MOLES
I.
Substances Are Made of Atoms
Objectives:
 State the three laws that support the existence of atoms.
 List the five principles of John Dalton’s atomic theory.
A. Atomic Theory
1. Scholarly thinkers who considered the mysteries of life were known as philosophers (lovers of
wisdom).
2. Democritus (460-370 B.C.)
a)
Matter is composed of empty space through which atoms move.
b)
Atoms are solid, homogeneous, indestructible, and indivisible.
c)
Different kinds of atoms have different sizes and shapes.
d)
Apparent changes in matter result from changes in the grouping of atoms and not from
changes in the atoms themselves.
3. Aristotle (384-322 B.C.)
a)
One of the most influential philosophers.
b)
Wrote extensively on many subjects, including politics, ethics, nature, physics, and
astronomy.
c)
Most of his writings have been lost through the ages.
d)
He rejected the atomic theory and won the approval of other philosophers.
4. Aristotle concluded that matter was composed of a combination of Earth, water, air, and fire.
5. Philosophers were not scientists because they could not/did not perform controlled experiments to
substantiate their ideas.
6. The Law of Definite Proportions states that two samples of the same compound are made of the
same elements in exactly the same proportions by mass regardless of the size of the sample.
7. The Law of Conservation of Mass states that the mass of the in a reaction equals the mass of the
products.
8. The law of Multiple Proportions states that if two or more compounds are composed of the same
two elements, the ratio of the masses of the ratio of the second element is in different, small, whole
number than that of the first element.
B. Dalton’s Atomic Theory Contains Five Principles
1. John Dalton (1766-1844)
a)
Was an English school teacher.
b)
Conducted scientific research on Democritus’ ideas.
2. The main points of Dalton’s Atomic Theory are:
a)
1. All matter is composed of extremely small particles called atoms.
b)
2. All atoms of a given element are identical, having the same size, mass, and chemical
properties. Atoms of a specific element are different from those of any other element.
c)
3. Atoms cannot be created, divided into smaller particles or destroyed.
d)
4. Different atoms combine in simple whole number ratios to form compounds.
e)
5. In a chemical reaction, atoms are separated, combined, or rearranged.
3. Dalton’s atomic theory explains the conservation of mass in a chemical reaction.
II.
Structure of Atoms
Objectives:
 Describe the evidence for the existence of electrons, protons, and neutrons, and describe the properties of
these subatomic particles.
 Discuss atoms of different elements in terms of their numbers of atomic number, and mass number.
 Define isotope, and determine the number of particles in the nucleus of an isotope.
A. Subatomic Particles
1. Experiments, by several scientists, in the mid 1800’s, brought the first challenges to Dalton’s
atomic theory when they revealed that atoms are made of several subatomic particles.
a)
Since that time, many types of subatomic particles have been discovered (fermions, leptons,
quarks, bosons, baryons, hadrons, mesons)
b)
The three particles that are most important for chemistry are the proton, neutron, and electron.
2. Electrons Were Discovered by Using Cathode Rays
a)
Like many other life changing discoveries, the electron was discovered accidentally.
b)
A cathode ray tube is an evacuated glass tube with electrodes on each end.
c)
The negative electrode is called the cathode and the positive electrode is called the anode.
d)
Because the radiation of energetic particles now known as electrons originates at the negative
terminal, this tube is called a cathode ray tube.
e)
The English physicist J.J. Thomason (1856-1940) is credited with discovering the first
subatomic particle, the electron.
f)
In 1909, the American physicist Robert Millikan (1868-1953) determined the charge and mass
of the electron in the classic oil droplet experiment.
g)
The first model of an atom was the short lived plum pudding model of the atom proposed by
J.J. Thomason.
3. Rutherford Discovered the Nucleus
a)
In 1911, James Rutherford (1871-1937 performed his famous “gold foil” experiment. (Board)
b)
Rutherford concluded that atoms are mostly empty space with a tiny, very dense region in
their center called a nucleus.
c)
The nucleus of an atom is so dense that if one were the size of the dot in an exclamation point
at the end of a sentence, it would have a mass equal to 70 automobiles.
d)
An atom’s diameter is approximately 10,000 times that of its nucleus.
e)
If an atom had the diameter of two football fields, its nucleus would be the size of a nickel.
4. Protons and Neutrons Compose the Nucleus
a)
The positively charged particles in an atom’s nucleus are called protons.
b)
A proton has a charge of +1 (exactly equal to and opposite from an electron).
c)
A proton has a mass 2000 times greater than an electron.
d)
Because protons and electrons have equal but opposite charges, an atom must have equal
numbers of protons and electrons.
e)
Neutrons are subatomic particles found in the nucleus that are equal in mass to the proton but
have no electrical charge.
5. Protons and Neutrons Can Form a Stable Nucleus
a)
Coulomb’s Law states that as the distance between two charges decreases, the force of
repulsion between them increases.
b)
There exists a fundamental force that attracts all subatomic particles called the strong nuclear
force.
c)
The addition of some neutrons to the nucleus makes it possible for the strong nuclear force to
overcome the force of repulsion due to like charges.
d)
All atoms with more than one proton also have neutrons.
