INTRODUCTION TO ELECTRON ARRANGEMENTS
I. INTRODUCTION TO ELECTRON CONFIGURATION OF AN ATOM or how
electrons are arranged in the atom. Describes the motion of an electron about an atom. It is not a
definite path like that of the earth’s orbit around the sun. An electron location can be given only
in terms of probabilities, similar to population densities of a map of Kansas.
Recall: the 1st energy level in an atom has a maximum of 2 electrons in it; the second energy
level has a maximum of 8 electrons in it; the 3rd energy level has a maximum of 18 electrons in it;
the 4th a maximum of 32 electrons. The formula to find the number of electrons in the orbital =
2n2 electrons where “n” is the energy level.
Definition: Valance electrons – the “s” and “p” electrons in the outermost energy level.
Definition: Ground State and Excited State - “Ground state” of an atom is when the atom is in its
lowest energy state. When an atom is not at ground state, we say it is “excited”. As electrons
jump to higher energy levels, they absorb energy. When electrons return to a lower level, they
release light energy (photons) equal to the amount of energy that was gained.
Quantum Mechanical Model of the Atom – Developed by Erwin Schrodinger, who combined
Heisenberg's Uncertainty Principle, which states that we can never simultaneously know an
electron’s velocity and position.
This model:
 treats electrons as waves and uses mathematics to calculate probability densities of
finding the electron in a particular region in the atom
 The mathematical solutions give regions in space of high probability for finding the
electron - these are called atomic orbitals
 Each orbital can be identified by a set of values called quantum numbers.
 The energy levels are labeled by the principal quantum number (n). These are
assigned the values n = 1,2,3,4 and so forth.
 For each energy level, there may be several orbitals with different shapes and
different energies – these are called the sublevels.
 each sublevel corresponds to an orbital of a different shape.
 Different orbitals are denoted by letters:
 s = 1 kind (spherical in shape)
perfect!

p = 3 kinds (dumbbell or peanut shaped)
kind of 
 d = 5 kinds (4 of which are shaped like a clover leaf)
 f = 7 kinds (complex shaped) I’m not even going to try 
okay…maybe not 
II. Intro to electron configuration of an atom or how electrons are arranged in the atom
1. AUFBAU PRINCIPLE states electrons will enter unoccupied orbitals of
lowest energy first.
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2. HUND’S RULE states electrons having parallel spins will enter
unoccupied orbitals of the same energy level one at a time before pairing up.
3. PAULI EXCLUSION PRINCIPLE states that no two electrons can be
described by the same set of four quantum numbers. They can have the
same energy level, the same orbital, the same axis, but would have to have
opposite spin.
C. ORDER OF FILLING ELECTRONS for electron configuration.
1. The representative elements are the “s” and “p” electrons
2. Groups 1 and 2 of the representative elements are known as the “s” area.
3. Groups 3-8 of the representative elements are known as the “p” area.
4. The transition elements are known as the “d” area.
5. The actinide and lanthanide series (inner transition) are known as the “f”
area.
D. ORDER OF FILLING ORBITALS-Aufbau Principle-lowest energy first
1S
2S
3S
4S
5S
6S
7S
[2 electrons]
2P
[8 electrons; 2 from s and 6 from p]
3P
[8 electrons; 2 from s and 6 from p]
3D 4P
[18 electrons; 2 from s, 10 from d & 6 from p]
4D 5P
[18 electrons; 2 from s, 10 from d & 6 from p]
4F 5D 6P [32 electrons; 2 from s, 14 from f, 10 from d & 6 from p]
5F 6D 7P [32 electrons; 2 from s, 14 from f, 10 from d & 6 from p]
EX: Digit before letter indicates main energy level (n). Letter indicates orbital
shape. Superscript tells number of electrons.
Recall that s has one orbital, p has 3 orbitals, d has 5 orbitals and f has 7 orbitals.
Practice: write the electron configuration for the element nickel.
Write the electron configuration for the element bismuth
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III. ORBITAL NOTATION –an unoccupied orbital is represented by a line ___ with the
sublevel designation written below it.
