Download Bohr Model, Quantum Mechanical Model

Bohr Model, Quantum Mechanical Model
1. Neils Bohr -connected spectra, and the quantum ideas of Einstein and Planck: the single
electron of the hydrogen atom could occupy only certain energy states, stationary states. This
explained the line spectra.
2. He explained that each line corresponded to a circular, fixed energy orbit around the nucleus.
3. There were many limitations to this model. Today we know that:
a. electrons exist only is discrete energy levels which are described by quantum numbers
b. energy is involved in moving an electron from one level to another.
4. Heisenberg Uncertainty Principle- It is impossible to know the momentum (mass of electron
times velocity) of an electron and its position in space at the same time. One or the other.
5. Quantum Mechanical Model- a mathematical model that incorporates both wave and particle
characteristics of electrons in atoms. Arises from work of Schrodinger which describe wave
functions and their corresponding energy characteristic and shape. This gives rise to orbitals,
which are the calculated probability of finding an electron of a given energy in a region of
space.
6. Quantum Numbers- a series of numbers which describe properties of orbitals.
a. principal quantum number (n)- an integer greater than zero, related to the size of the orbital.
As n increases:
i. orbital becomes larger
ii. electron spends more time further from the nucleus
iii. Higher energy because the electron is less tightly bound to the positive nucleus.
iv. A collection of orbitals with the same value of n is called an electron shell.
b. angular momentum quantum number (l)- is an integer from 0 to (n-1). This quantum number
defines the shape of the orbital
c. magnetic quantum number (ml)- is an integer with values between –l to l, including zero,
related to the orientation of the orbital in space.
d. spin magnetic quantum number- (ms) – can be +1/2 or -1/2, indicating opposite directions that
an electron can spin.
Exercise 6
Electron Subshells
For principal quantum level n = 5, determine the number of allowed subshells (different
values of ℓ), and give the designation of each.
ℓ = 0; 5s
ℓ = 1; 5p
ℓ = 2; 5d
ℓ = 3; 5f
ℓ = 4; 5g
7. Orbital Shapes and Energies
a. . Size of orbitals
1. Defined as the surface that contains 90% of the total electron probability
2. Orbitals of the same shape (s, for instance) grow larger as n increases
b. s Orbitals
1. Spherical shape Occur in levels n=1 and larger.
2. Nodes (s orbitals of n=2 or greater)
c. p Orbitals
1. Two lobes each
2. Occur in levels n=2 and greater
3. Each orbital lies along an axis (2px, 2py, 2pz)
d. d Orbitals1. Occur in levels n=3 and greater
2. Two fundamental shapes
a. Four orbitals with four lobes each, centered in the plane indicated in the orbital
label.
b. Fifth orbital is uniquely shaped - two lobes along the z axis and
a belt centered in the xy plane
e. f Orbitals
1. Occur in levels n=4 and greater
2. Highly complex shapes
3. Not involved in bonding in most compounds
8. Orbital Energies
1. All orbitals with the same value of n have the same energy
a. "degenerate orbital" (hydrogen only!)
2. The lowest energy state is called the "ground state"
3. When the atom absorbs energy, electrons may move to higher energy
orbitals - "excited state"
9. Electron Configuration- distribution of electrons among various orbitals of an atom.
10. Paramagnetic- an atom having one or more un-paired electrons.
Dimagnetic- all electrons in an atom are paired.
11. Nobel Gas configuration- show only electrons occupying the outermost sublevels. Insert
[nobel gas element closest with lower atomic number] + outermost sublevels.
Ex. [Ne] 3s1 --takes the place of 1s22s22p6
12. Penetration- Electron density profiles show that s electrons penetrate to the nucleus
more than other orbital types
a. Closer proximity to the nucleus = lower energy
b. Closer proximity = greater attraction