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Anomalies
Cr: 4s13d5 Mo: 5s14d2 Cu: 4s13d10 Ag: 5s13d10 Au: 6s13d10
Categories of electrons
Inner (core) electrons are those seen in the previous noble
gas and any completed transition series. They fill all the lower
energy levels of an atom.
Outer electrons are those in the highest energy level (highest
n value). They spend more of their time farthest from the
nucleus.
Valence electrons are those involved in forming compounds.
Among the main-group elements, the valence electrons are the
outer electrons. For the transition elements, all the (n -1)d
electrons are counted among the valence electrons also, even
though the elements Fe (Z = 26 through Zn (Z = 30) use only a
few of them in bonding as we will see later towards the last
weeks of the course.
A schematic of the SCF method for obtaining the orbitals of a
polyelectronic atom.
The radial distribution of electron probability density for the sodium
atom. The shaded area represents the 10 core electrons. The radial
distributions of the 3s, 3p, and 3d orbitals are also shown.
Note the difference in the penetration effects of an electron in these
three orbitals.
Radial probability distributions for the 3d and 4s orbitals. Note that the
most probable distance of the electron from the nucleus for the 3d orbital
is less than that for the 4s orbital. However, the 4s orbital allows more
electron penetration close to the nucleus and thus is preferred over the 3d
orbital.
Periodic Trends
• Ionization Energy
• Electron Affinity
• Atomic Radius
Ionization Energy
• Energy in kJ required for the complete removal of 1 mol of electrons
from 1 mol of gaseous atoms or ion
• Many-electron atoms can lose more than one electron. The first
ionization energy (IE1) removes an outermost electron (highest sublevel) from the gaseous atom: Atom (g) → ion+ (g) + e ∆E = IE1
• The second ionization energy (IE2) removes a second electron:
Ion+ (g) → ion2+ (g) + e ∆E = IE2 (IE2 > IE1)
X(g) → X+(g) + e–
Mg → Mg+ + e–
I1 = 735 kJ/mol(
1st IE)
Mg+ → Mg2+ + e–
I2 = 1445 kJ/mol
(2nd IE)
Mg2+ → Mg3+ + e–
I3 = 7730 kJ/mol
*(3rd IE)
*Core electrons are bound much more tightly than valence electrons.
Ionization Energy
• In general, as we go across a period from left to right, the
first ionization energy increases.
• Why?
– Electrons added in the same principal quantum level
do not completely shield the increasing nuclear charge
caused by the added protons.
– Electrons in the same principal quantum level are
generally more strongly bound from left to right on the
periodic table.
Ionization Energy
• In general, as we go down a group from top to bottom,
the first ionization energy decreases.
• Why?
– The electrons being removed are, on average, farther
from the nucleus.
The Values of First Ionization Energy for the Elements in the First Six
Periods
Explain why the graph of ionization energy versus
atomic number (across a row) is not linear.
electron repulsions
Where are the exceptions?
some include from Be to B and N to O
Electron Affinity
The electron affinity (EA) is the energy change in kJ
accompanying the addition of 1 mol of electrons to 1 mol of
gaseous atoms or ions.
The first EA refers to the formation of 1 mol of mono-valent (1 –)
gaseous anions: Atom (g) + e → ion- (g) ∆E = EA1
In most cases energy is released when an e is added because it
is attracted to the atom’s nuclear charge. The second EA2 must
always be positive because energy must be absorbed to
overcome electrostatic repulsions and add an e to a negative
ion.
a) elements in group 6A and
especially in 7A have high IE and
high EA. They lose e with
difficulty but attract e strongly.
In their ionic compounds they
form negative ions.
b) Elements in groups 1A and 2A
have low IE and slightly negative
EA. They lose e readily but
attract e very weakly. In their
ionic compounds they form
positive ions.
c) Noble gases (Group 8A) have
high IE and slightly positive EA.
They tend not to lose or gain e .
Electron Affinity
X(g) + e– → X–(g)
• In general as we go across a period from left to right, the
electron affinities become more negative.
• In general electron affinity becomes more positive in
going down a group.
Defining covalent and metallic radii
The metallic radius is ½ the distance
between nuclei of adjacent atoms in a
crystal of the element.
The covalent radius is ½ the distance
between bonded nuclei in a molecule
of the element.
C; in a covalent compound, the bond
length and known covalent radii are
used to determine other radii.
Bond length of C – Cl : 177 pm
Covalent radius of Cl: 100 pm
∴ covalent radius of C: 177 – 100 = 77
pm
−
−
Trends among main-group elements
1. Changes in n: as n increases the
probability that the outer electrons spend
more time farther from the nucleus
increases and the atoms are larger.
2. Changes in Zeff : as the Zeff - the
positive charge “felt” by an e
a) Down a group, n dominates. Atomic
radius generally increases in a group from
top to bottom.
b) Across a period Zeff dominates. Atomic
radius generally decreases in a period from
left to right.
Trends among transition metals
As we move from left to right, size
shrinks through the first two or three
transition elements because of the
increasing Zeff. But from then on, the
size remains relatively constant
because shielding by the inner d
electrons counteracts the increasing
Zeff.
The Periodic Table – Final Thoughts
1. It is the number and type of valence electrons that
primarily determine an atom’s chemistry.
2. Electron configurations can be determined from the
organization of the periodic table.
3. Certain groups in the periodic table have special names.
The trend in acid-base behavior of element oxides
•
As the elements become more metallic down a group, the oxides
become more basic (blue)
•
As the elements become less metallic across a period, their oxides
become more acidic (red)
Sb2O5: weakly basic, SiO2, As2O3: weakly acidic
Al2O3: amphoteric
The Period 4 Crossover in Sublevel Energies
For main-group, s-block metals
remove all electrons with the
highest n value , p-block metals
remove np electrons before ns
For transition metals (d-block)
remove ns electrons before (n-1)d
electrons
For nonmetals add electrons to the
p orbitals of the highest value
−
Magnetic Properties
A species with unpaired electron exhibits
paramagnetism, attracted by an external magnetic field.
A species with all electrons paired exhibits
diamagnetism, it is not attracted and, in fact, is slightly
repelled by a magnetic field.
Ionic radius
Ionic vs. atomic radius
The Alkali Metals
• Li, Na, K, Rb, Cs, and Fr
– Most chemically reactive of the metals
 React with nonmetals to form ionic solids
– Going down group:




Ionization energy decreases
Atomic radius increases
Density increases
Melting and boiling points smoothly decrease