CHEMICAL

PERIODICITY
CHEM 16
OVERVIEW
• ELECTRON CONFIGURATION
• PERIODIC PROPERTIES OF THE ELEMENTS
• METALS, NON-METALS, METALLOIDS
RECALL
Quantum Mechanical Model of the Atom
Electrons in an atom move in three-dimensional space around the nucleus, but NOT in
an orbit that has a definite radius

Orbital
• Region in space where the electron is most likely to be found
• Characterized by the four quantum numbers

Quantum Number Symbol Values

Principal n 1, 2, 3,…
Angular Momentum ! 0 to n⎯1
Magnetic m! ⎯! to 0 to +!
Spin ms ⎯½, +½
ELECTRON CONFIGURATION
The electron configuration of an atom is a designation of how electrons are
distributed among various orbitals in principal shells and subshells.

Ground state electron configuration – most stable electron configuration

RULES FOR ASSIGNING ELECTRONS TO ORBITALS

1. Electrons occupy in a way that minimizes the energy of the atom.
2. No two electrons in an atom can have all four quantum numbers alike.
3. When orbitals of identical energies (degenerate orbitals) are available, electrons
initially occupy these orbitals singly.
AUFBAU PRINCIPLE
RULES FOR ASSIGNING ELECTRONS TO ORBITALS
1. Aufbau Principle: Electrons occupy in a way that minimizes the energy of the
atom. Orbitals are filled in order of increasing energy, with no more than two
electrons per orbital.

• Energy increases with increasing values of n.

• For a given n, energy increases with increasing value of !

Madelung’s rule (n + ℓ rule)

• Lower n + ℓ (lower energy): first to be filled
e.g. 4s vs. 3d fill 4s first

• Equal n + ℓ values: the orbital with a lower n value is filled first

Order of filling electronic subshells e.g. 2p vs. 3s fill 2p first
PAULI EXCLUSION PRINCIPLE
RULES FOR ASSIGNING ELECTRONS TO ORBITALS
2. Pauli Exclusion Principle: No two electrons in an atom can have all four quantum
numbers alike.

(n, !, m!, ms) for electrons in 1s:

(1, 0, 0, +½) and (1, 0, 0, ⎯½)

(n, !, m!, ms) for electrons in 2s:

(2, 0, 0, +½) and (2, 0, 0, ⎯½)
HUND’S RULE (OF MAXIMUM MULTIPLICITY)
RULES FOR ASSIGNING ELECTRONS TO ORBITALS
3. Hund’s Rule: When orbitals of identical energies (degenerate orbitals) are
available, electrons occupy these orbitals singly first before pairing begins.
• This is to minimize repulsion of negative electrons

DEGENERATE ORBITALS

Order of filling 2p orbitals in a carbon atom using the Hund’s rule

PERIODIC TABLE He is
part of
s-block
s-block
p-block

d-block
(Transition metals)

f-block (Lanthanides)
WRITING ELECTRON CONFIGURATIONS
METHODS OF REPRESENTING ELECTRON CONFIGURATIONS
1. Expanded notation

Example: Write the expanded electron configuration of phosphorus (P).

Step 1: Determine the total number of electrons of the element.

P (Z = 15) has 15 electrons.

Step 2: Apply the Aufbau principle and assign the total number of possible electrons in
each orbital (maximum: s=2, p=6, d=10, f=14).
1s2 2s2 2p6 3s2 3p3
WRITING ELECTRON CONFIGURATIONS
METHODS OF REPRESENTING ELECTRON CONFIGURATIONS
2. Orbital diagram
Example: Illustrate the correct ground state atomic orbital diagram of phosphorus.
Step 1: Draw a box (or a line) for each orbital in a subshell. Aufbau principle should not be
violated.

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Step 2: Fill each orbital with two electrons by not violating Pauli exclusion principle and Hund’s
rule.
↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑ ↑ ↑
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WRITING ELECTRON CONFIGURATIONS
METHODS OF REPRESENTING ELECTRON CONFIGURATIONS
3. Condensed notation (or the noble gas core electron configuration)
Example: Write the condensed electron configuration of phosphorus.

Step 1: Determine the noble gas core (i.e., the nearest noble gas that precedes the
element).
For P: Noble gas core is Ne.

