Download Chapter 8 - Foothill College

Electron Configuration and Periodic Properties
Key Concepts/Goals:
1.
2.
3.
4.
5.
6.
Understand the “4th” quantum number, ms.
Understand and be able to describe the Pauli exclusion principle.
Understand and apply the Aufbau principle and Hund’s rule for electron filling of energy levels.
Be able to write orbital box diagrams and electron configurations for elements that follow the
expected filling order. Know the exceptions for the first row d elements.
Understand and be able to predict and explain trends in effective nuclear charge, Zeff.
Understand and be able to predict and explain the periodic trends in: atomic radii, ion radii,
ionization energy, electron affinity and properties of elements (metals, nonmetals, groups).
Periodic Table: first proposed in 1869 separately
by Dmitri Mendeleev in Russia and Lothar Meyer
in Germany. The Periodic Table proposed by
Mendeleev and Meyer was arranged in order of
increasing atomic weight. Some elements seemed
“out of order” though. The modern period table is
arranged by rows and columns in order of
increasing ATOMIC NUMBER. The properties of
the elements tend to repeat, are periodic, from row
to row.
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Electron Spin - Fourth Quantum Number, ms
An electron in an atom has the magnetic properties expected for a spinning, charged
particle. Electrons in effect behave as tiny electromagnets. Experiments have shown that
relative to an applied magnetic field, only two orientations are possible for the magnetic
moment associated with this “electron spin”: aligned with the field or opposed to the field.
This gives us a fourth quantum number, ms to define the final state of an electron.
ms is the electron spin magnetic quantum number.
ms is assigned a value of either +1/2 or -1/2 (We signify these as “spin up” or “spin down”.)
Evidence for the magnetic behavior of electrons:
+1/2
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-1/2
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Electron Spin - Fourth Quantum Number, ms
We now have 4 quantum numbers to express the energy, location and “spin” of an
electron. The Pauli Exclusion Principle states that no two electrons in an atom
can have the same four quantum numbers.
1.
2.
3.
4.
n = principle (orbital energy, size)
l = angular momentum (orbital type, shape)
ml = magnetic (specific orbital orientation)
ms = electron spin (direction of electron spin, ±1/2)
This leads to the conclusion that any atomic orbital defined by (n, l, ml) can contain
a maximum of two electrons.
Experiments have shown that when two electrons occupy the same orbital, they
have opposite spin orientations. Their electron spins are said to be “paired”, that
means that the magnetic field of one electron is cancelled by the magnetic field of
the second of opposite spin.
Example: An electron in a 3d orbital might have the following four quantum
numbers, n = 3, l = 2, ml = 0, ms = +1/2.
A second electron in the same 3d orbital would have the following quantum
numbers, n = 3, l = 2, ml = 0, ms = -1/2.
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Subshell Energies in Multi-Electron Systems Compared to Hydrogen
Hydrogen atom, single electron
Multi electron atom
Subshell energies different
in multi-electron systems
than in the H atom:
• Electron-electron
interactions complicate
the analysis.
• Fortunately, we can still
describe the electronic
structure of many electron
atoms in terms of orbitals
like those of the hydrogen
atom.
For multi-electron systems the energy of each orbital decreases (more negative) compared to the orbitals of
hydrogen. Increased nuclear charge causes this: For example, the 1s orbital of H is higher in energy than 1s
orbital of He. (Why do you think this occurs?)
Different subshells within each principle energy level (n) no longer have the same energy. For a given n:
s electrons closer than p electrons, s orbital lower in energy more stable.
p subshells lower in energy than d subshells.
d subshells lower in energy than f subshells.
These effects can be explained by the concept effective nuclear charge.
Effective Nuclear charge: The nuclear charge experienced by a particular electron in a multielectron atom, as
modified by the presence of other electrons.
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Energy Level Diagrams: Hydrogen Compared to Systems with More than One Electron
Do the energy level diagrams explain the differences in the number of spectral lines observed?
Emission Spectrum of Hydrogen
Hydrogen atom, single electron
Multi electron atom
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Effective Nuclear Charge, Zeff
The splitting of the principle energy level into the s, p, d, and f energy sublevels is best
explained by using the concept of “effective” nuclear charge, Zeff.
An electron in a higher energy level is “screened” from seeing 100% (all the protons) of the
nuclear charge by the electrons in lower energy levels. We usually talk about the valence
electrons and how they are screened from experiencing the complete nuclear charge. This
screening depends on the sublevel (orbital type) that the electron being screened is
occupying.
