Download Chapter 8 - Atomic Electron Configuration and Chemical Periodicity

Chapter 8 - Atomic Electron Configuration and Chemical Periodicity
Atomic orbitals are described by ________ quantum numbers
Schrodinger equation cannot be solved exactly for multi-electron atoms. The
approximate solutions show that the atomic orbitals of multi-electron atoms are
_______________ and we can use the same quantum numbers to describe them.
But the things to consider are
(a) a limit on the number of electrons that are allowed in an orbital, which
requires the introduction of a ___________ quantum number and
(b) a more complex set of orbital energy levels.
Electron Spin and Magnetism
Electron has a spin just as the earth has a spin. The electron spin is represented by
the fourth quantum number called the _________________________ quantum
number, ms. It can have one of two possible values, ____, ____. Hence the
complete description of an electron in an atom requires four quantum numbers (n,
l, ml, and ms).
Most substances are slightly repelled by a strong magnet; they are called
____________.
Some metals and compounds are attracted to a magnetic field; they are called
___________________.
When an electron is assigned to an orbital in an atom, its spin orientation can take a
value of + ½ or - ½.
Atoms with a single electron in their outer shell are paramagnetic; when an
external magnetic field is applied, the electrons align with the field an experience
an attractive force. But the ones with two electrons are diamagnetic. The two
electrons have opposite spin orientations (spins are paired).
Substances which have many unpaired electrons that have all had their spins
aligned are called ________________________. They can be used to make
magnets since they do not require an external magnetic field to keep the electron
spins aligned (iron, cobalt, nickel are few examples).
The Pauli Exclusion Principle
Austrian Physicist Wolfgang Pauli formulated the exclusion principle which states
that
“_________________________________________________________________
_____ __________________”.
Electron in
ground state of
H atom
1st e- of He atom
2nd e- of He atom
3rd e- of Li atom
5th e- of B atom
n
Values for
l
ml
ms
Each orbital can hold only two electrons. Hence the maximum number of electrons
in each shell is ____.
Atomic subshell Energies and Assignments
The electrons are arranged in such a way that the total energy of the atom is as low
as possible. Bohr’s model of the atom states
that the energy of single electron atoms
depends only on n (E =___________). The
situation is more complex for multi-electron
atoms. From experimental observations it
has been shown that the
_______________________________
______________________.
- electrons are assigned to subshells in
increasing order of “______” values
- when two subshell has the same
“_____” values, then the electrons are
assigned first to the subshell of lower
___.
This is explained by aufbau principle. It is often illustrated as shown in the figure.
The orbital can be represented by a box and the electrons by an arrow (↑ represents
+ ½ and ↓ represents – ½)
Two electrons in 1s orbital would be represented as
This is called the orbital box diagrams (or box notation). This can be written in
spdf notation as ____.
Maximum no. of electrons in s orbital is ___
Maximum no. of electrons in p orbital is ___
Maximum no. of electrons in d orbital is ___
Maximum no. of electrons in f orbital is ___
Fill the electrons in the boxes (orbitals) for the
first 4 elements of the periodic table
Corresponding electron configuration in
spdf notation
Each successive electron is assigned to a different orbital of the subshell until the
subshell is half full. Additional electrons must then be added to each orbital in the
subshell as a second electron. This assignment follows the Hund’s rule which
minimizes the total energy of an atom:
Draw orbital box diagram for C and N following Hund’s rule
Write the electron configuration of
Ti
Pb
Electron configuration for atoms with higher atomic numbers is often lengthy;
hence they are written in abbreviated form by combining the noble gas notation
with the spdf notation. The electrons included in the noble gas configuration are
often referred to as the ___________ ____________ and the electrons beyond the
core electrons are called _________ ___________, which determine the chemical
properties of an element.
Write the abbreviated electron configuration of
Ti
Pb
Electron configuration and the periodic table
Conveniently, it turns out that the periodic table is arranged according to electron
configuration.
1s1
1s2
2s1 2s2
2p1 2p2 2p3 2p4 2p5 2p6
3s1 3s2
3p1 3p2 3p3 3p4 3p5 3p6
4s1 4s2 3d1 3d2 3d3
3d5 3d6 3d7 3d8
3d10 4p1 4p2 4p3 4p4 4p5 4p6
5s1 5s2 4d1 4d2
4d5
4d10 5p1 5p2 5p3 5p4 5p5 5p6
6s1 6s2 5d1 5d2 5d3 5d4 5d5 5d6 5d7
5d10 6p1 6p2 6p3 6p4 6p5 6p6
7s1 7s2 6d1 6d2 6d3 6d4 6d5 6d6 6d7 6d8 6d9 6d10
4f 3 4f 4 4f 5 4f 6 4f 7
4f 9 4f 10 4f 11 4f 12 4f 13 4f 14 5d1
5f 6 5f 7
5f 9 5f 10 5f 11 5f 12 5f 13 5f 14 6d1
s – block elements
d – block elements (transition metals)
p – block elements
f – block elements
Exceptions to Aufbau principle
________________ and ________________ are most stable. Hence atoms
sometimes break the aufbau order to attain half-filled or filled subshells.
e.g.