Step 2: Enclose the noble gas core in a square bracket and write the rest of the
electron configuration
[Ne] 3s2 3p3 Note: 1s2 2s2 2p6 is equivalent to the electron configuration of Ne.
WRITING ELECTRON CONFIGURATIONS
Orbital Diagram Expanded notation
Condensed
Condensed Notation
Configuration
[He] 2s 1
[ He] 2s 2
[ He] 2s 2 2p 1
[ He] 2s 2 2p 2
[ He] 2s 2 2p 3
[ He] 2s 2 2p 4
[ He] 2s 2 2p 5
[ Ne] 2s2 2p6 = [Ne]
[He]
1s 2s 2p [Ne] will be the next noble gas core
EXCEPTIONS
There is an extra measure of stability associated with half-filled and completely filled d and f
orbitals.
Examples:
For Cr, we might expect that the electron configuration is [Ar] 4s2 3d4
However, a half full d subshell is more stable. Therefore, the preferred electron
configuration is: [Ar] 4s1 3d5
[Ar] ↑ ↑ ↑ ↑ ↑ ↑
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half-filled degenerate d orbitals

For Cu, we might expect that the electron configuration is [Ar] 4s2 3d9
However, a completely filled d subshell is more stable. Therefore, the preferred electron
configuration is: [Ar] 4s1 3d10
[Ar] ↑ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓
4# 3(
completely filled degenerate d orbitals
VALENCE ELECTRON CONFIGURATION
Valence electrons (also known as outermost electrons) are electrons found at the
outermost shell of an atom. These are the electrons involved in chemical reactions.

For example, F (with an electron configuration of 1s2 2s2 2p5) contains 7 valence
electrons found at its outermost shell (n = 2). Its valence configuration is 2s2 2p5.

Na (electron configuration: 1s2 2s2

2p63s1) has one valence electron. To derive the valence electron configuration of an
Valence configuration: 3s1. element, we use the following table:
Valence electron configuration
valence e- s-block nsx
p-block ns2 npx
d-block ns2 (n-1)dx
f-block ns2 (n-1)d1 (n-2)fx
n = outermost shell
ELECTRON CONFIGURATION OF IONS
Ions are charged species which are products of either a gain or a loss of electron/s of
the parent element.

Cation: loss of valence electrons in the outermost shell of atoms to gain positive
charge
e.g. Na+ (Na lost an electron)
valence e-
Na: 1s2 2s2 2p6 3s1 Na+: 1s2 2s2 2p6

Anion: gain of electrons in the outermost shell of atoms to gain negative charge
e.g. F– (F gained an electron)
F: 1s2 2s2 2p5 F– : 1s2 2s2 2p6
ISOELECTRONIC
Isoelectronic – state by which two species are of the same electron configuration
e.g. Na+, F-, Al3+, O2-, and N3- are all isoelectronic with Ne.
Electron configuration: 1s2 2s2 2p6

Complete the table below:

Ion formed after Electron Atom

Electron
Element gain/loss of configuration of isoelectronic
configuration
electron ion with

Li 1s2 2s1 Li+ ? ?

S 1s2 2s2 2p63s2 3p4 S2- ? ?
ISOELECTRONIC

Ion formed after Electron Atom

Electron
Element gain/loss of configuration of isoelectronic
configuration
electron ion with

Li 1s2 2s1 Li+ 1s2 He

S 1s2 2s2 2p63s2 3p4 S2- 1s2 2s2 2p63s2 3p6 Ar

Stable ions are formed when a noble gas electron configuration is attained upon losing
or gaining electron/s.
e.g. K+, Ca2+, Cl-, S2- are stable ions that are isoelectronic with Ar.
ELECTRON CONFIGURATION OF d-BLOCK IONS
Transition metals (d-block elements) form cations by losing its ns electrons
first before losing the appropriate number of its (n⎯1)d electrons

e.g. Mn3+ (Mn lost 3 electrons)

Mn: [Ar] 4s2 3d5 Mn3+: [Ar] 3d4
The two 4s electrons are
removed first before removing
Fe3+ (Fe lost 3 electrons) an electron from 3d.
Fe: [Ar] 4s2 3d6 Fe3+: [Ar] 3d5

