Atomic Electron
Configurations and
Chemical Periodicity
Goals:
1. Understand the role magnetism plays in
determining and revealing atomic structure.
2. Understand effective nuclear charge and its role
in determining atomic properties.
3. Write the electron configuration for elements
and monoatomic ions.
4. Understand the fundamental physical properties
of the elements and their periodic trends.
Arrangement of Electrons
in Atoms
Electrons in atoms are arranged as:
• Shells (n)
• Subshells (l)
• Orbitals (ml)
Electrons have _____.
• ms, __________________
quantum number, = +1/2 and -1/2
Complete description of electrons
requires _______ quantum numbers.
Electron Spin Magnetic Quantum
Number
Electron Spin and Magnetism
• ____________:
NOT attracted to a
magnetic field
• ___________:
substance is attracted
to a magnetic field.
• Substances with
unpaired electrons are
______________.
Electron Spin and Magnetism
• H atoms, each has a single electron,
they are paramagnetic – when an
external magnetic field is applied, the
electron magnets align with the field.
• He atoms, with two electrons, are
diamagnetic.
– We assumed opposite spin
orientations – spins are __________.
The Pauli Exclusion Principle
No two electrons in an atom can
have the same set of four quantum
numbers.
Therefore,
Each orbital can be assigned no
more than ____ electrons!
Orbital Box Diagrams
When n = 1, then l = 0
• this shell has a single orbital (1s) to which
2e- can be assigned.
H
(1e) n=1, l =0, ml =0
ms = +1/2
1s
He (2e) n=1, l=0, ml = 0, ms = +1/2
n=1, l=0, ml = 0, ms = -1/2
1s
Electrons in Atoms
A 2p electron can be designated by which
set of quantum numbers?
a.
b.
c.
d.
e.
n
1
2
2
3
3
l
0
1
2
1
2
ml
0
0
+1
+2
+1
ms
+1/2
+1/2
-1/2
+1/2
+1/2
Students should be familiar with the values
and meaning of quantum numbers.
Electrons in Atoms
• Electrons generally assigned to orbitals
of successively higher energy.
• For H atoms, E = - C(1/n2). E depends
only on _____.
• For many-electron atoms, energy depends
on both ____ and _____.
Assigning Electrons to Subshells
• In many-electron atom:
a) subshells increase in energy
as value of _______increases.
b) for subshells of same _____,
subshell with lower n is lower
in energy.
1 e- atom
Effective Nuclear Charge, Z*
• Z* - the nuclear charge experienced
by a particular electron in a
multielectron atom, as modified by
the presence of the other electrons.
• Li has 3 p (+) and 3 e (-)
2 e in 1 s orbital ; 1 e in 2 s orbital
e- in 2s should “see” a +1 charge,
but it sees 1.28
• C has 6 p (+) and 6 e (-)
2 e- in 1s ; 2 e- in 2s ; 2 e- in 2p
e- in 2s should “see” +3, but see 3.22
e- in 2p should “see” a +2 charge, but
see 3.14
Electron
cloud for
1s
electrons
Effective Nuclear Charge, Z*
• Z* is _______ for s electrons
than for p electrons.
– s electrons always have a lower
energy than p electrons in the
same quantum shell.
• The Z* _________ across a
period.
• The 2s electron PENETRATES
the region occupied by the 1s
electron.
– 2s electron experiences a
_________ positive charge
than expected.
Atomic Electron Configurations
• The arrangements of __________ in the elements in
the ground state.
• In general, electrons are assigned to orbitals in order
of increasing ________.
• Electron configuration can be given with the orbital
box diagram, or with the spdf notation.
for H, atomic number = 1
Orbital Box notation
spdf notation
1
1s
1s
value of n
no. of
electrons
label of l
Electron Configurations
• The outermost electrons of an element are
assigned to the indicated orbitals.
Lithium
Group 1A
Atomic number = 3
________ ---> 3 total electrons
3p
3s
2p
2s
1s
Boron
Group 3A
Atomic number = 5
___________ --->
5 total electrons
3p
3s
2p
2s
1s
Carbon
Group 4A
Atomic number = 6
___________ --->
6 total electrons
3p
3s
2p
2s
1s
Here we see for the first
time ___________ RULE:
When placing electrons in a
set of orbitals having the
same energy, we place them
singly as long as possible.