Cr is not [Ar]4s23d4; it is
Cu is not [Ar]4s23d9; it is
All of these exceptions are elements from either the ‘d-block’ or the ‘f-block’, so
they are either transition metals (Cr, Cu, Nb, Mo, Ru, Rh, Pd, Ag, Pt, Au),
lanthanides (Ce, Pr, Gd), or actinides (Th, Pa, U, Np, Pu, Cm).
Electronic configuration of ions
To form a cation from an atom, one or more of the ____________ electrons is
removed. The electrons are removed from the electron shell of highest ___. If there
are several subshells present within the nth shell, the electrons of maximum ___ are
removed.
For e.g.
Na: 1s22s22p63s1 → Na+: 1s22s22p6 + eAs: [Ar] 3d104s24p3 → As3+: [Ar] 3d104s2 + 3eAs: [Ar] 3d104s24p3 → As5+: [Ar] 3d10 + 5eNote that the ____ electrons are removed before the ____ ones; whereas when the
electrons are filled in the subshells, the ____ electrons gets filled first. This is due
to the fact that the 4s orbital in group 1 & 2 is lower in energy than the empty 3d
orbital. However when we get to the transition series, the 3d orbitals begin to fill
and their electrons become increasingly attracted to the greater nuclear charge. As
a result, a changeover in orbital energy takes place and the energy of 3d orbital
becomes ________ than that of 4s.
Write the electron configuration of the following atoms and ions
Li
+
as a neutral atom
as an ion
P3Ti+
Sn2+
Sn4+
Which of the following ions are likely to form? For those which are not, what ion
would you expect to form from that element?
a) O2c) Cl+
e) Pb4+
b) Mg6-
d) Ca+
f) Ga3+
Effective Nuclear Charge, Z*
Most electrons do not ‘feel’ the full positive
charge of the nucleus. Other electrons in the
atom (particularly those in lower energy
orbitals) ‘shield’ some of this charge. The
___________ __________ __________ is the
charge experienced by an electron in a multielectron atom, as modified by the presence of
other electrons. The order of electron filling
and many atomic properties can be explained
by Z*.
Consider Li atom (1s22s1): It can be
seen from the surface density plot that the 2s
electron density region penetrates the 1s region. If the 2s electron is a large
distance from the nucleus, it would experience less charge (+1) due to the shielding
by the two 1s electrons. On the other hand if it penetrates the 1s electron region it
would experience increasingly higher positive charge to a maximum of +3.
This is also the reason for 4s orbital having lower energy than the 3d orbital;
since 4s orbital has __________ node, the 4s electron spends a small but
significant amount of time close to the nucleus compared to the 3d electron which
has no _________ node.
Values of Z* for the s and p electrons of the most second-period elements are listed
in the table below.
Atomic Properties and Periodic Trends
Similarities in properties are the result of _________________________________
___________________________________________.
Atomic Radius
• _________________ between centers of the two atoms in a diatomic molecule
• Estimated from atom-atom distance in a crystal of mono-atomic element (like
metals)
• The atomic radii ____________ going down the group as more electron shells
are added
• ______________ going across the period as more electrons are added to the
same shell and the effective nuclear charge for the valence electrons increases.
Ionization energy (IE)
• Energy required to remove an electron from an atom in the gas phase.
• Energy required to remove the second valence electron is called the second
ionization energy (IE2).
• IE ______________ across the period as the effective nuclear charge increases
• IE ______________ down the group as the valence electrons are increasingly
farther from the nucleus and are more shielded from the nucleus
Electron Affinity (EA)
• Energy released in a process in which an electron is acquired by the atom in the
gas phase.
• EA values are always _____________; hence an increase in EA means more
________value.
• EA _____________ across the period (except for noble gases) as the effective
nuclear charge increases
• EA ____________ down the group as the valence electrons are increasingly
farther from the nucleus and are more shielded from the nucleus
Ionic Radii
• Periodic trends in the sizes of ions in the same group are the same as those for
_____________________.
• The radius of a cation is always _____________ than that of the neutral atom.
Removing an electron without changing the nuclear charge ____________ the
Z*
• The radius of the anion is always ______________ than the atoms from which
they are formed. Adding an electron without changing the nuclear charge
_______________ the Z* due to increased shielding.
Important Concepts from Chapter 8
• diamagnetic vs. paramagnetic vs. ferromagnetic substances
• electron spin
• Pauli exclusion principle
• drawing orbital box diagrams
• writing electron configurations
• using the aufbau order and predicting exceptions to it
• using electron configuration to predict if an ion will form
• core vs. valence electrons
• effective nuclear charge
• periodic trends (atomic radius, ionization energy, electron affinity, ionic radius)