Electronic Structure of Atoms(Chapter 7)
Light- Electromagnetic radiation , Consists of photons.
One photon has energy: E = h; h is Planck’s constant (6.63 x 10-34J-sec)
E = h is the energy for 1 quantum. At subatomic level, energy is quantized.
Different colors of light have different wavelengths. ( is Greek letter nu)
Properties of Light
Light spectra in visible region Balmer Series:
Violet
Blue
Green
Yellow
Orange
Red 700nm
400nm
6→2
5→2
4→2
5→3
3→2
Wavelength- Distance between two consecutive crestpoints or troughs of a wave. Units
of wavelength are in meters(m), nanometers(nm), and angstroms(Å).
Frequency(, f)- Number of wavelengths that pass a point per second. Unit is
reciprocal seconds or hertz.
Velocity- is the product of frequency and wavelength(v = ).
When v is measured in a vaccum, v = c =  = 3.00 x 108 m/s
Bohr Model Of the Atom 2πr = nλ
The Bohr Hydrogen Atom
Light- Electromagnetic radiation , Consists of photons.
One photon has energy: E = h; h is Plank’s constant(6.63 x 10-34J-sec)
E = h is the energy for 1 quantum. At subatomic level, energy is quantized.
Different colors of light have different wavelengths. ( is Greek letter nu)
1. Only certain orbits are permitted for electrons. The most inner orbit is called the
K shell, 2nd most inner is L shell, 3rd most inner is M shell, etc
2. Electron in permitted orbits do not radiate energy
3. Electron absorbs or emits light energy that is congruent with the difference in
energy between orbits.
Energy of the Hydrogen Atom:(orbits are called orbitals)
E = (-2.18 x 10-18J)(1/n2); n = 1, 2, 3…
When electrons absorb or emit light energy, the change in energy is:
E = h = hc/ = (-2.18 x 10-18J)(1/n2f - 1/n2i); nf is final state or orbital, ni is initial.
Ex. n3  n1
E = (-2.18 x 10-18J)(1/12 – 1/32) = -1.94 x 10-18J
Matter Behaves Like a Wave
At subatomic level, matter (like electrons, protons) have wave properties like
wavelengths ( The De Broglie equation).
Ex.  = h/mv; m is mass (kg) and v is velocity (m/s).
If v is 5.97 x 106m/s for an electron, m is 9.11 x 10-31kg
Then  = 1.22 x 10-10m
Heisenberg’s Uncertainty Principle says that position and speed of subatomic matter
can not be simultaneously determined: ∆x•∆mv ≥ ħ
Four Quantum Numbers of an Electron in an Atom
Each electron in an atom has 4 quantum numbers.
1. n is the principle quantum number(QN).; gives information about the energy of the
electron in an orbital; n has values from 1, 2, 3….
2. l is the azimuthal QN; Defines shape of orbital. For a given n, l has values that
range from 0 to (n-1) in increments of one.
Ex. If n = 3 , then possible values of l are 0, 1, 2. These are known as subshells.
l = 0 is called the s subshell or s orbital.
l = 1 is called the p subshell or p orbitals. There are three p orbitals.
l = 2 is called the d subshell or d orbital. There are five d orbitals.
l = 3 is called the f subshell or f orbital. There are seven f orbitals.
The maximum number of electrons an orbital can hold is 2 electrons.
3. ml is the magnetic quantum number; has values that range from –l to +l, including
0. This range indicates number of orbital for the subshell, as above.
Ex. For l = 2, ml has values (-2, -1, 0, 1, 2), which indicates five d orbitals.
4. sl is the spin quantum number, only 2 possible values (+1/2 or –1/2).
s (l=0)
m=0
m=0
s
pz
p (l=1)
m=±1
px
d (l=2)
m=±1
m=0
py
dz2
dxz
dyz
m=±2
dxy
dx2-y2
m=0
fz3
m=±1
fxz2
fyz2
f (l=3)
m=±2
fz(x2fxyz
2
y )
m=±3
fx(x2- fy(3x
2
3y )
2
y )
=1
=2
=3
=4
=5
=6
=7
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For n = 1 to n = 4 the Following Apply
1. The value of n is the number of types of subshells in the shell.
Ex. K shell; n = 1; l = 0  s orbital
L shell; n = 2; l = 0, 1  s and p orbitals (2 types of subshells).
M shell; n = 3; l = 0, 1, 2  s, p and d orbitals (3 types)
N shell; n = 4; l = 0, 1, 2, 3  s, p, d, and f orbitals (4 types)
2. For n, the total number of orbitals is n2; the maximum number of electrons is 2n2.
s orbital is spherical; has nodes. => О
p orbitals start with n = 2. The 3 orbitals are like dumbbell shape(each with 2 lobes)
with 1 node and are aligned along the 3 cartesian axes (px , py , pz) => ∞
d orbitals start with n = 3. Four of the five orbitals have 4 lobes. The fifth orbital is
dumbbell shaped with a donut-like shape around the node. ОlО
ОlО
f orbitals are for n = 4. There are seven orbitals and they have many lobes and are
complicated.
