Chemistry
Periodic Table and
Electrons
Essential knowledge and skills:
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Distinguish between a group and a period.
Identify key groups, periods, and regions of elements on the periodic table.
Identify and explain trends in the periodic table as they relate to ionization energy, electronegativity,
shielding effect, and relative sizes.
Compare an element’s reactivity to the reactivity of other elements in the table.
Relate the position of an element on the periodic table to its electron configuration.
Determine the number of valence electrons and possible oxidation numbers from an element’s electron
configuration.
Write the electron configuration for the first 20 elements of the periodic table.
Essential understandings:
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The periodic table is arranged in order of increasing atomic numbers.
The names of groups and periods on the periodic chart are alkali metals, alkaline earth metals, transition
metals, halogens, and noble gases.
Metalloids have properties of metals and nonmetals. They are located between metals and nonmetals on
the periodic table. Some are used in semiconductors.
Periods and groups are named by numbering columns and rows. Horizontal rows called periods have
predictable properties based on an increasing number of electrons in the outer energy levels. Vertical
columns called groups or families have similar properties because of their similar valence electron
configurations.
The Periodic Law states that when elements are arranged in order of increasing atomic numbers, their
physical and chemical properties show a periodic pattern.
Periodicity is regularly repeating patterns or trends in the chemical and physical properties of the
elements arranged in the periodic table.
Atomic radius is the measure of the distance between radii of two identical atoms of an element. Atomic
radius decreases from left to right and increases from top to bottom within given groups.
Electronegativity is the measure of the attraction of an atom for electrons in a bond. Electronegativity
increases from left to right within a period and decreases from top to bottom within a group.
Shielding effect is constant within a given period and increases within given groups from top to bottom.
Ionization energy is the energy required to remove the most loosely held electron from a neutral atom.
Ionization energies generally increase from left to right and decrease from top to bottom of a given
group.
Electron configuration is the arrangement of electrons around the nucleus of an atom based on their
energy level.
Electrons are added one at a time to the lowest energy levels first (Aufbau Principle). Electrons occupy
equal-energy orbitals so that a maximum number of unpaired electrons results (Hund’s Rule).
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Energy levels are designated 1–7. Orbitals are designated s, p, d, and f according to their shapes and
relate to the regions of the Periodic Table.
An orbital can hold a maximum of two electrons (Pauli Exclusion Principle).
Atoms can gain, lose, or share electrons within the outer energy level.
Loss of electrons from neutral atoms results in the formation of an ion with a positive charge (cation).
Gain of electrons by a neutral atom results in the formation of an ion with a negative charge (anion).
History of the Periodic Table
Dimitri Mendeleev 1869, Professor of Chemistry at the University of Saint Petersburg (Leningrad). Mendeleev
stated that the elements vary periodically (in cycles) according to their atomic masses.
Mendeleev separated his elements and left spaces on his table in order for the periodicity to continue. He
then predicted that elements would be discovered to fill these "gaps" in the table. Mendeleev even accurately
stated the properties of these elements. Scandium(ekaboron), galluim(ekaaluminum), and
germanium(ekasilicon). By 1886 all of the elements predicted by Mendeleev had been isolated.
When Mendeleev's notes show that the periodic system was created in a single day, February 17, 1869. He
arrived at his system by puzzling over cards containing the names of the 63 known elements along with their
atomic weights and important chemical and physical properties.
Henry Moseley-1914 was a student of Rutherford. Moseley was studying X-ray formation by high energy
electron bombardment. When the atoms were plotted according to atomic number, then a linear relationship was
established. Moseley stated, "There is every reason to suppose that the integer that controls the X-ray spectrum
is the charge on the nucleus."
Periodic Law - The properties of the chemical elements are a periodic function of atomic number.
Why is Mendeleev given Credit in Modern Text Books?
Mendeleev's Table allowed for and was capable of adjusting to future discoveries:
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noble gases, new column in 1894-1901
incorporation of the rare earth elements
Moseley's atomic number in 1914
Bohr atom and electronic structure in 1913
discovery of synthetic elements 1939 to present (element 110, 1994)
The Periodic Table
Group
a vertical column of elements in the periodic table; also called a family
Period
a horizontal row of elements in the periodic table
Metals
one of a class of elements that includes a large majority of the known
elements; metals are characteristically lustrous, malleable, ductile, and good
conductors of heat and electricity
Metalloids
The elements that border the stair-stepped line are classified as metalloids. The
metalloids, or semimetals, have properties that are somewhat of a cross
between metals and nonmetals.