Cr3+ (Cr lost 3 electrons)

Cr: [Ar] 4s1 3d5 Cr3+: [Ar] 3d3
MAGNETIC PROPERTIES
Paramagnetism – exhibited by atoms/ions with unpaired electron/s
Paramagnetic materials are weakly attracted by an externally applied magnetic
field, and form internal, induced magnetic fields in the direction of the applied field.
e.g. Al is paramagnetic.
↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑
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Diamagnetism – exhibited by atoms/ions with no unpaired electrons

Diamagnetic materials create an induced magnetic field opposite to an externally
applied magnetic field, and are repelled by the applied field.
e.g. Zn is diamagnetic.

↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓
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SAMPLE PROBLEMS
A. Which of the following atoms/ions is/are isoelectronic?
Kr, Br⎯, Se, Sr2+, Cs+, Ba

B. Which of the following atoms/ions is/are paramagnetic?

Mn2+, Ne, Cr3+

ANSWERS:
A. Kr, Br ⎯, and Sr2+ are isoelectronic with each other.
B. Mn2+ and Cr3+ are paramagnetic (they have unpaired electrons).
PERIODIC PROPERTIES
Periodic Trends
• properties of elements that behave predictably as we go vertically or horizontally
through the periodic table

1. Atomic size
2. Ionic size
3. Ionization energy
4. Electron affinity
5. Electronegativity
EFFECTIVE NUCLEAR CHARGE, Zeff
Effective nuclear charge (Zeff) is the net positive charge experienced by an
electron in a multi-electron atom

)*++ = ) − .
where /011 = effective nuclear charge
Z = nuclear charge
3 = shielding constant

Shielding effect – describes the decrease in attraction between an electron and the
nucleus in any atom with more than one electron shell
– results from electron⎯electron repulsions, which cancel some of the
attraction of the electron to the nucleus
EFFECTIVE NUCLEAR CHARGE, Zeff nucleus
Na has 11 protons Z = +11

Na has 11 electrons 1 valence electron and 10 inner

(core) electrons. valence e-
The valence electron of Na (3s1) will be shielded from
the nucleus by the 10 core electrons )*++ < )

Mg has 12 protons Z = +12

Na: 1s2 2s2 2p6 3s1
Mg has 12 electrons 2 valence electrons and 10 inner
(core) electrons.

Since Mg has more protons than Na, and the number of Higher )*++ means greater
core electrons for both atoms is the same, )*++ of Mg is attraction between the nucleus
higher than that of Na. and the electrons.
EFFECTIVE NUCLEAR CHARGE, Zeff
increasing Zeff
Zeff increases in the same period
because the number of core electrons
is the same while the number of

increasing Zeff
protons increases.

Zeff increases gradually down a group

because the more diffuse core electron
cloud is less able to screen the valence
electrons from the nuclear charge.

TREND:
• Across a period: Zeff increases
• Down a group: Zeff increases
ATOMIC RADIUS (SIZE)
Atomic radius – half the distance between covalently bonded atoms

increasing atomic radius

increasing atomic radius

TREND:
• Across a period: atomic
radius decreases
• Down a group: atomic
radius increases
ATOMIC RADIUS

Across a period: Electrons are

added to the same shell
increase number of protons but same
number of core electrons greater
nuclear attraction (↑ Zeff) valence
electrons are pulled closer to the
nucleus decrease in atomic radius
Down a group: Electrons are added
to a higher shell
increase n increase orbital size
increase in atomic radius
IONIC RADIUS
TREND:
1. Cations are smaller than parent atoms due
to removal of electron/s (increase Zeff).
2. Anions are larger than parent atoms due to
addition of electron/s (increase repulsion,
decrease Zeff).
3. For isoelectronic ions (same number of
electrons), increase number of protons
increase Zeff decrease in ionic size
e.g. O2- > F- > Na+ > Mg2+ > Al3+
4. Going down a group, ionic size increases
due to increasing n.
IONIZATION ENERGY
Ionization energy is the minimum energy required to remove an electron from
a gaseous atom or ion.

