Download Chemistry 102 Summary June 25th - Bohr model only works for one

Chemistry 102 Summary June 25th
-
Bohr model only works for one electron systems.
Current model called the Wave or Quantum mechanical model.
Several people contributed to its development.
Bohr: Energy is quantized, energy levels in atoms are quantized.
deBroglie: All matter exhibits both particle and wave properties.
deBroglie equation: λ = h/mν where h = Planck’s constant (6.626 x 10-34 Js)
m = mass (kg)
ν = frequency (s-1)
Example:
-
λball = 4.4 x 10-34 m
λelectron = 2.4 x 10-9 m
For a very massive ball, the wavelength is so short that we cannot
detect the wave motion (basically moves in a straight line).
For very small particles like electrons, the wave properties are
extremely important when discussing overall characteristics.
The movement of electrons about the nucleus cannot be
contained to simple circles as Bohr predicted but must include the
wave motion of the electron.
Heisenberg: Formulated the “Heisenberg Uncertainty Principle” which
states that it is impossible to know the exact location and momentum of
an electron at a given time.
- When we talk about the location of an electron, we always refer to the
probabilities of finding an electron.
Schrödinger: assigned wave properties to the electron.
- devised the Shrödinger equation:
-
Specific wave functions are called orbitals.
Orbitals define the allowed energy states where electrons can
reside.
There are four basic shapes: s, p, d and f
Shapes represent where an electron will reside 90 % of the time in
that allowed energy state.
From Heisenberg – the exact location cannot be determined but
instead the probability (Ψ2).
Each orbital (solution to Schrödinger equation) is characterized by a
series of quantum numbers which describe various properties of the
orbital.
We will only focus on n.
n is called the principle quantum number.
As n increases, the size and energy of the atomic orbitals also
increases.
The greater the value of n the further the electrons are from the
nucleus.
Example: A 2s orbital is larger and higher energy than a 1 s orbital.
Type of Orbital
Number of Degenerate
Orbitals (same Energy)
s (sphere)
p (peanut)
d (cloverleaf
except dz2)
f
-
1
3 (px, py, pz)
5 (dxy, dxz, dyz, dx2-y2, dz2)
Total
electron
capacity
2
6
10
Examples
1s, 2s, 3s, 4s etc.
2p, 3p, 4p, etc.
3d, 4d, 5d, etc.
7
14
4f, 5f, etc.
You are expected to know size and shape of all orbitals (except f
orbitals).
Shapes are related to the probability.
Each atomic orbital can hold a maximum of 2 electrons.
p orbitals contain nodes – regions of zero probability of finding an
electron.
Electron Configurations
Aufbau Principle: electrons occupy the lowest energy orbitals first.
- When electrons are in the lowest energy levels, called the ground
state which is the most stable arrangment.
- Any other allowed arrangement represents an excited state.
-
Two electrons can occupy any one atomic orbital.
-
When two electrons are in an atomic orbital they have opposite
spins.
Hund’s Rule: The lowest energy configuration for an atom is the one
having the maximum number of unpaired electrons.
-
The energy order of orbitals is:
1s<2s<2p<3s<3p<4s<3d<4p<5s<4d<5p<6s<4f<5d<6p<5f<6d<6f
Can memorize the order or use the periodic table.
Simple Examples of Electron Configurations:
(a) Hydrogen (H)
(b) Helium (He)
(c) Lithium (Li)
(d) Beryllium (Be)
(e) Boron (B)
(f) Carbon (C)
(g) Nitrogen (N)
(h) Oxygen (O)
(i) Fluorine (F)
(j) Neon (Ne)
For elements with a greater number of electrons, noble gas shorthand can
be used:
(a) Magnesium (Mg)
(b) Calcium (Ca)
(c) Chlorine (Cl)
(d) Bromine (Br)
Valence electrons: the electrons in the highest n value orbitals.
-
Notice that Mg and Ca and Cl, Br are in the same group and
contain the same number of valence electrons.
-
Similar numbers of valence electrons give rise to similar chemical
properties.
The group number gives the number of valence electrons.
More complicated examples:
(a) Platinum (Pt)
(b) 117
(c) At(d) Sn2+
*There are exceptions to the energy filling order but will not discuss in
detail.
Exceptions you are expected to know:
(i) Chromium: [Ar] 4s23d4 but is actually [Ar] 4s13d5
(ii) Copper: [Ar] 4s23d9 but is actually [Ar] 4s13d10
•
some extra stability with filled and half filled orbitals.
(iii) transition metal ions lose s electrons before d electrons when forming
cations.
Example: Mn: [Ar] 4s23d5
Mn2+: [Ar] 4s03d5
Mn4+: [Ar] 4s03d3
Cu: [Ar] 4s13d10
Cu+: [Ar] 4s03d10
Cu2+: [Ar] 4s03d9
*An apparent shifting in energies between s and d orbitals when ions are
formed.
- There are more exceptions, especially in the transition metal block of
elements but we will assume that the electron configuration expected
from the periodic table is correct (except for the exceptions mentioned).