Chapter 8
Review of Quantum Numbers
Principal Quantum Number (n)
-tells you the energy level
-n can be equal to 1, 2, 3, 4, 5, 6, 7…
-distance e- is from the nucleus inc. as n inc.
Angular Momentum Quantum Number (l)
-determines the shape of the orbital
-shapes are s, p, d, or f
-when given a value of n, l can be any integer
including zero up to n - 1
Value of l
Shape of orbital
l=0
s
l=1
p
l=2
d
l=3
f
Magnetic Quantum Number (ml)
-specifies the orientation of the orbital
-equal to integer values, including zero ranging
from +l to -l
Examples:
1) What are the quantum numbers of the orbitals
in the 3rd energy level?
n=3
l = 2, 1, 0
ml = +2, +1, 0, -1, -2 (represents d orbitals)
+1, 0, -1
(represents p orbitals)
0
(represents s orbital)
2) What are the quantum numbers of the 4p
orbital?
n=4
l = 1 (because it is in the p orbital)
ml = +1, 0, -1 (p orbitals have three orientations)
Spin Quantum Number (ms)
-electron spin is represented by the direction of
the arrow (which represents the electrons)
-all e- have the same amount of spin
-electrons can only spin in one of two directionsspin up or spin down
ms = +1/2 (spin up)
ms = -1/2 (spin down)
**Write the electron configurations for
potassium and titanium.
potassium 19e1s22s22p63s23p64s1
**shorthand way = [Ar] 4s1
titanium 22e1s22s22p63s23p63d24s2
**shorthand way = [Ar] 3d24s2
Question: What are the four quantum numbers
for each of the two e- in a 4s orbital?
n=4
l=0
ml = 0
ms= +1/2
n=4
l=0
ml = 0
ms= -1/2
Orbital Diagrams
-shows arrangement of electrons in orbitals
-symbolizes electrons as arrows and orbitals as
boxes
Rules for orbital diagrams:
1) Aufbau Principle
-electrons enter orbitals of lower energy first
spdf
-atomic orbitals are represented as boxes
s = 1 box (1 orbital)
p = 3 boxes (3 orbitals)
d = 5 boxes (5 orbitals)
f = 7 boxes (7 orbitals)
2) Pauli-Exclusion Principle
-an atomic orbital can hold at most 2 electrons
-electrons are represented as arrows
-spins are opposite
-first electron is +1/2 ↑
-second electron is -1/2 ↓
-number of e- must equal number of arrows
3) Hund’s Rule
-one electron enters each orbital of equal energy
until orbitals contain one electron, then they
can hold two e-it is more stable to have partially filled orbitals
than empty orbitals
**Draw orbital diagrams for beryllium and
sulfur.
**Draw orbital diagrams for potassium and
titanium.
Electron Configuration and the Periodic Table
periodic property- property that is predictable
based on an element’s position within the
periodic table
Modern periodic table is set up according to
Dmitri Mendeleev’s:
periodic law- when elements are arranged in
order of increasing mass, they arrange into
groups with other elements having similar
properties
-Henry Moseley later said it would be better to
arrange according to increasing atomic number
because not all masses are greater as you move
across
Ex- tellurium and iodine
valence electrons- electrons in the outermost
energy levels (highest energy level)
-important for chemical bonding because they
are held most loosely and are easier to share or
lose
-elements in the same group have similar # of
valence e- and similar chemical properties
-in transition metals the d e- are included in the
valence electrons even though they are not in
the outermost energy level
core electrons- all other e- besides the valence e-
*Identify the valence and core e- for potassium,
titanium and germanium
K = 1 valence e- and 18 core eTi = 4 valence e- and 18 core eGe = 4 valence e- and 28 core e-
-electron configurations can determine the group
of the element on the periodic table
alkali metals = ns1
alkaline Earth metals = ns2
transition metals = d block
halogens = np5
noble gases = np6
inner transition metals = f block
**Predict the outer e- config for each element:
1) strontium
2) bromine
3) cadmium
1) 5s2
2) 4s24p5
3) 4d105s2
Summary
-periodic table is divided into four blocks (s, p,
d, and f)
-the group # of a main-group element is equal to
the number of valence e-the row # of a main-group element is equal to
