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UNIT #3: Electrons in Atoms/Periodic Table and Trends
1. ELECTRON CONFIGURATION
Electrons fill the space surrounding an atom’s nucleus in a very specific order following the
rules listed below:
a) Aufbau Principle: Each electron occupies the lowest energy orbital available. The
orbitals closest to the nucleus have the lowest energy; the orbitals farthest from the nucleus
have the highest energy.
Order of increasing energy:
1s→2s→2p→3s→3p→4s→3d→4p→5s→4d→5p→6s→4f→5d→6p→7s→5f→6d→7p
b) Pauli Exclusion Principle: A maximum of two electrons may occupy a single orbital, but
only if the electrons have opposite spins. Each electron in an atom has an associated spin,
similar to the way a top spins on its axis. Like a top, an electron can spin in only one of
↑ for an electron
two directions. In an orbital diagram, this is represented by an arrow up
↓ for an electron spinning in the opposite
spinning in one direction, and an arrow down
direction.
c) Hund’s Rule: Single electrons with the same spin must occupy each equal-energy orbital
before additional electrons with opposite spins can occupy the same orbitals. This is due
to the fact that electrons carry like negative charges and thus, repel each other. An electron
will pair up with another electron within a given sublevel (s,p,d,f) only when necessary and
in doing so, adopts the opposite spin.
Key Terms:
1. Principle Energy/Quantum Level: Major energy levels surrounding the nucleus of an
atom. Consists of n=1, n=2, n=3, n=4, n=5, n=6, n=7 (corresponding to periods 1 through
7 on the periodic table).
2. Energy Sublevels: Within a principle energy level, electrons occupy sublevels labeled s,
p, d or f according to the shape of the atom’s orbital. S-orbitals are spherical in shape; p-
orbitals are dumbbell shaped; d and f orbitals have varying shapes.
3. Orbitals: Within a sublevel, electrons occupy a specific number of orbitals, each of which
contain up to one pair of electrons with opposite spins. The number of orbitals within a
sublevel is as follows:
S-sublevel: Contains one orbital which contains a maximum of 2 electrons.
P-sublevel: Contains three orbitals, each of which contains a maximum of 2 electrons.
Maximum number of p-sublevel electrons is six.
D-sublevel: Contains five orbitals, each of which contains a maximum of 2 electrons.
Maximum number of d-sublevel electrons is ten.
F-sublevel: Contains seven orbitals, each of which contains a maximum of 2 electrons.
Maximum number of f-sublevel electrons is fourteen.
4. Valence Electrons: Electrons occupying the outermost principle energy level.
Electron Configuration: Denotes the filling of electrons according to the rules listed above.
The configurations depict the principle energy level of each electron (coefficient 1 through
7), followed by the sublevel (s,p,d,f), followed by a superscript that represents the number of
electrons. NOTE: Electrons filling sublevel d drop one energy level and electrons filling
sublevel f drop two energy levels.
Order of filling sublevels according to aufbau principle:
Period 1 atoms: 1s
Period 2 atoms: 1s, 2s, 2p
Period 3 atoms: 1s, 2s, 2p, 3s, 3p
Period 4 atoms: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p
Period 5 atoms: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p
Period 6 atoms: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p
Period 7 atoms: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p
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Ex. He: 1s (2 electrons in atom)
2 2 6
Ne: 1s 2s 2p (10 electrons in atom)
2 2 6 2 6
Ar: 1s 2s 2p 3s 3p (18 electrons in atom)
2 2 6 2 6 2 10 6
Kr: 1s 2s 2p 3s 3p 4s 3d 4p (36 electrons in atom)
2 2 6 2 6 2 10 6 2 10 6
Xe: 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p (54 electrons in atom)
2 2 6 2 6 2 10 6 2 10 6 2 14 10 6
Rn: 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p (86 electrons in atom)
NOTE: In these examples, each atom (other than helium) contains 8 valence electrons.
This is the stable octet that all other atoms strive to achieve. When atoms become ions,
they either lose electrons (metals) or gain electrons (non-metals) to achieve a stable
principle energy level similar to their closest noble gas.
More examples of neutral atoms versus their corresponding ions:
2 2
Be 1s 2s neutral beryllium atom with 4 electrons
2+ 2
Be 1s beryllium ion with 2 electrons (lost 2)
2 2 6 1
Na 1s 2s 2p 3s neutral atom with 11 electrons
+ 2 2 6
Na 1s 2s 2p sodium ion with 10 electrons (lost 1)
2 2 4
O 1s 2s 2p neutral oxygen atom with 8 electrons
2- 2 2 6
O 1s 2s 2p oxide ion with 10 electrons (gained 2)
2 2 6 2 3
P 1s 2s 2p 3s 3p neutral phosphorous atom with 15 electrons
3- 2 2 6 2 6
P 1s 2s 2p 3s 3p phosphide ion with 18 electrons (gained 3)
Orbital Diagrams: Denotes each orbital within a sublevel and the electrons occupying those
orbitals (indicated by an up arrow ↑ or a down arrow ↓). Electrons fill orbitals singularly at
first, then pair as necessary with an opposite spin.
