Transcript 5-1 Quantum Theory of the atom

5-1 QUANTUM THEORY OF THE ATOM

Bohr Model of the Atom

Proposed that electrons orbited the nucleus in
circular paths.

Ground state- lowest allowable energy states of an
atom.

Excited state- atom gains energy; H atoms can have many
different excited states although it contains 1 e-.
Electrons move around a H atom in circular orbit
 Orbits equal to a principal quantum number n, where
n=1 is lowest energy level, closest to nucleus.

BOHR MODEL OF THE ATOM

n =6
n =5
n =4
n =3
n =2
Orbits/ levels are like rungs
in step ladder
Cannot stand b/w rungs, ecan’t exist b/w levels (orbits).
 E- move from 1 orbit to the
next emitting or absorbing
certain amounts of energy
(quanta).


n =1

nucleus
The smaller the e- orbit, the
lower the energy state/level
The larger the e- orbit, the
higher the energy state/level
5-1 QUANTUM THEORY AND THE ATOM

Quantum mechanical model is the modern
atomic model and comes from
Louis De Broglie: radiation (energy) behaves like
particles and vice versa.
A.
1.
2.
All particles w/ a mass have wave characteristics
E- move around nucleus in a wave-like manner
Heisenberg uncertainty principle- impossible
to know both the velocity and position of an e- at
the same time.
C. Shrodinger: e-’s energy are limited to certain
values (quantum) but does not predict path
B.
1.
2.
Treated e-’s as waves
Created wave function = predicts probability of finding
e- in a volume of space (location)
HYDROGEN’S ATOMIC ORBITALS
Shrodinger’s wave eqn predicts atomic orbitals
 Atomic orbital - 3D regions around the nucleus
that describes the e-’s probable location.

a.
b.
c.
atomic orbital = fuzzy cloud
Do not have a defined size
Shape = volume that contains 90% of the probable
location of e-’s inside that region.
QUANTUM MECHANICAL MODEL

Like Bohr, electrons occupy space surrounding
the nucleus and exist in several principal energy
levels = principal quantum number (n)
Relative size and energies of atomic orbital
 n = 1,2, 3, etc. = period


Principal energy levels consist of energy sublevels
with different energy values.

Energy sublevels – shape of the atoms’ orbitals
s = spherical
p = dumbbell
d, f= different shapes
QUANTUM MECHANICAL MODEL

Principal energy levels have specific allowed
sublevels - shapes.
s
p
s
d
p
s
p
s

f
4
d
3
2
n= 1
s sublevel is lower in energy and f has higher
energy
QUANTUM MECHANICAL MODEL

Sublevels consist of orbitals of different orientation.
Orbitals in same sublevel are = in energy (no matter
orientation)
 Orbitals only hold 2e- maximum with opposite spins (+ or –
spins).
Sublevel
Orientations/ Orbitals
Max # es
1
2
p
3
6
d
5
10
f
7
14

ORIENTATIONS/ ORBITALS PER SUBLEVEL

s- spherical only 1 orbital orientation

p- dumbbell has 3 orbital orientations

d- 2dumbbells with 5 orbital orientations

f- 3dumbbells with 7 orbital orientations

http://winter.group.shef.ac.uk/orbitron/AOs/1s/in
dex.html
BOHR MODEL OF THE ATOM

Hydrogen’s Line Spectrum (AES)
At n= 1 H atom is in ground state
 When energy is added, e- moves to higher energy level, n=2
(excited state).
 e- drop back to lower energy level n=1 and emitts a photon
equal to the difference b/w levels.

A photon is emitted
with E= hυ
A photon is
absorbed
HYDROGEN’S LINE SPECTRUM

Lines which show up have specific energies which
correspond to a frequency of a color of light.
Energy of Hydrogen Atom
E= 4.85 x 10-19 J
n
6
5
4
3
E= 3.03 x 10-19 J
2
1
A photon is
emitted with
E= hυ for each
frequency
5-2 ELECTRON CONFIGURATIONS
Electron configuration – arrangement of e- in
atoms; lower nrg arrangements
 Arrangements defined by:

Aufbau principle – e- occupy lowest nrg orbital
available
1.
a.
b.

c.
d.
a.
All orbitals in a sublevel are = in nrg (px py pz )
Sublevels within an energy level have different energies
Ex: 2s lower in nrg than 2p
Order of energy = s, p, d, f
Sublevels in one energy level can overlap with sublevels in
another principal energy level.
Ex: 4s lower in nrg than 3d
AUFBAU DIAGRAM
ELECTRON CONFIGURATIONS
2.
3.
Pauli exclusion principle – a max of 2 emay occupy a single orbital only if they have
opposite spins.
Hund’s rule – energy charged e- repel each
other.