Metalloids tend to be economically important because of their unique
conductivity properties (they only partially conduct electricity), which make
them valuable in the semiconductor and computer chip industry. The
metalloids are shown in the following illustration.
Nonmetals
one of a class of elements that are not lustrous and are generally poor
conductors of heat and electricity; nonmetals are grouped on the right side of
the periodic table
Alkali metals
any metal in Group 1 of the periodic table. (soft, malleable, lustrous, good
conductors, MOST REACTIVE family of metals)
Alkaline earth metals
any metal in Group 2 of the periodic table. (higher densities and melting points
than alkali metals; not as reactive as alkali)
Halogens
any member of the nonmetallic elements in Group 17 in the periodic table. (
MOST REACTIVE Non-Metals; do not occur free in nature; commonly found in
sea water, minerals, & living tissues)
Noble gases
any member of a group of gaseous elements in Group 18 in the periodic table.
(VERY INACTIVE elements, used in balloons, scuba diving tanks, light bulbs)
Electrons in Atoms
Wave Nature of Light
3 waves/sec
Wavelength (λ)
Frequency (ν)
5 waves/sec
wavelength…
amplitude
Basic wave characteristics:
wavelength – shortest distance between equivalent points on a continuous wave
 symbol…. λ....Greek letter lambda
 usually expressed in m, cm or nm
frequency – the number of waves that pass a point per second
 Greek symbol υ Greek letter nu
 expressed in Hz (hertz) 1 Hz = 1 s-1
amplitude – wave’s height from the origin to the crest or trough
speed – all electromagnetic waves travel at the speed of light
 symbol….c = 3.00 x 108 m/s in a vacuum
Electromagnetic Spectrum
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all forms of electromagnetic radiation traveling in waves
vary by the frequency and wavelength
energy increases with increasing frequency
Particle Nature of Light
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the wave behavior of light cannot explain why HEATED OBJECTS give off
distinct colours (specific frequencies) of light
Max Planck (1858 – 1947)
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conclusion: matter can gain or lose energy only in specific amounts
QUANTUM
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– minimum amount of energy that can be gained or lost by an atom
hot objects emit light in quantized amounts
Dual wave-particle model of light
 light can act as a particle or a wave
Atomic Emission Spectra
 a set of frequencies of waves emitted by atoms
 observed as a set of lines of individual colours
 each element has a unique emission spectrum
 can be used to identify an element
Continued Development of Atomic Theory
Remember the models of Dalton, Thomson and Rutherford.
Bohr’s Model
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defined energy levels
electrons only move in distinct orbits around the nucleus
nicknamed the “Planetary Model”
Bohr’s model explained the Hydrogen emission spectrum
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e – are naturally in a ground state (lowest possible energy)
e – gain energy to reach an excited state (higher energy)
e – release energy in a different form to achieve a lower energy state
The Quantum Mechanical Model
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Bohr’s model could only be used to describe the H emission spectrum
led to the “wave-particle duality of electrons”
Heisenberg Uncertainty Principal
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It is impossible to know the velocity and position of a particle at the same
time
Therefore we can not predict EXACTLY where the electron is located
AND how fast it is moving.
Orbital - a specific wave function (we called it an area of high probability).
Heisenberg uncertainty principle - a fundamental limitation to just how precisely we
can know both the position and momentum of a particle at a given time.
Describing the location of the electron
Atomic orbital – three-dimensional region around the nucleus that describes the electron’s
probable location
Quantum numbers
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describe the location of the electron in four categories
each category gets more specific
Energy Level (principal quantum number)
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n
which can have values 1,2,3,4,5,6,7
defines the size, as n increases the energy level gets larger
Energy Sublevels (angular momentum quantum number)
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the sublevel have labels
1st one in each level…s
2nd…p
3rd …d
4th…f
Atomic Orbitals (magnetic quantum number)
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each sublevel has a fixed number of orbitals
s …. 1 orbital
p …. 3 orbitals
d …. 5 orbitals
f …. 7 orbitals
Atomic Spin (spin quantum number)
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each orbital can hold a maximum of 2 electrons
each energy level has “2n2” number of electrons
Electron configurations
1. the arrangement of electrons in an atom
2. lowest energy and most stable
Rules of electron arrangement
1.
