Download Chapter 5 Electrons in Atoms

Bio/Chem
Electrons in Atoms
I
The development of atomic models:
A. J.J. Thompson – proposes the “plum pudding” model
- negatively charged particles are stuck into a lump of positively charged
mass.
B. Rutherford – proposes the “nuclear atom” electrons surround a dense nucleus
C. Niels Bohr – (1913) Danish Physisist
- proposes that e-(s) are arranged in concentric circles (orbits) around the
nucleus
- often called the “planetary model”
1. An energy level is a region in space where an e- is likely moving.
2. Each energy level has a fixed energy
3. Electron(s) can move from one energy level to another if it gains or looses the
correct amount of energy.
4. A “quantum” of energy is the amount of energy required to move an e- from
its present energy level to the next higher one. (“Quantum Leap” – a large
abrupt change)
5. The energy levels in an atom are NOT evenly spaced. (They become closer
the farther you get from the nucleus. )
II
The Quantum Mechanical Model of the Atom
A. Erwin Schrodinger ( 1926) used the quantum theory to write and solve a
mathematical equation describing the location & energy of an electron in a
hydrogen atom.
*** It is primarily mathematical; it has few analogies in the visible world.
B. It does not define the exact path an electron takes around the nucleus, rather it
estimates the probability of finding an electron in a certain position.
C. The probability of finding an electron within a certain volume of space is
portrayed as a fuzzy cloud.
III
Atomic Orbitals
A. Energy levels are designated by principal quantum numbers (n).
1. these are assigned values of increasing energy
2. n = 1, 2, 3, 4, 5, … (Rows on the P.T.)
3. the average distance of the e-(s) from the nucleus also increases.
B. Within each principal energy level the electrons occupy energy sublevels (the
number of sublevels is the same as the principal quantum no.)
C. Atomic Orbitals (Cloud Shapes)
1.
2.
3.
4.
s orbitals are spherical with 1 orbital
p orbitals are dumbbell shaped with 3 orbitals
d orbitals are cloverleaf shaped with 5 orbitals
f orbitals are very complex with 7 orbitals
“Stupid pigeons don’t fly”
Principal E Level
No. of Sublevels
n=1
n=2
n=3
n=4
IV
1
2
3
4
Types of Sublevels
1s
2s 2p
3s 3p 3d
4s 4p 4d 4f
Max. No of e- (2n2)
2
8
18
32
Electron Configuration – the ways in which e- (s) are arranged around the nuclei
A. Three Rules Explain
1. Aufbau Principle - e- (s) enter orbitals of lowest energy first.
2. Pauli Exclusion Principle – an atomic orbital may describe at most 2e- (s)
3. Hund’s Rule – when e- (s) occupy orbitals of equal energy one e- enters each
orbital until until all the orbitals contain one e- with parallel spins.
Second e- (s) then add to each so that their spins are paired in the orbital.
*** See Diagram
B. Shorthand
1. Involves the energy level & the symbol for every sublevel occupied
by an e2. A superscript indicates the number of electrons
Ex. Carbon with 6 e- and in the 2nd row
_____ _____ _____ _____ _____
1s
2s
2p
** The sum of the superscripts always equals the number of electron(s).
C. Shorthand of the Shorthand
Carbon 1s2 2ss 2p2
1. find the Noble Gas that precedes the element, place [ ].
2. Add just the new energy level
[He] 2s2 2p2
V. Exceptions – there are some. Some “half-filled” levels are more stable than others.
i. e. - chromium
copper
VI
Light & Atomic Spectra
A. Every element emits light by passing electricity through its gas or vapor (See Lab)
1. Passing the light emitted by an element through the prism gives the atomic
emission spectrum of the element.
2. The emission spectrum of each element is unique to that element.