All same nrg orbitals are filled first with econtaining same spin before extra e- can occupy the
same orbital with opposite spins.

Ex: 3 orbitals of 2p
2px 2py 2pz
FILLING SUBLEVELS WITH ELECTRONS
Energy sublevels are filled from lower energy to
higher energy following the diagram.

ALWAYS start at the beginning of each level and
follow it until all e- in an element have been placed.
Increasing Energy

1s
2s
3s
4s
5s
6s
7s
2p
3p
4p
5p
6p
7p
3d
4d 4f
5d 5f
6d
ORBITAL DIAGRAM AND E- CONFIGURATIONS

Orbital diagram for Fe:
Iron has how many e- ?
 26 e
1s 2s

3s
3p
4s
3d
Electron configuration for Fe:



2p
Iron has 26 e1s2 2s2 2p6 3s2 3p6 4s2 3d6
Shortcut to the E- config. for Fe is Noble gas notation
Group 18 or 8A are the Nobel Gases
 Argon has 18 e1s2 2s2 2p6 3s2 3p6
 Iron has 26 e1s2 2s2 2p6 3s2 3p6 4s2 3d6
 Noble gas notation: [Ar] 4s2 3d6

[
]
VALENCE ELECTRONS AND ELECTRON DOT
STRUCTURES

Valence electrons – outer energy level/orbital
electrons which are involved in bonding.
Valence electrons = groups 1A to 8A
 B GROUPS DO NOT COUNT


E- dot structures- consists of the element’s:
a.
b.
c.
Symbol - represents the atomic nucleus & innerlevel electrons
Surrounded by dots- represent the valence
electrons.
Ex: O = 1s2 2s2 2p4 or [He]2s2 2p4 ve- =6 in grp 6A
O
PERIODIC TABLE SHORTCUT
Periods = Energy Level
1A
8A
Groups (A only) = Valence e2A
3A 4A 5A 6A 7A
Energy level = n-1 for d sublevel
Energy level =
n-2 for f sublevel
5-3 Light and Quantized
Energy
Some elements emit visible light when
heated with a flame.
 This chemical behavior is due to the
arrangement of e- in atoms.

ELECTROMAGNETIC RADIATION
Form of energy that exhibits wave-like behavior
as it travels through space.
 There are many types of electromagnetic
radiation and all are represented in the

ELECTROMAGNETIC SPECTRUM
ELECTROMAGNETIC SPECTRUM
PARTS OF A WAVE

Frequency (v, nu) –The number of complete
wavelengths that pass a given point each second.
 Units: wave/second = 1/s = s-1 = Hertz (Hz)

Wavelength (l, lambda) – The distance between
identical points on successive waves. (crest to
crest or trough to trough)
 Units: meters (m)
c=lv
c = speed of light, 3.00 x 108 m/s
WAVE NATURE OF LIGHT

Max Planck theorized that all matter
can gain/ lose energy in small “chunks”
of light (quanta).

Quantum- minimum amt of energy that
can be gained or lost by an atom.
Ex: Iron when hot appears red or blue, emits
energy that is quantized has a specific
frequency.
o Heating water – temp increases by molecules
absorbing a specific amt or quanta.
o

Calculated as follows:
Equantum= hv
E = Energy (J)
o h = Planck’s constant 6.626 x 10-34 (J s)
o v = frequency ( Hz or s-1)
o
PARTICLE NATURE OF LIGHT
Photoelectric effect – electrons are emitted from a
metal’s surface when light of a specific frequency
shines on the surface.
 Albert Einstein (1905) assumed that light
travelled as a stream of tiny particles or packets of
energy called photons.
 Photons- EM radiation w/ no mass that carries a
quantum of energy.

EM radiation has both wavelike and particle- like nature.
 Ephoton= hv
 Photon = quantum of energy

ATOMIC EMISSION SPECTRA
Set of frequencies of light waves emitted by an
atom of an element.
 Line spectrum – consists of several individual
lines of color from light energy emitted by excited
unstable atoms


Only certain colors (frequencies) appear in an
element’s AES & it can be used to identify the
element.