Hund’s Rule
• The most stable arrangement of
electrons is that with the __________
_____________________, all with
the same ________ direction.
• This arrangement makes the total
energy of an atom as low as possible.
Nitrogen and Oxygen
Group 5A
Atomic number = 7
_________--->
7 total electrons
3s
3p
Group 6A
3s
3p
2p
2p
2s
2s
1s
1s
Fluorine and Neon
Group 7A
Atomic number = 9
____________--->
9 total electrons
Group 8A
Atomic number = 10
_____________ --->
10 total electrons
3p
3p
3s
3s
2p
2p
2s
2s
1s
1s
• Note that we have
reached the end of the
2nd period, and the
2nd shell is ________!
Sodium and Potassium –
Noble gas notation
Na - Group 1A
Atomic number = 11
1s2 2s2 2p6 3s1
or “neon core” + 3s1
[Ne] 3s1 (uses rare gas notation)
Note that we have begun a new period (3rd )
All Group 1A elements have
[core]ns1 configurations, n=period number
K – Atomic number = 19, 4th period
___________
Phosphorus
Group 5A
Atomic number = 15
spdf: ______________
short: __________
All Group 5A elements
have [core ] ns2 np3
configurations where n
is the period number.
3p
3s
2p
2s
1s
Transition Metals
3d orbitals used for ScZn (Table 8.4)
and so are d-block elements.
Transition Metals
All 4th period elements have the configuration
[argon] nsx (n - 1)dy and so are d-block elements.
Chromium
Iron
Copper
26e-
Lantanides and Actinides
4f orbitals used for Ce - Lu and
5f for Th - Lr (Table 8.2)
and so are f-block elements.
Lantanides and Actinides
All these elements have the configuration
[core] nsx (n - 1)dy (n - 2)fz and so are
f-block elements.
Cerium
[Xe] 6s2 5d1 4f1
Uranium
[Rn] 7s2 6d1 5f3
Ion Configurations
To form cations from elements remove 1 or more efrom subshell of highest n [or highest (n + l)].
P
---> P3+
[Ne] 3s2 3p3
- 3e[Ne] 3s2 3p0
3p
3p
3s
3s
2p
2p
2s
2s
1s
Ion Configurations
For transition metals, remove ns electrons and then
(n - 1) electrons.
Fe [Ar] 4s2 3d6
loses 2 electrons ---> Fe2+ [Ar] 4s0 3d6
Fe2+
Fe
4s
3d
To form cations, always
remove electrons of
highest n value first!
4s
3d
Fe3+
4s
3d
Practice
• Which substance will be paramagnetic?
V+5 or Fe+3
Students should be familiar with writing
electronic configurations and identifying
diamagnetic vs. paramagnetic materials.
Ion Size Makes a BIG Difference
• About 20% of the CO2 binds to
hemoglobin and is released in the
lungs. About 70% is converted by
Carbonic Anhydrase into HCO3- ion,
which remains in the blood plasma
until the reverse reaction releases
CO2 into the lungs.
• Carbonic Anhydrase catalyzes the
reversible hydration of CO2 to form
bicarbonate anion and a proton:
CO2 + H2O <==> HCO3- + H+
• Toxic metals like Cd2+ replace Zn2+
inactivating the enzyme.
PERIODIC
TRENDS
Periodic Trends
• Atomic and ionic size
• Ionization energy
• Electron affinity
Higher effective nuclear charge
Electrons held ______ tightly
Larger orbitals.
Electrons held ____
tightly.
Atomic size
• Size goes ____ on
going down a group.
See Figure 8.9.
• Because electrons are
added further from
the nucleus, there is
______ attraction.
• Size goes _______ on
going across a period.
Effective Nuclear Charge, Z*
• Atom
•
•
•
•
•
•
•
Li
Be
B
C
N
O
F
Z* Experienced by Electrons in
Valence Orbitals (Outermost)
+1.28
------Increase in
+2.58
Z* across a
+3.22
period
+3.85
+4.49
+5.13
[Values calculated using Slater’s Rules]
Atomic size- Transition Metals
• 3d subshell is inside the 4s subshell.
