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The highlights are:
1. Pauli's Exclusion principle
2. Aufbau principle
3. Hund's Rule
4. Electronic configurations and the periodic table
1.
3.9 Electron Configurations
The electron configuration describes how the electrons are
distributed in the various atomic orbitals.
In a ground state hydrogen atom, the electron is found in the 1s
Ground state electron
configuration of
principal (n = 1) hydrogen number of electrons in
1s1 the orbital or subshell
2s 2p 2p 2p
angular momentum (l = 0)
The use of an up arrow indicates an electron
with ms = + ½
1s
2.
Electron Configurations
If hydrogen’s electron is found in a higher energy orbital, the atom
is in an excited state.
A possible excited state electron
configuration of hydrogen
2s1
2s 2p 2p 2p
1s
3.
Electron Configurations
The helium emission spectrum is more complex than the hydrogen
There are more possible energy transitions in a helium atom
because helium has two electrons.
4.
Electron Configurations
In a multi-electron atoms, the energies of the atomic orbitals are
Splitting of energy levels refers to
the splitting of a shell (n=3) into
subshells of different energies
(3s, 3p, 3d)
Text practice: 3.90
5.
Electron Configurations
According to the Pauli exclusion principle, no two electrons in an
atom can have the same four quantum numbers.
The ground state electron
configuration of helium
2p 2p 2p
2s
1s2
Quantum number
Principal (n)
1 1
1s Angular moment (l)
describes the 1s orbital 0 0
Magnetic (ml)
0 0
describes the electrons in the 1s orbital Electron spin (ms) +½ ‒½
6.
Electron Configurations
The Aufbau principle states that electrons are added to the lowest
energy orbitals first before moving to higher energy orbitals.
Li has a total of 3 electrons
The ground state electron
configuration of Li
2p 2p 2p 1s22s1
2s The third electron must go in the
next available orbital with the
1s lowest possible energy.
The 1s orbital can only accommodate 2
electrons (Pauli exclusion principle)
7.
Electron Configurations
The Aufbau principle states that electrons are added to the lowest
energy orbitals first before moving to higher energy orbitals.
Be has a total of 4 electrons
The ground state electron
configuration of Be
2p 2p 2p
1s22s2
2s
1s
8.
Electron Configurations
The Aufbau principle states that electrons are added to the lowest
energy orbitals first before moving to higher energy orbitals.
B has a total of 5 electrons
The ground state electron
configuration of B
2p 2p 2p
1s22s22p1
2s
1s
9.
Electron Configurations
According to Hund’s rule, the most stable arrangement of
electrons is the one in which the number of electrons with the same
spin is maximized.
C has a total of 6 electrons The ground state electron
configuration of C
1s 2s 2p
2 2 2
2p 2p 2p
2s The 2p orbitals are of equal energy, or degenerate.
1s Put 1 electron in each before pairing (Hund’s rule).
10.
Electron Configurations
According to Hund’s rule, the most stable arrangement of
electrons is the one in which the number of electrons with the same
spin is maximized.
N has a total of 7 electrons The ground state electron
configuration of N
1s 2s 2p
2 2 3
2p 2p 2p
2s The 2p orbitals are of equal energy, or degenerate.
1s Put 1 electron in each before pairing (Hund’s rule).
11.
Electron Configurations
According to Hund’s rule, the most stable arrangement of
electrons is the one in which the number of electrons with the same
spin is maximized.
O has a total of 8 electrons The ground state electron
configuration of O
1s 2s 2p
2 2 4
2p 2p 2p
2s
Once all the 2p orbitals are singly occupied, additional
electrons will have to pair with those already in the
1s orbitals.
12.
Electron Configurations
According to Hund’s rule, the most stable arrangement of
electrons is the one in which the number of electrons with the same
spin is maximized.
F has a total of 9 electrons The ground state electron
configuration of F
1s22s22p5
2p 2p 2p
2s
When there are one or more unpaired electrons, as
1s in the case of oxygen and fluorine, the atom is
called paramagnetic.
13.
Electron Configurations
According to Hund’s rule, the most stable arrangement of
electrons is the one in which the number of electrons with the same
spin is maximized.
Ne has a total of 10 electrons The ground state electron
configuration of Ne
1s22s22p6
2p 2p 2p
2s
When all of the electrons in an atom are paired, as
1s in neon, it is called diamagnetic.
14.
Electron Configurations
General rules for writing electron
1) Electrons will reside in the available
orbitals of the lowest possible energy.
