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What is causing the light to appear when the atom is excited by electricity?
Why does the line spectrum of hydrogen have very distinct lines of color and not the complete rainbow, which includes all colors?
The Quantum Theory
• The main idea that comes from the Quantum Mechanical Theory and the Schrodinger Equation is Quantum numbers that define the surfaces of the electron densities in which the electrons reside.
• There are 3 main quantum numbers, n, l and ml, and a fourth quantum number, ms
Principle quantum number, nSchrödinger’s equation provides a number of possible functions that define the shape of the electron cloud about the nucleus, these shapes are the orbitals. The functions can be arranged in sets and subsets. Orbitals are classified in three ways:
An electron shell is a collection of orbitals, identified by an integer, n, which ranges in value infinitely up from 1. Electrons with the same value are in the same shell. In chemistry, important shells have n values between 1 and 7.
Angular Momentum, l
Subshells are groups of orbitals within an electron shell. They are identified by the letters s, p, d, f and through the alphabet beginning with g. In chemical systems, only s, p, d, and f subshells are important. The number of subshells within an electron shell is equal to the n value of the shell.
Magnetic, ml
Individual orbitals are identified by their direction in space using Cartesian coordinates. The number of orbitals within a subshell depends on the subshell type: s subshells have a single orbital, p subshells have 3 orbitals, d subshells 5 and f subshells 7 orbitals.
Examplen value l value ml value
n = 1, 2,… l = 0, 1, .. n-1 ml = 0, 1, …l n = 1 l = 0 ml = 0
n = 2 l = 0 ml = 0
l = 1 ml = +1, 0, -1
n = 3 l = 0 ml = 0
l = 1 ml = +1, 0, -1
l = 2 ml = +2,+1,0,-1,-2
n value l value Orbital designation n = 1 l = 0 1s
n = 2 l = 0 2s
l = 1 2px, 2py, 2pz
n = 3 l = 0 3s
l = 1 3p
l = 2 3dxy, 3dz2, 3dyz, 3dxz, 3dx
2-y
2
n = 4 l = 0 4s
l = 1 4p
l = 2 4d
l = 3 4f
The Energy Levels of the orbitals
In hydrogen the energy of the orbitals and the order in which they fill are dependant only upon their principle quantum number; however, in atoms with more electrons the order changes.
Energy levels in atoms of more than one electron.
Notice the shifting of the energy levels, as n increases, energy increases, as l increases, energy increases, but mixing of the energy levels still occur.
Orbitals10s
9s 9p 9d 9f 9g 9h 9i 9j 9k
8s 8p 8d 8f 8g 8h 8i 8j
7s 7p 7d 7f 7g 7h 7i
6s 6p 6d 6f 6g 6h
5s 5p 5d 5f 5g
4s 4p 4d 4f
3s 3p 3d
2s 2p
1s
Increasing Energy
• Electrons fill different orbitals in a subshell until the subshell is half - filled. (one in each orbital before pairing begins)
• If an atom has a filled valence shell, the atom is more stable.
• Each electron or the space it occupies can be defined by its four quantum numbers (n, l, ml, ms)
• No two electrons can occupy exactly the same space or have the same four quantum numbers.
• Some electron configurations defy expectations created by the periodic table, particularly for heavier elements.
Why do we have thoserows at the bottom?
Rb
H
Li
Na
Cs
K
LaBa
Fr
Be
Mg
Sr
Ca
Y
AcRa
Sc
TlHgAuHf PtIrOsReWTa
He
RnAtPoBiPb
CdAgZr PdRhRuTcMoNb
ZnCuTi NiCoFeMnCrV
In XeITeSbSn
Ga KrBrSeAsGe
Al ArClSPSi
B NeFONC
Gd
Cm
Tb
Bk
Sm
Pu
Eu
Am
Nd
U
Pm
Np
Ce
Th
Pr
Pa
Yb
No
Lu
Lr
Er
Fm
Tm
Md
Dy
Cf
Ho
Es
This arrangement takes toomuch space and is hard to read.
This arrangement takes toomuch space and is hard to read.
Hund’s RuleA full energy level is very stable, of lowest energy, a half filled shell is almost as stable as a fully filled shell and a partially filled shell is the least stable of highest energy. The orbitals will go from partially filled to half filled by putting one electron in each orbital before pairing the spins.
Electron Configurations
• Complete electron configurations
• noble gas electron configurations showing the noble gas and the remaining valence electrons
• valence electron configuration
Rules for assigning oxidation numbers
The sum of the oxidation number of all atoms must equal the net charge of the species.
In compounds: In compounds: Group IA are +1.Group IIA are +2.B and Al are +3, and F is -1.Hydrogen is +1 except when combined with a metal. Then
it is -1.Oxygen is -2 except for peroxides and superoxides.
Elements in their elemental state have an oxidation number of zero.
Tl+3+1
Hg+2+1
Au+3+1
Hf+4
Lu+3
Li+1
Na+1
Cs+1
Rb+1
K+1
Fr+1
Pt+4+2
Ir+4+3
Os+8+6
Re+7+6+4
W+6+4
Ta+5
H+1
He
RnAt-1
Po+2
Bi+5+3
Pb+4+2
Cd+2
Ag+1
Zr+4
Y+3
Pd+4+2
Rh+4+3+2
Ru+8 +6
+4+3
Tc+7+6+4
Mo+6+4+3
Nb+5+4
Lr+3
Ba+2
Be+2
Mg+2
Sr+2
Ca+2
Ra+2
Zn+2
Cu+2+1
Ti+4+3+2
Sc3+
Ni+2
Co+3+2
Fe+3+2
Mn+7 +6+4 +3
+2
Cr+6+3+2
V+5 +4
+3+2
In+3
Xe+6+4+2
I+7 +5
+1-1
Te+6+4-2
Sb+5+3-3
Sn+4+2
Ga+3
Kr+4+2
Br+5+1-1
Se6+4+2-
As5+3+3-
Ge+4-4
Al+3
ArCl
+7 +5+3 +1
-1
S+6 +4
+2-2
P+5+3-3
Si+4-4
B+3
NeF-1
O-1-2
N+5 +4+3 +2+1 -3
C+4-2-4
Common oxidation numbersCommon oxidation numbers
Oxidation numbers and the periodic table
Some observed trends in compounds.Some observed trends in compounds.
Metals have positive oxidation numbers.
Transition metals typically have more than one oxidation number.
Nonmetals and semimetals have both positive and negative oxidation numbers.
No element exists in a compound with an oxidation number greater than +8.
The most negative oxidation numbers equals 8 - the group number