Chapter 8
Atoms and Periodic Properties
The Atom
• The concept that matter is made up of
atoms dates back to the Greeks 2,500
years ago.
• Democritus in 460-362 BC came up with
the atomic model in which the atom was
indivisible.
• In the early 1800’s John Dalton
reintroduced the concept of the atom.
Dalton’s Atomic Theory
• Indivisible particles called atoms make up all
matter.
• All the atoms of an element are exactly alike
in shape and mass.
• The atoms of different elements differ from one
another in their masses.
• Atoms chemically combine in definite whole
number ratios to form chemical compounds.
• Atoms are neither created or destroyed in
chemical reactions.
• Bolded items were later found to be untrue
Discovery of the Electron
• Thomson discovered the existence of the
electron in 1897.
• He placed charged metal plates on each side of
a beam from a high voltage electrical source in a
vacuum and found that the beam was deflected
away from the negative plate.
• Since like charges repel, this meant that the
beam was made up of negatively charged
particles.
The Nucleus
• Ernest Rutherford discovered the nucleus in
1907.
• Alpha particles from a radioactive source were
allowed to move as a narrow beam of fast
moving particles through a small opening in a
lead container directed to a very thin gold foil.
• The alpha particles were detected by plates that
produced small flashes of light when struck by
alpha particles.
• Most of the alpha particles went right through the
foil. Some of the particles were deflected at
different angles, some even backwards.
Conclusions of Rutherford’s
Experiment
• The positive alpha particles that were
repelled was because they encountered
massive positive charge concentrated in a
small region of the gold atom.
• He concluded that the atom contains a
tiny, massive and positively charged
nucleus surrounded by electrons.
• The volume of the atom is mostly empty
space.
The Nucleus
• The nucleus contains protons and neutrons.
• The protons are positively charged while the
neutrons do not have a charge, they are neutral.
• The number of protons is equal to the atomic
number of the element.
• The number of protons is equal to the number of
electrons, since all atoms are neutral.
• All atoms of the same element have the same
atomic number, and hence number of protons.
• Every element has a different atomic number.
• There are 113 different elements.
Isotopes
• The neutrons in the nucleus, along with the
protons, contribute to the mass of the atom.
• The number of neutrons in an atom of a given
element may vary.
• Atoms of the same element which have different
number of neutrons are called isotopes.
• Isotopes have the same number of protons and
electrons but different number of neutrons.
Atomic Mass Unit
• The unit of mass used for atoms is the atomic mass unit
(AMU)
• This is defined as 1/12 of the mass of C-12, the
isotope of carbon which contains 6 protons and 6
neutrons.
• The mass number of atoms is equal to the number of
protons plus the number of neutrons expressed in AMU.
• The atomic weight or atomic mass of an atom is the
weighed average of all the mass numbers of the
isotopes of that element. Both the mass number and the
% abundance of each isotope of the element are taken
into account to arrive at the atomic weight. For example,
the element chlorine has a Cl-35 (75%) and a Cl-37
(25%) isotope. The atomic mass found in the periodic
table for chlorine is 35.45.
Protons, Neutrons, Electrons
Particle
proton
Symbol
p+
Charge
+1
Mass
1 amu
neutron
no
0
1 amu
electron
e-
-1
0 amu
The number of protons = number of electrons.
The number of neutrons varies. The atomic number is the
number of protons, the mass number is the number of
protons plus neutrons.
Bohr Model of the Atom
Electrons
(Total of 7)
14
7
N
Atomic weight =
14.0067 amu
Protons
(total of 7)
Neutrons
(total of 7)
Bohr Theory of the Atom
• An electron can revolve around an atom only in specific
allowed orbits.
• An electron in an orbit does not emit radiant energy as
long as it remains in the orbit.
• An electron gains or loses energy only by moving from
one allowed orbit to another. The energy an electron has
depends on which orbit it occupies The only way that an
electron can change its energy is to jump from one orbit
to another in “quantum leaps”.
• An electron must acquire energy to jump from a lower
orbit to a higher one, and the acquisition of energy must
be sufficient for the jump to occur all at once.
• The energy that the electron acquires comes from high
temperatures or from electrical discharges.
• An electron jumping from a higher to a lower orbit gives
up energy in the form of light. A single photon is emitted
when a downward jump occurs and the energy of the
photon is exactly equal to the difference in the energy
level of the two orbits. This is a quantum of energy.