The Effective Nuclear Charge is the NET
NUCLEAR charge an electron experiences
when other electrons “screen” the nuclear
charge. An analogy is looking at a lightbulb that
is covered by a frosted-glass lamp shade. The
lampshade “screens” our eyes from the full
brightness of the lightbulb.
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Zeff - Effective Nuclear Charge
Sodium valence electron Zeff at 100% screening: Zeff = 11-10 = +1
Actual Na atom 3s valence electron
screening by core electrons:
Zeff = 11-8.49 = +2.51
Lower energy (inner) electrons “shield” higher
energy (outer) electrons from seeing a full
nuclear charge. This screening is not 100%.
Zeff = Z - S
where Z is the atomic number (number of
protons) and S is the screening constant.
S is a positive number with a value
that is dependent on the energy
subshell occupied by the electron.
Electrons in the same valence shell screen each
other very little, but do have a slight screening
effect. For valance (highest energy, outermost)
electrons, the core electrons provide most of the
shielding.
Screening electron density
from the core electrons:
1s, 2s, and 2p
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Trends in Effective Nuclear Charge
Graph showing the variations in effective nuclear charge
for outer electrons in period 1 through 3 elements.
Period 3
Period 2
Period 1
Questions:
•What do you notice about Zeff for hydrogen compared to atomic
number that is different compared to the other elements? What is
the reason for this difference?
•What is the trend in Zeff across a period? Can you explain the
trend?
•What happens to Zeff at the beginning of a period? Can you explain
this?
•How does Zeff for the 1s electron in K compare to its atomic
number? Can you explain why they are close in value, but not
equal?
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Other Details: “Splitting” of Subshell Energies with the Same n Value
Remember that for many electron atoms, the energies of orbitals with the same n value
increase in the order ns < np < nd < nf.
Graph showing the 1s, 2s and 2p radial
probability functions.
This can be explained by the following:
• In general, for a given n value:
s electrons penetrate closer to the nucleus than p
p electrons penetrate closer to the nucleus than d
d electrons penetrate closer to the nucleus than f
•
Thus, for a given n value, the attraction
between the the electron and the
nucleus decreases in the order:
ns > np > nd > nf
•
One result is that the ns orbitals are
lower in energy then the (n-1)d
orbitals.This is why we fill the 4s before Orbital shape causes electrons in some
the 3d, 5s before 4d, etc.
orbitals to “penetrate” close to the nucleus.
Penetration increases nuclear attraction and
decreases shielding.
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Multi Electron Atoms: Writing Electron Configurations
“Aufbau” or “building up” principle.
Electron configurations for each element are built upon the previous element in the table.
One electron at a time is placed in the lowest energy orbital available.
This gives the ground state electron configuration.
The order of orbital filling is given by the periodic table with some exceptions!
To write electron configurations using spdf notation, read the periodic table from left to right, row by
row, filling the orbitals from the lowest energy up! Its that easy!
Hund’s Rule (Bus seat rule):
For orbitals of the same energy (degenerate orbitals), the lowest energy is attained when the number
of electrons with the same spin is maximized. In other words, electrons occupy orbitals of the same
energy one at a time until the sub-shell is half-filled (Can you think of the reason why this should
occur?) In addition, all unpaired electrons will all have the same spin (parallel spins).
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Orbital Filling and the Periodic Table
The order in which the orbitals are
filled can be obtained directly from
the periodic table.
• Write condensed electron configurations and valence level orbital diagram for:
C, K, Fe, As, and Cd.
• Some exceptions to the rule for transition elements: Cu, Cr.
(Know exceptions for first row transition elements!)
• On exams, you will not be asked to write electron configurations beyond Barium.
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Ground State Electron Configurations
Some of the exceptions to
the expected filling order are
outlined in red.
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General Electron Configurations of Groups of Elements
Categories of Electrons
• Inner (core) electrons are those an atom has in common with the previous noble gas and any
•
•
completed transition series.
Outer electrons are those in the highest energy level (highest n value).
Valence electrons are those involved in forming compounds.
- For main group elements, the valence electrons are the outer electrons.
- For transition elements, the valence electrons include the outer electrons and any (n -1)d
electrons.
• The use of noble gas notation is useful for clearly showing the valence electrons. The valence
electrons determine the chemical properties of an element. Why is this?
You must familiar with the following terminology and general electron configurations of groups:
• Main Group Elements (Representative Elements): s and p block
• Groups 1A (alkali metals) and 2A (alkaline earth metals) are the s-block elements. The last electrons
are added to s orbitals. They have the general valence electron configuration: nsx
• Groups 3A through 8A are the p-block elements. The last electrons are added to p orbitals. They
have the general valence electron configuration: ns2npx
• For the main group elements, the number of valence electrons equals the group number.