2) Each orbital can accommodate a
maximum of two electrons.
3) Electrons will not pair in degenerate
orbitals if an empty orbital is available.
4) Orbitals will fill in the order indicated
in the figure.
15.
Worked Example 3.10
Write the electron configuration and give the orbital diagram of a calcium (Ca)
atom (Z = 20).
Setup Because Z = 20, Ca has 20 electrons. They will
fill in according to the diagram at right. Each s subshell
can contain a maximum of two electrons, whereas each p
subshell can contain a maximum of six electrons.
16.
3.10 Electron Configurations and the Periodic Table
The electron configurations of all elements except hydrogen and
helium can be represented using a noble gas core.
The electron configuration of potassium (Z = 19) is
Because 1s22s22p63s23p6 is the electron configuration of argon, we
can simplify potassium’s to [Ar]4s1.
The ground state electron configuration of K:
1s22s22p63s23p64s1
[Ar] [Ar]4s1
17.
Electron Configurations and the Periodic Table
Elements in Group 3B through Group 1B are the transition metals.
18.
Electron Configurations and the Periodic Table
Following lanthanum (La), there is a gap where the lanthanide
(rare earth) series belongs.
19.
Electron Configurations and the Periodic Table
After actinum (Ac) comes the actinide series.
20.
Electron Configurations and the Periodic Table
21.
Electron Configurations and the Periodic Table
There are several notable exceptions to the order of electron filling
for some of the transition metals.
Chromium (Z = 24) is [Ar]4s13d5 and not [Ar]4s23d4 as expected.
Copper (Z = 29) is [Ar]4s13d10 and not [Ar]4s23d9 as expected.
The reason for these anomalies is the slightly greater stability of d
subshells that are either half-filled (d5) or completely filled (d10).
Cr [Ar]
4s 3d 3d 3d 3d 3d
Greater stability with half-filled
3d subshell
22.
Electron Configurations and the Periodic Table
There are several notable exceptions to the order of electron filling
for some of the transition metals.
Chromium (Z = 24) is [Ar]4s13d5 and not [Ar]4s23d4 as expected.
Copper (Z = 29) is [Ar]4s13d10 and not [Ar]4s23d9 as expected.
The reason for these anomalies is the slightly greater stability of d
subshells that are either half-filled (d5) or completely filled (d10).
Cu [Ar]
4s 3d 3d 3d 3d 3d
Greater stability with filled 3d
subshell
23.
Worked Example 3.11
Write the electron configuration for an arsenic atom (Z = 33) in the ground state.
Setup The noble gas core for As is [Ar], where Z = 18 for Ar.
The order of filling beyond the noble gas core is 4s, 3d, and 4p. Fifteen electrons
go into these subshells because there are 33 – 18 = 15 electrons in As beyond its
noble gas core.
Text Practice 3.93 3.98 3.104a, b, c, d 3.106 3.113
24.
Study Guide for sections 3.9-3.10
DAY 7, Terms to know:
Sections 3.9-3.10 electron configuration, Aufbau principle, Hund’s rule, orbital
DAY 7, Specific outcomes and skills that may be tested on exam 1:
Sections 3.9-3.10
•Be able to use the Pauli exclusion principle and Aufbau principle to give a
complete or abbreviated electron configuration for an atom in either its ground state
or one possible excited state
•Given an electron configuration, be able to give a complete elemental symbol for
an atom
•Be able to use the Pauli exclusion principle, Aufbau principle, and Hund’s rule to
give a complete or abbreviated orbital diagram for an atom either its ground state
or one possible excited state
•Given an orbital diagram, be able to give a complete elemental symbol for an atom
or ion
•Be able to recognize and explain how Cr and Cu are exceptions to the Aufbau
25.
Extra Practice Problems for sections 3.9-3.10
Complete these problems outside of class until you are confident you have learned
the SKILLS in this section outlined on the study guide and we will review some of
them next class period. 3.95 3.97 3.99 3.101 3.103 3.105 3.115 3.117 3.119
3.137 3.139
26.
On day 8, we will have exam 1
Prep for day 9
Must watch videos:
http://www.learnerstv.com/video/Free-video-Lecture-3354-Chemistry.htm (ionization energy)
http://www.learnerstv.com/video/Free-video-Lecture-3355-Chemistry.htm (periodic trends)
Other helpful videos:
lecture-9/ (MIT)
http://ps.uci.edu/content/chem-1a-general-chemistry (UC-Irvine lectures 7)
Read sections 4.1-4.6
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