A photon possesses a quantum
of energy. The frequency is
proportional to the amount of
energy the photon possesses.
Bohr Orbits
n=4
n=3
n=2
n=1
Nucleus
The further away from
The nucleus the higher the
frequency and
energy of the electron.
Quantum Mechanical Model of the
Atom
• This is highly mathematical treatment of where the
electrons are located in the atom.
• The Bohr model of the atom considered electrons as
particles in circular orbits that could be only certain
distances from the nucleus.
• The quantum mechanical model considers the electron
as a wave and considers the energy of its harmonics.
• In the Bohr model the location of the electron was
certain, in an orbit, in the quantum mechanical model the
electron is a spread out wave and its location is
uncertain.
• The Bohr model dealt with orbits, the quantum
mechanical model deals with energy levels.
Quantum Mechanical Model of the
Atom
• Quantum mechanics indicates the
probability of finding an electron in a
certain region in space at a given instant.
• There are shapes associated with the
different energy levels where electrons
may be.
90% of the time an electron in a particular energy level
will be found in a particular region in space which has specific
boundaries. Increasing energy levels have the same shape but
occupy larger areas which reach further away from the nucleus.
The Periodic Table
• All the elements are located in the periodic table.
• The rows are called periods and are horizontal.
There are rows 1-7
• The columns are called groups or families and
are vertical. There are groups 1-8 A and the B
groups.
• The table is arranged with increasing atomic
numbers going left to right and coming around to
the next row.
• Elements in the same group have similar
chemical properties.
Periodic Table
• Elements are identified with a symbol consisting of 1 or 2
letters. The first one is always capitalized, the second
one, if present, is lower case.
• The symbols may come from latin or other language.
• The number above the symbol is the atomic number of
the element.
• The number below the symbol is the atomic weight of the
element.
• The elements to the left of the staircase are metals and
to the right are nonmetals.
• The majority of the elements are metals and solids.
Groups in the Periodic Table
• Group IA is called the alkali metals with the exception of
hydrogen. These are the most reactive metals. They
don’t occur by themselves in nature, only as compounds.
They are called the alkali metals because they react
violently with water to form alkaline or basic solutions.
Sodium is an example of this. Group IA are all solids.
• Group IIA is called the alkaline earth metals and they are
the second most reactive group of metals. Magnesium is
an example of this. Group IIA are all solids.
• Group VIIA is called the halogens. These are the most
reactive nonmetals. Fluorine and chlorine are greenish
poisonous gases. Bromine is a reddish-brown liquid and
iodine is a dark purple solid. They are used as
disinfectants and have other uses as well. They react
with metals to form salts. e.g. NaCl, sodium chloride,
which is table salt.
Groups in the Periodic Table
• Group VIIIA is the noble gases. These are the
most stable of all elements, not reactive at all,
and they are all gases.
• Groups IA-VIIIA are the representative elements
and the B groups are the transition elements, or
transition metals. The transition metals are all
solids except for mercury, which is a liquid. Many
of these metals are very common, such as
copper, silver, and gold, iron. Transition metals
are usually not very reactive and can sometimes
be found in nature in mines as elements.
• The two rows of elements along the bottom of
the periodic table are called the inner transition
elements. The first row is the lanthanides or rare
earths. The second row is the actinides. All
elements after uranium (U) are manmade.
Metals, nonmetals, and Metalloids
• Metals are shiny, all are solids except for mercury and all
conduct electricity. They can be drawn into wires (ductile)
and can be bent (malleable). About 80% of the elements
are metals.
• Nonmetals are dull, there are solids and gases and one
liquid, bromine. They do not conduct electricity and the
solids are brittle, they break rather than bend. Consider a
piece of charcoal, made up of C atoms.
• Metalloids are the elements on either side of the
staircase. They are semiconductors and they are all
solids. They have properties intermediate between
metals and nonmetals. They are used to make computer
chips, e.g., silicon.
Electrons in Atoms
• n=1 (energy level 1) can have a maximum
of 2 electrons.
• n=2 (energy level 2) can have a maximum
of 8 electrons.
• n=3 (energy level 3) can have a maximum
of 18 electrons
• n=4 (energy level 4) can have a maximum
of 32 electrons.
The Periodic Table
• Elements in the same group have the same
number of electrons in the last energy level.
These are called valence electrons.
• The number of valence electrons for the Group A
elements is equal to the group number. e.g. C is
in group IVA and has 4 valence electrons, Br is
in group VIIA and has 7 valence electrons.