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Ex. 2p ↑↓ ↑ ↑
2p 2p 2p
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3d ↑↓ ↑↓ ↑_ ↑_ ↑_
3d 3d 3d 3d 3d
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2. ELEMENTS AND THE PERIODIC TABLE
a) An element is a pure substance that cannot be separated into simpler substances by
physical or chemical means.
b) Each element has a unique chemical name and symbol. The chemical symbol consists of
one, two or three letters: the first letter is always capitalized and the remaining letter(s) are
always lowercase.
c) Seven elements occur in nature as diatomic molecules (2 atoms) because the molecules
formed are more stable than the individual atoms. They are Br , I , N , Cl , H , O , F .
Remember it as BrINClHOF. 2 2 2 2 2 2 2
d) On earth, 91 elements are naturally occurring and their abundance in the universe varies.
e) The Periodic Table organizes the elements according to increasing atomic number.
1. Elements are arranged in vertical columns called groups or families. Each group is
numbered 1 through 18.
2. Groups 1, 2, 13, 14, 15, 16, 17 and 18 are often referred to as the main group, or
representative elements, because they possess a wide range of chemical and physical
properties.
3. Groups 3, 4, 5, 6, 7, 8, 9, 10, 11 and 12 are referred to as the transition elements.
4. Elements in the same group have similar chemical and physical properties.
5. Elements are arranged in horizontal rows called periods. Beginning with hydrogen in
period 1, there are a total of 7 periods.
f) Classification of Elements
1. Metals are elements that are generally shiny when smooth and clean, solid at room
temperature, and good conductors of heat and electricity. Most metals are malleable
(can be pounded into thin sheets) and ductile (can be drawn into wires).
a) Used to transmit electrical power, ex. copper.
b) Can be formed into coins, tools, fasteners and wires.
c) Group 1 elements (except hydrogen) are known as the alkali metals.
d) Group 2 elements are known as the alkaline earth metals.
e) Both alkali and alkaline earth metals are chemically reactive, with alkali
metals being the more reactive group.
f). Groups 3 through 12 elements are divided into
1. transition metals-located in periods 4 through 7.
2. inner transition metals-two sets of inner transition metals, known as the
lanthanide and actinide series, appear at the bottom of the periodic table
and are usually offset from the numbered periods. These elements are
phosphors, substances that emit light when struck by electrons.
2. Nonmetals are elements that are generally gases or brittle, dull-looking solids. They
are poor conductors of heat and electricity. The only non-metal that is a liquid at room
temperature is bromine.
a) Group 17 elements are the halogens. These are the most reactive non-metals.
b) Group 18 elements are the noble gases-extremely unreactive due to the most
stable and complete electron configuration.
3. Metalloids or semimetals are elements with physical and chemical properties of both
metals and nonmetals.
a) Located on the right hand side of the periodic table and form a stair-step
pattern between the transition metals and the nonmetals.
b) Consists of B, Si, Ge, As, Sb, Te and At.
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3. COMPOUNDS AND LAWS OF DEFINITE/MULTIPLE PROPORTIONS
a) A compound is a combination of two or more different elements that are combined
chemically. Much of the matter of the universe are compounds; there are approximately
10 million known compounds.
Examples are water, table salt, table sugar, aspirin.
b) Compounds or elements that occur alone are referred to as pure substances. Compounds
or elements that occur in combination with other compounds or elements are referred to as
mixtures.
1. Homogenous mixture-one that has a uniform composition throughout and always has
a single phase; can be separated by physical means such as distillation (a technique
used to separate mixtures based on the differences in the boiling points of the
substances) or by evaporation (removing liquid component from solid component);
homogenous mixtures are also referred to as solutions.
Ex. salt water, sugar water, lemonade, gasoline, steel.
2. Heterogeneous mixture-one that does not have a uniform composition and in which
the individual substances remain distinct; can be separated by physical means such as
filtration (technique that uses a porous barrier to separate solids from liquids).
Ex. sand and water, dirt, Italian salad dressing.
c) Law of Definite Proportions
1. Elements making up compounds always combine in definite proportions by mass.
Regardless of the amount of a given compound, it is always composed of the same
elements in the same proportion by mass.
d) Law of Multiple Proportions
1. When different compounds are formed by combinations of the same elements,
different masses of one element combine with the same relative mass of the other
element in a ratio of small whole numbers.
2. Examples:
a) Water is H2O: 2 parts hydrogen to 1 part oxygen
Hydrogen Peroxide is H O 2 parts hydrogen to 2 parts oxygen
2 2:
Both compounds are comprised of the same elements; however, H O differs from
2 2
HO in that it has twice as much oxygen. When we compare the mass of oxygen
2
in H O to the mass of oxygen in H O, we get the ratio 2:1.
2 2 2
b) Methane is CH ; Carbon = 12amu and Hydrogen = 4amu; C : H = 12:4 or
4 mass mass
3:1
Ethane is C H ; Carbon = 24amu and Hydrogen = 6amu; C : H = 24:6 or
4:1 2 6 mass mass
4. PERIODIC TABLE TRENDS
a) Atomic Radius
1. The radius of an atom is one-half the distance between the nuclei of two atoms of the
same element when the atoms are joined; it is comparable to the radius of a circle which
is the length of a line from the center of the circle to its edge.
2. Radius decreases as you move across a period. As you move across a period, each
successive element has one additional proton in its nucleus; therefore, the positive
nuclear pull increases on the negative electrons surrounding the nucleus, causing the
radius to decrease.
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