Write the orbital notation for the element nitrogen
    
1s 2s 2p
Practice: Write the orbital notation for the element nickel
IV. ELECTRON DOT NOTATION (aka lewis dot) a notation of the outermost energy levels
only.
A. Write the symbol to designate the element.
B. Dots are used to represent the outermost “s” and “p” electrons only.
C. The number of “s” and “p” electrons used (valence electrons) is equal
to the group number in which the element is found. Only the representative
elements have “s” and “p” sublevels.
1. Group 1 has 1 valence electron.
2. Group 2 has 2 valence electrons.
3. Group 3 has 3 valence electrons.
4. Group 4 has 4 valence electrons.
5. Group 5 has 5 valence electrons.
6. Group 6 has 6 valence electrons.
7. Group 7 has 7 valence electrons.
8. Group 8 has 8 valence electrons. The outer shell is full.

EX: Write the dot notation for the element phosphorous
P:

Write the dot notation for the element calcium and chlorine.
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INTRODUCTION TO THE PERIODIC TABLE
Demitri Mendeleev (1834-1907), a Russian scientist published the first table of elements in 1869.
He listed the elements in order of increasing atomic number and categorized the elements by
matching similar properties. There were missing elements in which he “guessed” their properties
and was later found to be correct! This was no small feat.
I. PERIODIC LAW states that the properties of the elements show up periodically when
the elements are arranged in increasing order of atomic number.
A. GROUPS (a.k.a. FAMILIES) found in vertical (up and down) columns.
1. Elements are grouped by similar chemical properties.
2. Many groups have special names:
a. GROUP 1-ALKALI METALS- react with water to form strong bases
(alkai).
b. GROUP 2-ALKALINE EARTH METALS-also react with water to form
strong bases but are not as reactive as group 1 metals.
c. GROUP 7-HALOGENS- means “salt forming” so members of this group
react with metals to form salts
d. GROUP 8-NOBLE GASES (a.k.a. INERT GASES). They got their name
because they do not react well with other elements, so these snobs were
referred to as Noble. NOT REACTIVE.
e. GROUPS known as the “d” region are called TRANSITION
METALS. What does transition mean?
f. GROUPS known as the “f” region are called INNER TRANSITION
METALS
B. PERIODS (a.k.a. SERIES) are the horizontal rows, left to right.
1. PERIOD 1: H and He
2. PERIOD 2 and 3: 8 elements each
3. PERIOD 4 AND 5: 18 elements each
4. PERIOD 6: Contains the LANTHANIDE SERIES (a.k.a. inner transition
metals) with atomic numbers 57 – 71. There are 32 elements.
5. PERIOD 7: Contains the ACTINIDE SERIES (a.k.a. inner transition
metals) with atomic numbers 89 – 103. There are 32 elements.
C. LOCATIONS OF MAJOR CATEGORIES OF ELEMENTS
1. METALS are found on the left side of the zig-zag line. Properties
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2. NONMETALS are found on the right side of the zig-zag line. Properties:
3. METALLOIDS are elements that border the zig-zag line and show
properties of both metals and nonmetals. Aluminum is the exception-it is
considered a metal.
II. PERIODIC TABLE AND ELECTRON CONFIGURATION
A. Elements in the same group have the same number of electrons in their outermost
energy levels (a.k.a. valence electrons)
1. Group 1 has 1 valence electron.
2. Group 2 has 2 valence electrons.
3. Group 3 has 3 valence electrons.
4. Group 4 has 4 valence electrons.
5. Group 5 has 5 valence electrons.
6. Group 6 has 6 valence electrons.
7. Group 7 has 7 valence electrons.
8. Group 8 has 8 valence electrons. The outer shell is full.
B. ISOELECTRONIC – the name given to ions that have the same electron
configuration as atoms of noble gases.
1. IONS are charged particles that gain electrons (negative charge) or lose
electrons (positive charge) to obtain the same electron configuration as a
noble gas to become stable.