• The valence electrons determine the chemical
properties and reactivity of the elements, and
also some of the physical properties.
Electrons in Atoms
Period 1: H (1 e-) and He (2e-) n=1 only
Period 2: n=1(2e-) and n=2 (1-8 e-)
Period 3: n=1 (2e-), n=2 (8e-), n=3 (1-8e-)
IA
VIIIA
IIA
IIIA IVA VA VIA VIIA
Electrons in Atoms
Period 1: H (1 e-) and He (2e-) n=1 only
Period 2: n=1(2e-) and n=2 (1-8 e-)
Period 3: n=1 (2e-), n=2 (8e-), n=3 (1-8e-)
All elements want 8e- in last energy
Level, like the noble gases (group VIII)
Electron dot structures
• The group number of the representative elements
determines the number of electrons in the valence shell
(valence electrons). These are the electrons in the
highest energy level of the atom.
• Valence electrons can be represented using dot
structures.
• Atoms of representative elements form ions by losing or
gaining electrons so that they will have 8 valence
electrons, like the stable noble gases.
• Groups IA-IIIA will lose 1, 2, or 3 electrons so that they
will go to the previous energy level and have 8 electrons
there.
• Groups V-VIIA will gain 3, 2, or 1 electrons so that they
will end up with 8 valence electrons.
Ions
• Ions are atoms that have gained or lost electrons so that
the total number of electrons are no longer the same.
The number of protons is still the same.
• Every time an electron is lost the charge becomes +1.
Only metals will lose electrons. Every time an electron is
gained the charge becomes -1. Only nonmetals will gain
electrons. If more than one is lost or gained it will be an
additional +1 or -1 for each.
• The elements in groups IA-IIIA form +1, +2, and +3 ions
respectively. These are all metals. The group number
gives you the positive charge of the ion formed.
• The elements in groups V-VIIA will form -3, -2, and -1
ions respectively. These are all nonmetals.
• 8-group number gives you the negative charge of the
ion formed.
• The elements in group IVA will gain electrons if they are
nonmetals and lose if they are metals. The metalloids do
not tend to gain or lose electrons.
Ions formed by Representative
Elements
Element
Na
Mg
Al
N
O
Cl
Ne
Group Valence eIA
1
IIA
2
IIIA
3
VA
5
VIA
6
VIIA
7
VIIIA
8
Ion Formed
Na+
Mg 2+
Al 3+
N 3O 2Cl None
Element Names to Know
• Groups IA, IIA, VIIA, and VIIIA
• Periods 1, 2, 3, 4.
• Groups IB and IIB
Practice Exercises
• p. 226-228 Applying Concepts. # 2, 3, 4, 5, 6, 8,
9, 10, 13, 14, 15, 23, 24, 25, 26, 27, 28, 29, 30,
31.
• P. 229-230 Parallel Exercises Group A.
# 12, 13, 14, 15, 16, 17, 18, 19.
New Book: p. 243-246 # 3, 6, 7, 9, 10, 11, 12, 13,
17, 19, 22, 23, 24, 25, 26, 27, 28, 29, 30, 33, 39,
40, 41, 42, 43, 44, 48, 49, 50
Review for Chapter 8
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Dalton’s atomic theory
Thomson’s experiment: Discovery
of electrons.
Rutherford’s experiment: The
nucleus of the atom: protons and
neutrons.
Atomic number, mass number and
atomic mass. Atomic mass unit
(AMU), #p+, #e-, #no.
Isotopes
Symbols for atoms (isotopes)
showing atomic number and mass
number.
Bohr model of the atom.
Orbits, energy levels, electrons
absorbing and emitting energy as
a photon (a quantum of energy),
and corresponding to the colors of
the visible spectrum.
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Symbols, charges and masses for
protons, neutrons and electrons.
Quantum Mechanical Model of the
atom.
The periodic table: arrangement,
groups, periods, metals,
nonmetals and metalloids.
Electron dot structures for
elements in groups IA-VIIIA,
valence electrons.
Numbers of electrons in energy
levels 1-4.
Numbers of electrons in energy
levels for the first 18 elements.
Ions formed by elements in groups
IA-VIIIA.
Metals, Semimetals and
Nonmetals.
Alkali metals, alkaline earth
metals, halogens, noble gases,
representative elements, transition
elements, inner transition
elements, lanthanides (rare
earths) and actinides.