Na(9) + <=<>?@ → NaB9 +< C IE1 = +496 kJ/mol

increasing IE General Trend:

• Across a period:
IE increases
increasing IE

• Down a group: IE
decreases

Note: Lower IE means it’s

easier to remove an electron
easier to form a cation
IONIZATION ENERGY
In each period, the noble gases have the highest IE. They are the most stable (inert)
because they have completely filled electron shell.

General Trend
Across a period: greater nuclear
attraction (↑ Zeff) valence
electrons are held more tightly
(more difficult to remove)
increase in IE

Down a group: increase in atomic

radii valence electrons are
farther from the nucleus (easier to
remove) decrease in IE
IONIZATION ENERGY ⎯ DEVIATIONS FROM GENERAL TREND
Examples: IE1 of Be > IE1 of B

Be ↑↓
2p subshell is at a higher energy
$# $% than 2s subshell electron in 2p
is easier to remove
B ↑↓ ↑
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IE1 of N > IE1 of O

all three p electrons are
N ↑↓ ↑ ↑ ↑ unpaired electron repulsion
is minimized
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repulsion of paired electrons
O ↑↓ ↑↓ ↑ ↑ makes it easier to remove the
Atomic number 4th p electron.
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SUCCESSIVE IONIZATION ENERGIES
• First ionization energy (IE1): energy required to remove the outermost electron

Ca(9) + <=<>?@ → CaB9 +< C IE1 = +599 kJ/mol

• Second ionization energy (IE2): energy required to remove the subsequent high-
energy electron

CaB9 + <=<>?@ → CaPB

9 +<
C
IE2 = +1145 kJ/mol

• Generally, the next ionization energy is larger than the preceding (e.g. IE3 > IE2 >
IE1). With each successive removal, an electron is being pulled away from an
increasingly positive ion (higher Zeff), and this requires increasingly more energy.
ELECTRON AFFINITY
Electron affinity is the amount of energy change when an electron is added to a
gaseous atom or ion. It is a measure of an atom’s ability to form negative ions.

Cl(9) + < C → ClC

(9) + <=<>?@ EA = ⎯349 kJ/mol
negative sign denotes that energy is released

increasing EA General Trend*

(magnitude of EA):
• Across a period:
EA increases
increasing EA

• Down a group:
EA decreases
*exclude noble gases

Note: Higher EA value means

it’s easier to gain an electron
easier to form an anion
ELECTRON AFFINITY
For noble gases, EA is positive, which means an electron will not attach itself to a noble
gas atom. Noble gases are inert because they have stable electron configuration.

General Trend for magnitude of EA

(excluding noble gases):
Across a period: greater nuclear attraction
(↑ Zeff) easier to gain an electron
increase in EA
Also, elements toward the right (e.g.
halogens, chalcogens) could attain a more
stable noble gas configuration by gaining
electron/s.
Down a group: valence electrons are farther
Note that there are deviations from the general from the nucleus less attractive toward
trend, which can be explained by analysis of the
electron configuration electrons decrease in EA
ELECTRONEGATIVITY
Electronegativity is the tendency of an atom to attract electrons to itself when
chemically combined with another element
increasing electronegativity
Across a period (excluding noble
gases): ↑ proton number greater

increasing electronegativity
nuclear attraction increase in
electronegativity

Down a group: valence electrons

are farther from the nucleus less
“desire” to grab other electrons
decrease in electronegativity

TREND:
F is the most electronegative element.
• Across a period: electronegativity increases
• Down a group: electronegativity decreases
SAMPLE PROBLEMS
Arrange the following in decreasing order in terms of the given parameter:

A. Atomic radius
F, As, Ga, Rb
B. Ionic radius
S2-, K+, Ca2+, Cl-
C. Ionization energy
K, Se, F, Br
D. Electronegativity
Al, F, Na, P, O
SAMPLE PROBLEMS
Arrange the following in decreasing order in terms of the given parameter:

A. Atomic radius
F, As, Ga, Rb Answers:
B. Ionic radius A. Rb > Ga > As > F
S2-, K+, Ca2+, Cl- B. S2- > Cl- > K+ > Ca2+
C. F > Br > Se > K
C. Ionization energy
D. F > O > P > Al > Na
K, Se, F, Br
D. Electronegativity
Al, F, Na, P, O
Bring a Periodic Table for your quiz on Thursday.
-end