AUFBAU PRINCIPLE
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– each e – occupies the lowest energy orbital
each sublevel has a different ENERGY STATE
e – within a energy level fill in sub level order…s,p,d,f
2. PAULI EXCLUSION PRINCIPLE – an atomic orbital contains a maximum of two electrons
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the two e – will travel with opposite spins
direction of spin will be represented by
one pair of e –
3. HUND’S RULE – e – will individually occupy equal energy orbitals before forming a pair
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all orbitals of a sublevel are of EQUAL energy
Methods of notation
ORBITAL NOTATION – shows every electron with an arrow
_______ ______ ______ _____ _____
O
1s
2s
2p
2p
2p
ELECTRON CONFIGURATION – shows the total number of electrons in each sublevel as a
superscript.
Ar 1s2 2s2 2p6 3s2 3p6
ELECTRON DOT – shows each outer level electron as a dot.
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maximum number of 8 dots
Ne
Valence electrons
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electrons in the highest number energy level
found in the highest number s and p sublevels
the electrons used in electron dot notation
Atoms in the same group have similar chemical properties because they have the same number
of valence electrons.
Notice that the electron dot structures repeat as you move down the table.
PERIODIC TRENDS NOTES
Periodic Law – Similar properties recur periodically when elements are arranged according to
increasing atomic number.
Properties of elements – Primarily determined by the outer shell (valence) electrons.
Basic Assumptions
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The electrons are attracted to the nucleus
o energy is required to remove an electron from an atom
o energy is released when an atom gains an electron
The further away an electron is from the nucleus, the easier it is to remove
o The higher the energy level, the further away the electron (think of the energy
levels as shells)
Shielding Effect
o Inner core (inner shell) electrons shield the nucleus from the outer shell electrons
o Electrons in the same shell do not shield each other
o This shielding is not 100% complete
Effective Nuclear Charge
o The attractive force experienced by the outer shell electrons
Atomic Radius (AR)
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Group Trend: Down a group, AR increases, due to the addition of energy levels
Periodic Trend: Across a period, AR decreases, due to the increased effective
nuclear charge
Shielding Effect: The two main assumptions of the shielding effect are:
1. The inner core electrons shield the nucleus from the valence
electrons
2. The valence electrons do not shield each other
note: These assumptions are not 100% accurate
Ionic Radius (IR)
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Cations: Positive ions (shown in blue) are smaller than neutral atoms of the same
element, the atoms have lost outer shell electrons, often the entire outer shell
Anions: Negative ions are larger than neutral atoms of the same element, the
atoms have gained electrons, the nucleus cannot pull as strongly on each electron
Trends: The trends within the cations and within the anions are the same as for
atomic radius
Ionisation Energy (IE)
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The energy required to remove an electron, measured in kJ / mole
Process of removing an electron is endothermic (requires energy)
The larger the atom, the easier it is to remove an electron
General trends are the opposite of the atomic radius trends
o Group Trend: Down a group, IE decreases
o Periodic Trend: Across a period, IE increases
Electronegativity
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The relative tendency of an atom to attract a bonding pair of electrons when the
atom is chemically combined with another atom
High EN (non metals) – tendency to attract the bonding pair (often win the “tug of
war”)
Low EN (metals) – lower tendency to attract the bonding pair (often loses the
“tug of war”)
Two elements with very different EN values will form ionic bonds, a transfer of
electrons to form ions
Two elements with very close EN values will form covalent bonds, a sharing of
electrons to form molecules
General trends are the opposite of the atomic radius trends
o Group Trend: Down a group, EN decreases
o Periodic Trend: Across a period, EN increases
Physical and Chemical Properties
Non-metals:
 Melting point, boiling point, and the state of matter at room temperature change
gradually within a group
o Group Trend: Down a group, melting point and boiling point increases,
state goes from gas → (liquid) → solid
o Periodic Trend: No predictable trend
 Reactivity increases down a group
 Metallic character increases down a group
Metals:
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Metallic Character
o Group Trend: Down a group, metallic character increases
o Periodic Trend: Across a group, metallic character decreases
Reactivity increases down a group