2. Metals usually lose electrons to become isoelectronic with the noble gas in
the previous energy level.
EX: K loses one electron to become K 1+ and is isoelectronic with Ar.
3. Nonmetals usually gain electrons to become isoelectronic with the nearest
noble gas.
EX: Fluorine gains one electron to become F 1− and is isoelectronic with Ne.
III. PERIODIC TRENDS
A. ATOMIC RADIUS-NEUTRAL ATOMS
1. Top to bottom within a group the atomic radius gets bigger ( more energy levels)
1s1 H (smallest atomic radius) valence electrons are CLOSE to nucleus

7s Fr (largest atomic radius) valence electrons are FAR from the nucleus
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2. From left to right within a period, atomic radius gets smaller-protons and
electrons added as you go from left to right ( more attraction to nucleus) but no NEW
(higher) energy levels are added.
3. Positive protons are attracted to negative electrons so the atom squeezes in
closer –nuclear charge
B. IONIC RADII -charged atoms (no longer neutral)
1. Within a group top to bottom:
H+
Smaller
+
Li

+
Na
K+
Larger
2. Within a period left to right, metal ions are smaller than nonmetal ions.
Examples of ionic radius vs. atomic radius
Ion Na + 1s2 2s2 2 p6
{isoelectronic with Ne}
Atom Na
1s2 2s2 2p6 3s1
Ion Cl -
1s2 2s2 2p6 3s2 3p6 {isoelectronic with Ar}
Atom Cl
1s2 2s2 2p63s2 3p5
Sodium atom is larger than the chlorine atom HOWEVER sodium ion is smaller than the
chloride ion because Na lost its outer shell (n=3). Chloride has a n=3 valence electrons,
therefore, the 3p6 in the chloride ion is bigger than 2 p6 in the sodium ion.
Mg atom
1s2 2s2 2p6 3s2
Mg 2+ ion 1s2 2s2 2p6
S atom
{isoelectronic with Ne}
1s2 2s2 2p6 3s2 3p4
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S 2- ion
1s2 2s2 2p6 3s2 3p6 { isoelectronic with Ar}
C. IONIZATION ENERGY: energy needed to remove an electron from an atom.
1. Low ionization energy is characteristic of metals (left side). Why?
2. High ionization energy is characteristic of nonmetals (right side). Why?
3. Within a group of representative elements, the ionization energy generally
decreases with increasing atomic size.
4. Within a period, ionization energy generally increases since elements become
less willing to give up electrons ( get smaller in size).
Summary of Ionization energy (I.E.) (the energy needed to remove an electron).
1st I.E. refers to energy needed to remove 1e2nd I.E. refers to energy needed to remove a 2nd e3rd I.E. refers to energy needed to remove a 3rd e- and so on….
Na has a low 1st I.E. because it has one valence e- therefore it needs to lose one electron to have
an octet.
Na has high 2nd I.E. This would require tremendous energy to remove because Na + is an octet
(isoelectronic with a noble gas – Ne). Na ++ is highly unlikely
Ca has a low 1st and 2nd I.E., but a 3rd I.E. because it is in group 2 therefore losing another
electron when it is an octet is highly unlikely.
Nonmetals all have high ionization energies because they do not want to lose electrons, they
want to gain electrons to become octets.
Guided Practice:
Would you expect Al to have a low 1st ionization energy? _____ a low 2nd ionization
energy? ____ a low 3rd ionization energy?____ a low 4th ionization energy?____
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EX:
st
6. Does Li have a high or low 1 ionization energy? ________
Why?
7. Does K have a high or low 2nd energy level? __________Why?
8. Does Ga have a high or low
1st ionization energy?
2nd ionization energy?
3rd ionization energy?
4th ionization energy?
___________
___________
___________
___________
9. In general Groups V, VI, VII will have ionization energies that get progressively
higher from left to right within a period (more energy needed to remove electrons).
D. ELECTRONEGATIVITY: ability of the nucleus of one element to attract
electrons of another element to obtain an octet (become stable).