B. Atomic Number and Mass Number
1. The atomic number is the number of protons in an atom’s nucleus.
a)
Since all atoms are electrically neutral, the atomic number also discloses the number of
electrons an atom has.
b)
Atomic numbers are always whole numbers.
c)
To date, scientists have discovered 113 elements ranging in atomic number from 1-114.
2. The mass number of an atom is the total number of particles (protons and neutrons) in the nucleus.
a)
To determine the number of neutrons in an atom, subtract the atomic number from the mass
number.
3. Because of isotopes, it is possible for atoms of different elements to have the same mass number.
4. Atomic structure can be represented by symbols.
a)
The atomic number (number of protons) is written at the lower, left side of the element’s
symbol.
b)
1H 2He 3Li 4Be 5B
c)
The mass number (number of protons and neutrons) is written at the upper, left side of the
element’s symbol.
1
d)
H 2H 3He 4He 6Li 7Li 9Be 10B 11B
e)
Both notations may be used at the same time
1
f)
𝐻 42𝐻𝑒 73𝐿𝑖 94𝐵𝑒 115 𝐵𝑒
1
5. Isotopes of an element have the same atomic number.
a)
Although all atoms of a given element have the same number of protons, they can have
different numbers of neutrons.
b)
Atoms of an element with different numbers of neutrons are called isotopes.
III.
Electron Configuration
Objectives:
 Compare the Rutherford, Bohr, and quantum models of an atom.
 Explain how the wavelengths of light emitted by an atom provide information about electron energy levels.
 List the four quantum numbers and describe their significance.
 Write the electron configuration of an atom by using the Pauli exclusion principle and the aufbau principle.
A. Atomic Models
1. Because the atom is far too small to see, it is necessary to construct models of it for study and
explanation.
2. As we have read, J.J. Thomas’ plum pudding model was the first model of the atom.
3. Earnest Rutherford’s Model
a)
Replaced the plum pudding model with the nuclear model
b)
Provided for the electrons to orbit around the nucleus in undefined, circular, or elliptical
orbits.
4. Niels Bohr’s Model
a)
Replaced Rutherford’s model two years later.
b)
Provided for different energy levels for the electrons to travel in.
c)
Defined the difference in the amount of energy from one level to the next as a quantum of
energy.
i.
Compared energy levels to rungs on a ladder.
ii.
States that an electron can only be in a definite energy level and not in between energy
levels.
5.
Electrons Act Like Both Particles and Waves
a)
By using an electroscope, Thomas’s experiments demonstrated that electrons have mass (they
could turn the paddle wheel).
b)
Louis de Broglie also demonstrated by using Bohr’s model, that electrons have characteristics
like waves because, when confined in space, light can have only discrete wavelengths.
c)
The present Quantum Model of the atom takes into consideration both the particle and wave
properties of electrons.
i.
Electrons are found in orbitals (specifically shaped paths) in definitive regions that
correspond to energy levels.
ii.
Like the blades on a moving fan, the orbital shows where the electron is most likely to
be found at any given time.
B. Quantum Numbers
1. Scientists have assigned four quantum numbers to each electron to define the region where that
electron is most likely to be found.
a)
The principal quantum number (n number) indicates the energy level and can be any integer
(1, 2, 3, 4, ……)
b)
The angular momentum number (l number) indicates the shape of the orbital (sublevel) and
can have a value ranging from 0-n.
i.
s-sublevels are designated 0
ii.
p-sublevels are designated 1
iii.
d-sublevels are designated 2
iv.
f-sublevels are designated 3
c)
The magnetic quantum number (m number) indicates the numbers and orientations of the
orbitals and range from l-1 to l+1.
d)
The spin quantum number symbolized by +½ or -½ or (↑↓) represents the spin on a particular
electron and is important because a single orbital can only contain two electrons and they
must have opposite spins.
C. Electron Configurations
1. The Pauli exclusion principle (1925 Wolfgang Pauli) states that no more than two electrons may
occupy the same orbital.
2. The aufbau principal (German for “building up”) states that electrons fill orbitals that have the
lowest energy first.
3. Hund’s rule states that orbitals of the same n and l quantum numbers are each occupied by one
electron before any pairing occurs.
4. The arrangement of electrons in an atom is usually shown by writing an electron configuration.
(numerical representation)
5. The arrangement of electrons for an atom may also be drawn in and orbital diagram.
6. A Lewis diagram shows the valence electrons for an atom by drawing them around the atom’s
symbol.
IV.
Counting Atoms
Objectives:
 Compare the quantities and units for atomic mass with those for molar mass.
 Define mole, and explain why this unit is used to count atoms.
 Calculate either mass with molar mass or number with Avogadro’s number given an amount in moles.
A. Atomic Mass
1. Because an atom is so small, a special mass unit is used to express atomic mass.
a)
The atomic mass unit (amu). [used in this book}
b)
The Dalton (Da).
2. One atomic mass unit is equal to the mass of one proton or one neutron.
B. Introduction to the Mole
1. A mole is the number of atoms in exactly 12 grams of carbon-12.
2. The mole is simply a counting unit.
3. The molar mass of an element is the mass of exactly one mole of the element’s atoms.
4. The number of particles in a mole is called Avogadro’s number and is equal to 6.022 x 1023.
a)
[602200000000000000000000]
b)
602.2 sextillion