Orbital Designations: for n = 4 & l = 2 => 4d orbitals
The s-block, p-block, d-block and f-block of the periodic table.
Electronic Configuration – Shows the distribution of electrons in the orbitals of an
atom (horizontal energy level diagram; energy increases from left to right).
Ex. For Hydrogen [ ], therefore electronic configuration is 1s1
1s1 .
For Lithium [] [ ], therefore 1s2 2s1 (auf bau principle, ground state E.C.)
1s
2s
Pauli Exclusion Principle- No two electrons in an atom can have the same set of 4
quantum numbers.
Ex. For Helium [], electronic configuration is 1s2
1s2
For both electrons n = 1, l = 0, ml = 0
But sl,(spin quantum number) for the first electron(  ) is +1/2; for the second
electron(  ) it is –1/2.
Hund’s Rule- For degenerate orbitals (like p, d, & f ), electrons fill the orbitals as
unpaired(paramagnetic state) before pairing takes place(Completely paired electrons
is a diamagnetic state).
Ex. For Nitrogen [] [] [ ][ ][ ],
1s
2s
2px 2py 2pz
E.C. is 1s2 2s2 2p3
Condensed Electronic Configuration- Use nearest noble gas core as shorthand for
writing long electronic configurations. It also shows the number of valence
electrons.
Ex. For Sodium E.C. is 1s22s22p63s1 becomes [Ne]3s1 (one valence electron)
Electronic Configuration of Ions- Electrons are removed from the orbital with the
highest n value for cations OR added to next available orbitals for anions.
Ex. Zn  Zn+2 means:
[Ar]4s23d10  [Ar]3d10 (2 electrons removed)
-2
2 2
Ex. O => O means: 1s 2s 2p4 => 1s22s22p6 = [Ne] (2 electrons added)
Half-Filled Stability- some exceptions when writing electronic configuration. Some
subshells tend to have more stable electron configuration when they are half-filled.
Chromium: [Ar]4s13d5 , not [Ar]4s23d4
Copper: [Ar]4s13d10, not
[Ar]4s23d9
Electronic Configuration for Excited States- An electron or more are promoted out
of normal sequence to a higher orbital.
Ex. Ground state config. For Magnesium is [Ne]3s2
An excited state config. For Magnesium is [Ne]3s13p1
Isoelectronic Set, Paramagnetism, Diamagnetism
Isoelectronic Set – chemical species with same number of electrons.
Example( O-2, F-, Na+, Mg+2 )
Trends in the Periodic Table(Chapter 7 Pt II)
There are certain trends with respect to effective nuclear charge, size of atoms and ions,
ionization potentials, electron affinity and electronegativity in the periodic table.
Effective Nuclear charge- is the electric field that an electron feels in an atom. The
charge it feels is from the protons in the nucleus. However, this charge is deminished
because of the shielding effect that core electrons closer to the nucleus provide. As one
goes across the periodic table from left to right, the effective nuclear charge increases
while the number of CORE electrons stays the same per period.
As one goes down a family or group, the charge increases only slightly, or is nearly
constant.
Size of Atoms (atomic radius)- atomic size of atoms is measured in terms of bonding
atomic radius. The bonding atomic radius- radius of the atom when it is bonded to
another atom- is shorter than the nonbonding atomic radius. Think of an atom as a sphere.
Atomic size tends to decrease as one goes from left to right per period on the periodic
table. The bonding radii across period 2 from Li to Ne are respectively:
(1.34, 0.9, 0.82, 0.77, 0.75, 0.73, 0.71, 0.69 angstroms).
Atomic size increases as one goes from top to bottom, mainly because of the principle
quantum number of the outer electron. For alkali metals from Li to Rb bonding radii are:
(1.34, 1.54, 1.96, 2.11 angstroms).
Mono-atomic cations tend to be smaller than their parent atoms, and generally follow
the same periodic trends in size. (Ex. Na => 1.54 Å, Na+ => 0.97 Å; Mg => 1.30 Å, Mg+2
=> 0.66 Å; Al => 1.48 Å, Al3+ => 0.51 Å). The anions are generally larger than their
parent atoms (Ex. O => 0.73 Å, O2- => 1.40 Å; F => 0.71 Å, F- => 1.33 Å; S => 1.02 Å,
S2- => 1.84 Å).
Isoelectronic Series- is a set of atoms and/or ions that have the same number of
electrons.
Ex. O-2, F-, Na+, Mg2+, Al3+ (All species have 10 electrons each; they are isoelectronic.