1. Electronegativity generally increases from left to right within a period
(as the size of the atom decreases, there is more attraction to nucleus).
2. Electronegativity generally decreases from top to bottom within a group
( the size of the atom increases as more energy levels are added).
3. The most electronegative element is Fluorine and has a value of 4.0.
4. Metals have a small attraction for valence electrons (low electronegativity).
Metals want to lose electrons.
5. Nonmetals have a high attraction for valence electrons (high electronegativity).
Nonmetals want to gain electrons.
EX: Which atom has the higher electronegativity, aluminum or chlorine?
Why?
Which atom has the higher electronegativity, calcium or barium?
Why?
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QUANTUM NUMBER ASSIGNMENT-UNIT 2
2. What is the maximum number of electrons in the:
a. s orbital
b. p orbital
c. d orbital
d. f orbital
3. Write out the complete and correct order for filling the electron orbital configuration.
For the following elements, write the electron configuration and electron dot notation.
6. sodium
7. calcium
8. aluminum
9. sulfur
10. bromine
11. argon
12. carbon
13. francium
14. nitrogen
Define the following:
15. Pauli’s exclusion principle9
16. Aufbau’s principle –
17. Hund’s rule –
PERIODIC TABLE/ELECTRON CONFIGURATION ASSIGNMENT
Part A: Use the periodic table to help answer the following questions.
1. By what other name is Group 1 elements known?
2. What does the name for Group 1 mean?
3. By what other name is Group 2 known?
4. By what other name is Group 7 known?
5. What does the name for Group 7 mean?
6. By what other names are Group 8 known?
7. What is the symbol of the element in period 5 that has only 2 valence electrons?
8. What is the name of the element in period 6 that has only 4 valence electrons?
9. When bromine reacts to form an ion, it becomes isoelectronic with _________
10. When barium reacts to form an ion, it becomes isoelectronic with __________
11. Which element has the larger atomic radius, magnesium or sulfur? _________
12. Which element has the smaller atomic radius, francium or sodium? __________
13. Considering ionization energies, which element will require less energy to bond, lithium or
potassium?
____________
14. Considering electronegativity, which element is more reactive to bond, fluorine
or iodine? _____________
PART B: Relating electron configuration to the Periodic Table
Write electron configurations for the atoms that have the following atomic numbers.
15. atomic # = 12
16. atomic # =20
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17. atomic # = 38
18. What do all the above electron configurations have in common?
19. To what family or group do all the above atoms belong?
20. What are the names and symbols of each of the above atoms?
Write electron configurations for the atoms that have the following atomic numbers.
21. atomic # = 8
22. atomic # =16
23. atomic # = 34
24. What do all the above electron configurations have in common?
25. To what family or group do all the above atoms belong?
26. What are the names and symbols of each of the above atoms?
PART C: Ions-elements with a charge and the Periodic Table
27. For sodium, Na, element 11, write the electron configuration for the neutral sodium atom and
then write one for the sodium ion.
Na:
Na+
28. The sodium ion is isoelectronic with ______________.
29. Why does sodium tend to form a 1+ ion?
30. In what group of the periodic table is sodium located?
32. For chlorine, F, element 9, write the electron configuration for the neutral fluorine atom and
then write one for the fluorine ion.
F:
F1
33. The fluorine ion is isoelectronic with ______________.
34. In what group of the periodic table is fluorine located?
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PART D: Graphing analysis
36. The first seven elements in Period 3 have been estimated to have the following
radii measured in nanometers.
Na = 0.186 nm
Mg = 0.160 nm
Al = 0.143 nm
P = 0.110 nm
S = 0.104 nm
Cl = 0.099 nm
Si = 0.117 nm
a. Use the Periodic table to find the atomic number of each of these elements.
b. Label the X axis with the atomic numbers beginning with the smallest atomic
number on the graph below.
c. Label the Y axis with the values given for the radii above.
d. Graph
37. What do you notice about the relationship of atomic number to radii
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