• Transition Elements: For transition elements, the ns subshell fills before the (n-1)d subshell. The
transition elements are the d-block elements. The last electrons are added to d orbitals. The
transition elements have similar behavior because the electrons are being added to the (n-1) energy
level, not the n level.
• Lanthanides and Actinides: the 4f and 5f sublevels are being filled, respectively.
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Defining Atomic Radius
When two of the same atoms bond covalently, the covalent bond length (the
distance between the two nuclei) can be used to determine the covalent bonding
atomic radius of the atom. Bonding atomic radii are shorter than nonbonding
atomic radii due to the attractive forces that lead to the bond.
The covalent radius of
chlorine.
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Known covalent radii and distances
between nuclei can be used to find
unknown radii.
The metallic radius of
aluminum.
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Atomic Radius Trends (Outer Valence e–)
Atomic radii decrease along a row.
Why? Zeff increases as we add electrons to the same energy level. The increase in nuclear
charge as we move across a row is not completely screened by the additional valence
electrons so Zeff becomes larger for each valence electron. (Atomic radii of transition metals
decrease only slightly across a period.)
Atomic radii increase down a column.
Why? As we move down a column n increases for the valence electrons, hence the orbital size
also increases. Zeff also increases SLIGHTLY, but the valence electrons spend more time further
from the nucleus in the larger orbitals, 2s compared to 1s, etc.
Periodicity of Atomic Radius
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Periodic Trends - Ionization Energy, IE
Ionization Energy, IE, is the energy needed to remove an outer electron from an atom or ion
in the gas phase. Ionization energies are usually given in units of kJ/mole of electrons.
Do atoms with a low IE tend to form cations or anions?
Each atom can have a series of ionizations to produce a multi-charged cation.
For example consider the ionization of Mg(g):
1.
First:
Mg(g) —> Mg+(g) + e– IE1 = +738 kJ/mol
2.
Second: Mg+(g) —> Mg2+(g) + e–
IE2 = +1451 kJ/mol
Third: Mg2+(g)
IE3 = +7733 kJ/mol
3.
—>
Mg3+(g)
+
e–
1.
Why the increase from IE1 to IE2?
2.
Why the HUGE increase from IE2 to IE3?
Name the Period 3 element with the following ionization
energies (in kJ/mol) and write its electron configuration:
IE1
IE2
IE3
IE4
IE5
IE6
1012
1903
2910
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4956
6278
22,230
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Periodic Trends - First Ionization Energy, IE1
In general:
first ionization energy increases across
a row.
Zeff increases across a row. As Zeff
increases there is more attraction
of the electrons to the nucleus
thus more difficult to remove.
In general:
first ionization energy decreases down
a column.
The outer electrons are in higher
principle quantum shells and are
further from the nucleus. Less
attraction to the nucleus thus
easier to remove.
We see some exceptions however. For example, IE1 of N is greater than IE1 of O.
Why?
Half-filled p-sublevel for N is more stable than the partially filled p-sublevel for
O. In N, we have no e– - e– repulsive pairing energy since all p-orbitals have
only 1 e–. In O we have a p-orbital with two electrons, the pairing energy in this
p-orbital leads to a slightly less stable electron configuration and thus lower
ionization energy.
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Periodic Trends - Electron Affinity, EA
Electron affinity is the energy change when an electron is added to a neutral atom or ion in
the gas phase. Electron affinities are usually given in units of kJ/mole of electrons.
Do atoms with a high EA tend to form cations or anions?
For example:
F(g) + e– —> F–(g)
N(g) + e– —> N–(g)
EA = -328 kJ/mol
EA = + 7 kJ/mol
Questions:
All second electron affinities are positive. For example:
O-(g) + e– —> O2–(g)
EA2 = +710 kJ/mol
Does this make sense?
Why does O3– not exist in compounds in nature?
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Periodic Trends - Electron Affinity, EA
The trends are not as “regular” as for ionization energies.
Electron affinities of the main-group elements (in kJ/mol).
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Summary of Atomic Trends
Increase
Increase
Zeff
The individual atomic properties of atoms can be related to the observed macroscopic behavior of the elements.
The trends we observe across the periodic table help explain chemical behavior:
• Reactive nonmetals have high IEs and highly negative EA; they tend to form (-) ions.
• Reactive metals have low IEs and slightly negative EA; they tend to form (+) ions.
• Noble gases have very high IEs and slightly positive EA; they tend to neither lose or gain electrons.