But the proton number increases from 8, 9, 11, 12, 13, causing the ionic radii to be in
decreasing order respectively: 1.40, 1.33, 0.97, 0.66, 0.51 angstroms)
Ionization Potential(I.P.)- energy change associated with the removal of an electron
from an atom in the gaseous state, thereby making a cation. There can be successive
ionization potentials. There is usually no dramatic change in each successive ionization
potential for the removal of valence electrons. The number of valence electrons for an
atom is the same as its group number.
There is a dramatic change in ionization potentials for the removal of core electrons. Core
electrons are the other electrons (noble gas core in condensed E.C.), besides the valance
electrons.
For example, the 5th and 6th ionization potentials for the Sulfur{ [Ne]3s23p4 }atom are
respectively, 7010 kJ/mol and 8500 kJ/mol. But the 7th ionization potential (removal of
the first core electron) rises dramatically, 27,100 kJ/mol. The 1st, 2nd, and 3rd ionization
potentials for Aluminum { [Ne]3s23p1 }are respectively, 578 kJ/mol, 1820 kJ/mol &
2750 kJ/mol. But the 4th ionization potential also rises dramatically, 11,600 kJ/mol.
Ionization potentials tend to increase as one goes from left to right on the periodic table.
They tend to decrease as one goes from top to bottom.
Metals have the lowest I.P.’s, whereas nonmetals have the highest I.P.’s.
There are exceptions to the trend (Group 2 to Group 3 & Group 5 to Group 6)
Electron Affinity(E.A)- energy change associated with the addition of an electron to an
atom in the gaseous state. Almost the opposite of I.P., but not quite.
Electron affinity tends to increase from left to right on the periodic table. It tends to
decrease from top to bottom. Non metals have higher E.A. than metals in same period.
There are exceptions to the trend (Group 1 to Group 2 & Group 4 to Group 5)
Electronegativity- Tendency of an atom(in a compound or ion ) to attract electrons.
Fluorine is the most electronegative atom on the periodic table. Francium is the least
electronegative atom on the periodic table. Non metals tend to be more electronegative
than metals. Metals are more electropositive.
Electronegativity tends to increase from left to right in the periodic table. It tends to
decrease from top to bottom.
Metals
Metals are generally electropositive and have low ionization potentials, low electron
affinity, good electron conductors, lustrous, maliable . They tend to react with non metals
to form ionic compounds.
Metal oxides are considered to be basic because they react with water to form bases.
Na2O(s) + H2O(l)  2NaOH(aq)
BaO(s) + H2O(l)  Ba(OH)2(aq)
Al2O3(s) + 3H2O(l)  2Al(OH)3(aq)
Metal oxides react with acids to form the metal salt and water.
NiO(s) + 2HCl(aq)  NiCl2(aq) + H2O(l)
ZnO(s) + H2SO4(aq)  ZnSO4(aq) + H2O(l)
CaO(s) + 2HNO3(aq)  Ca(NO3)2(aq) + H2O(l)
Some metals form other oxygen compounds.
2Na(s) + O2(g)  Na2O2(s) Sodium peroxide
K(s) + O2(g)  KO2(s) Potassium superoxide (Also for Rb & Cs).
Peroxides and superoxides are unstable and react with themselves to form an oxide and
oxygen gas.
2H2O2(aq)  2H2O(l) + O2(g)
2Na2O2(s)  2Na2O(s) + O2(g)
4KO2(s)  2K2O(s) + 3O2(g)
NonMetal
Nonmetals are generally electronegative, have high ionization potentials, high electron
affinity, are poor electron conductors, come in different colors. They react with metals to
form ionic compounds. They react with other nonmetals to form molecular or covalent
compounds.
The chalcogens (O, S…) and halogens (F, Cl, Br, I) are typical nonmetals.
Fluorine is the most electronegative atom on the periodic table; very reactive. Chlorine is
more useful; it is added to water as a disinfectant:
Cl2(g) + H2O(l)  HCl(aq) + HClO(aq)
Nonmetal oxides are considered to be acidic because they react with water to form acids.
P4O10(s) + 6H2O(l)  4H3PO4(aq)
CO2(g) + H2O(l)  H2CO3(aq)
N2O5(s) + H2O(l)  2HNO3(aq)
Nonmetal oxides react with bases to form salts and water.
CO2(g) + 2NaOH(aq)  Na2CO3(aq) + H2O(l)
SO3(g) + 2KOH(aq)  K2SO4(aq) + H2O(l)
N2O5(g) + Ca(OH)2(aq)  Ca(NO3)2(aq) + H2O(l)
Semimetals or metalloids (B, Si, Ge, As, Sb, Te, Al, Zn) have intermediate properties
(amphoteric).