the highest principle quantum # of that element
Periodic Trends
1) Atomic Size
-looking at atomic radius:
-half the distance between the nuclei of two
atoms bonded together
Trend:
1) atomic radius tends to increase as you move
down a group
-as you move down a group, the n value (energy
level) increases resulting in larger atoms
2) atomic radius tends to decrease as you move
across a period
-because there are more valence e- as you move
across, there is a stronger attraction between
the outermost e- and the nucleus and it makes
it more tightly bound and therefore smaller
Examples:
1) Choose the larger atom for these pairs:
a) nitrogen or fluorine
N
b) carbon or germanium
Ge
c) nitrogen or aluminum
Aℓ
d) aluminum or germanium
unable to tell based on trends
2) Choose the larger atom:
a) tin or iodine
b) germanium or polonium
c) iron or selenium d) chromium or tungsten
a) Sn
b) cannot tell c) Fe
d) W
3) Place in order of decreasing radius:
sulfur, calcium, fluorine, rubidium, silicon
4) Place in order of increasing radius:
nitrogen, lithium, carbon, oxygen, beryllium
3) Rb, Ca, Si, S, F
4) O, N, C, Be, Li
Electron Configs and Magnetic Properties of Ions
-remember ions are atoms or groups of atoms that
have either lost or gained e- and have a charge
-when forming ions, atoms try to achieve e- config
of closest noble gas
Ex- write e- config of a fluoride ion
F1- = 1s22s22p6  e- config of neon
Try an aluminum ion
Aℓ3+ = 1s22s22p6  e- config of neon
-transition metal cations lose e- in a different
way
-the 4s will lose its e- before the 3d even though
the 4s is in a higher energy level
Ex- vanadium ion = V2+
1s22s22p63s23p63d3
paramagnetic- when atoms or ions have
unpaired e- in their e- configs (have an s, p, d
or f orbital only partially filled)
diamagnetic- when all e- are paired in an atom
or ion’s e- config (all orbitals contain max
amount of e-)
Write e- configs and orbital diagrams for the
following ions and determine if they are
diamagnetic or paramagnetic.
1) Ga3+
1s22s22p63s23p63d10
diamagnetic
2) S21s22s22p63s23p6
diamagnetic
3) Fe3+
1s22s22p63s23p63d5
paramagnetic
Trends Continued:
2) Ionic Size
-cations are much smaller than their
corresponding atoms
-anions are much larger than their corresponding
atoms
**as you move across a period the ionic size will
decrease (comparing cations to cations or
anions to anions)
**as you move down a group ionic size will
increase
1) Choose the larger ion or atom
a) S or S2- b) Ca or Ca2+ c) Br or Bra) S2b) Ca
c) Br2) Arrange in order of increasing ionic size:
Ca, Sr, Be, Mg, Ba (all have 2+ charge)
Be, Mg, Ca, Sr, Ba
Trends Continued:
3) Metallic Character
-as you move down a group, metallic character
increases
-as you move across a period, metallic character
decreases
-makes sense with distribution of metals on the
periodic table
Examples:
Choose the more metallic element:
1) tin/tellurium
2) phosphorus/antimony
3) germanium/indium4) sulfur/bromine
1) Sn
2) Sb
3) In
4) cannot tell
Arrange in order of increasing metallic character:
silicon, chlorine, sodium, rubidium
Cℓ, Si, Na, Rb
Trends Continued:
4) Ionization Energy (IE)
-the energy required to remove an e- from the
atom or ion in the gaseous state
first IE- energy needed to remove the first e-IE tends to decrease as you move down a family
b/c e- in the outermost energy level become
farther away from the + charged nucleus and
are held less tightly
-IE tends to increase as you move across a period
b/c valence e- experience greater attraction
with the nucleus
Examples:
Choose the element with the higher first IE.
1) aluminum/sulfur 2) arsenic/antimony
3) nitrogen/silicon 4) oxygen/chlorine
5) tin/iodine
6) carbon/phosphorus
1) S
2) As
3) N
4) cannot tell 5) I
6) cannot tell
Put in order of decreasing first IE:
sulfur, calcium, fluorine, rubidium, silicon
F, S, Si, Ca ,Rb
second ionization energy- energy needed to
remove the second electron
**Read page 347 on second and successive IE’s
Electron affinity (EA)
- energy change associated with the gaining
of an e- by an atom in the gaseous state
**Read pages 352-356