• Compounds formed by a metal and a nonmetal tend to be ionic substances.
• Compounds formed by two nonmetals tend to be molecular substances.
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Atomic Properties of Metals and Nonmetals & Reactivity
(We will skip Redox Behavior of Main-Group Elements and Acid-Base Behavior of Oxides for now - page 345)
Some Properties of Metals:
1. Low ionization energies - tend to be oxidized
(lose electrons)
2. Form ionic compounds with non-metals
3. Metallic bonding in elemental form
Trends in Metallic Behavior
Metallic behavior
decreases across a
period.
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Metallic
behavior
increases
down a
group
Some Properties of Nonmetals:
1. Vary greatly in appearance
2. High electron affinities - tend to be reduced
(gain electrons).
3. Form ionic compounds with metals
4. Compounds of nonmetals are typically
molecular substances (covalent bonding)
Can you give an example of:
• an ionic compound formed by the elements
shown here?
• a covalent compound formed by the elements
shown here?
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Electron Configurations of Ions
Main Group Elements: electrons are lost or gained so that the electron configuration of the
ion matches that of the nearest Noble Gas. The ion is said to be isoelectronic (same
electron configuration) with the Noble Gas.
Metals lose electrons to become cations. Nonmetals gain electrons to become anions. We
can use spdf notation to show this.
1. Al —> Al3+ + 3e–
[Ne]3s23p1 —> [Ne] + 3e– (loses the 3s and 3p electrons)
2. Ca —> Ca2+ + 2e–
[Ar]4s2 —> [Ar] + 2e– (loses the 4s electrons)
3. O + 2e– —> O2[He]2s22p4 + 2e– —> [Ne] (gains two e– to fill the 2p shell)
In general, what type of orbitals are filled when nonmetals gain electrons?
Transition metals (d-block) lose the (n+1)s electrons first!
1. Fe —> Fe2+ + 2e–
[Ar]4s23d6 —> [Ar]3d6 + 2e– (loses the 4s electrons)
2. Fe —> Fe3+ + 3e–
[Ar]4s23d6 —> [Ar]3d5 + 3e– (loses the 4s & a 3d electron)
Write electron configurations for Li+, Zn2+, Mn4+, P3–, Sn2+ and Sn4+.
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Ionic Size Trends
Ions show a trend in ionic size:
• Cations are smaller than the atoms
they come due to an increase in
Zeff. Also, electron-electron
repulsions are reduced.
• Anions are larger than the atoms
they come from because of
increased electron-electron
repulsions. Also, Zeff decreases for
added valence electrons.
about 2x size
about 1/2
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Periodic Trends - Comparison of Atomic and Ionic Radii
(Units are picometers.)
Grey: Cation radius
Blue: Atomic radius
Grey: Anion radius
Red: Atomic radius
An isoelectronic series
Questions:
• What is the trend in ionic
radius down a group, as n
increases?
• What is an isoelectronic
series?
• Within a isoelectronic
series, what trend do you
notice for ion size? Does
this trend make sense?
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Magnetic Properties
A species with one or more unpaired electrons exhibits paramagnetism – it is attracted
by a magnetic field.
A species with all its electrons paired exhibits diamagnetism – it is not attracted (and is
slightly repelled) by a magnetic field.
The apparent mass of a
diamagnetic substance is
unaffected by a magnetic
field.
Determine if each of the following will have
a mass that is unaffected by a magnetic field
(Provide reasoning for each):
• Fe
The apparent mass of a
paramagnetic substance
increases because it is
attracted by the magnetic
field.
• ZnCl2
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Questions
• Compare B, Al, and C
Which has the largest atomic radii?
Which has the highest electron affinity?
Rank them in order of INCREASING first ionization energy.
Which has the most metallic Character?
• Which experiences the greatest effective nuclear charge, a 2p electron in F−, a 2p electron
in Ne, or a 2p electron in Na+?
• Consider S, Cl and K and their most common ions.
(a) List the atoms in order of increasing size.
(b) List the ions in order of increasing size.
(c) Explain any differences in the orders of the atomic and ionic sizes.
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Questions
•
Consider the first ionization energy of neon and the electron affinity of fluorine.
(a) Write the balanced chemical equations for each process. Include phase labels.
(b) These two quantities will have opposite signs. Which will be positive and which will
be negative?
(c) Would you expect the magnitudes of these two quantities to be the same. If not,
which one would you expect to be larger and why?
•
While the electron affinity of bromine is a negative quantity, it is positive for Kr. Use the
electron configurations of the two elements to account for this observation.
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