• 4s electrons feel a more or less constant Z*.
• Sizes stay about the same and chemistries are similar!
Ion Sizes
+
Li,152 pm
3e and 3p
Li + , 78 pm
2e and 3 p
Forming
a cation.
• CATIONS are __________ than the
atoms from which they come.
• The electron/proton attraction has gone
______ and so size ____________.
Ion Sizes
F, 71 pm
9e and 9p
F- , 133 pm
10 e and 9 p
Forming
an anion.
• ANIONS are __________ than the
atoms from which they come.
• The electron/proton attraction has
gone ______and so size __________.
• Trends in ion sizes are the same as
atom sizes.
Ion Sizes
Redox Reactions
Why do metals lose
electrons in their
reactions?
Why does Mg form Mg2+
ions and not Mg3+?
Why do nonmetals take
on electrons?
Ionization Energy
IE = energy required to remove an
electron from an atom in the gas
phase.
Mg (g) + 738 kJ ---> Mg+ (g) + e-
Ionization Energy
Mg (g) + 738 kJ ---> Mg+ (g) + e-
Mg+ (g) + 1451 kJ ---> Mg2+ (g) + eMg+ has 12 protons and only 11 electrons.
Therefore, IE for Mg+ > Mg.
Ionization Energy
Mg (g) + 735 kJ ---> Mg+ (g) + eMg+ (g) + 1451 kJ ---> Mg2+ (g) + e-
Mg2+ (g) + 7733 kJ ---> Mg3+ (g) + eEnergy cost is very high to dip into a shell of
lower n.
This is why oxidation number = Group number.
Effective nuclear charge
• Sodium
valance A valance electron in an atom is
attracted to the nucleus of the atom
1and it is repelled by the other
10electrons in the atom: inner e- shield
11+
or screen the outer electrons from
attraction of the nucleus.
Effect = 11-10 = +1
core
Radial electron density
Effective nuclear charge
[Ne] core
3s
For the 3s e- (valance e- of Na) there is a probability
of being found close to the nucleus – there is a
probability of experiencing a greater attraction than
suggested.
Zeff = +2.5
Electrons in 3s orbitals has a higher Zeff than 3p
orbitals: subshells energy trend is: ns < np < nd
Atomic size
Radius (pm)
250
K
Na
2nd period
Size decreases across a
period owing to increase in
Z*. Each added electron
feels a greater and greater
+ charge.
1st transition
series
3rd period
200
Li
150
Kr
100
Ar
Ne
50
He
0
0
5
10
15
20
25
30
35
Atomic Number
Large
Small
Increase in Z*
40
Ionization Energy
IE ___________ across a
period and ___________
down a group.
Ionization Energy
As Z* increases, orbital energies “drop” and IE increases.
Trends in Ionization Energy
• IE increases across a period because Z*
increases.
– Metals lose electrons more easily than nonmetals.
– Metals are good reducing agents.
– Nonmetals lose electrons with difficulty.
• IE decreases down a group .
– Because size increases.
– Reducing ability generally increases down the
periodic table (easier to give e-).
– See reactions of Li, Na, K
Ionization Energy
Electron Affinity
A few elements GAIN electrons to form
anions.
Electron affinity is the energy change
that occurs when _______________
to ___________________________.
A(g) + e- ---> A-(g)
Electron Affinity = ∆E
Electron Affinity
Electron Affinity
Affinity for electron
_________ across a
period (EA becomes
more negative).
Affinity __________
down a group (EA
becomes less
negative).
Summary
Practice
• Which of the following elements has the greater
difference between its first and second ionization
energies: C, Li, N, Be?
• Which should be smaller: the sulfide ion, S2-, or a
sulfur atom, S?
Students should be familiar with periodic
trends – IE, EA, Atomic size.
Remember
• Go over all the contents of your
textbook.
• Practice with examples and with
problems at the end of the chapter.
• Practice with OWL tutor.
• Work on your assignment for Chapter 8.
• Practice with the quiz on CD of
Chemistry Now.