Chapter 6 Over-View
Part 1: ALREADY DID THIS! CHAPTER 2!!
classical view of atomic structure
Part 2: 6.1-6.2 light and quantum theory
Part 3: 6.3-6.4 quantum view of atomic structure
Part 4: Application: Trends!
I. History of the Atom (chapter 2!)
A.
B.
C.
D.
E.
Democritus: Smallest part of matter (atomos)
Dalton: Small solid spheres, Atomic theory
Thomson: Electrons, Plum pudding, cathode ray tube
Millikan: electron charge (oil drop experiment)
Rutherford: +, center, dense, nucleus
II. Nuclear Atom
A. Ernest Rutherford (1871-1937)
“It is about as incredible as if you had fired a 15-inch shell at a piece of
tissue paper and it came back and hit you”
B. Gold Foil
gold foil
C. Bohr: Planetary model
1. Answered question of what keeps
e- from falling into the nucleus
a. e-have a fixed amount of
energy in energy levels
-region around the
nucleus that e- is moving
-NOT a concentric circle!
2. electrons can skip to higher energy levels
a. rungs of a ladder
b. quantum leap
-energy required to move e- from
one rung to the next
c. Energy levels increase in energy
the farther you get away from the
nucleus
-”rungs” become closer at top
-easier for e- to escape (farther
from nucleus)
III. Influences on Quantum Mechanical Model
A. Study of light
1. Isaac Newton (1642-1727)
a. light made up of particles
2. Thomas Young (1802)
a. light made up of waves!
http://www.gilestv.com/tutorials/quantum.html
B. Light waves
1. Light made up of electromagnetic waves
2. Wave parts
a. Amplitude: height from origin to crest
b. Wavelength: distance b/twn crests
c. Frequency: Cycles/second (s-1 or Hz)
3. Product of frequency and wavelength (inverse)
always equaled the same number…3.0 x 1010cm/s
Note: frequency and wavelength are inverse
Example
Example
An FM radio station found at 103.1 on the FM
dial broadcasts at a frequency of 1.031x 108
s-1 (103.1 MHz). What is the wavelength of
these radio waves in meters?
4. Electromagnetic spectrum
a. visible, UV, infrared, microwave radiation.
b. Visible 700nm(red)
c. ROY G BIV
380nm(blue)
C. Young’s experiment
1. Light shined at 2 slits
a. If particles you should see 2
spots on the wall
b. Results:
-interference pattern: light/dark bands
c. Conclusion:
-wave nature to light!
Figure 5: Young’s
Experiment
NOTE: Continuous spectrum of various
types of radiation!
Flame Tests
D. Atomic emission spectra
1. Aka line spectra or discontinuous spectra
2. Elements give a specific light
when excited with electricity
-this light can be broken up to
see the emission spectra
Question: why is this spectrum discontinuous?
Classical, wave theory, could not explain line spectra, black body
radiation, or photoelectric effect…enter Max Planck (1900)
E. Max Planck (1858-1947)
1. Why emission not continuous?
a. Ex: iron is heated and undergoes a series of color
changes: red-yellow-white-blue
-black body radiation
2. Energy in a body must change in small units!
-like bricks in a wall.
-absorb or emit certain
“quanta” of energy.
3. Planck Math
a. Energy absorbed/emitted is directly
related to frequency of the waves.
b. Vibrating atoms in a heated solid absorb/emit
energy in discrete amounts (quanta)
c. E = h x v
E = energy in Joules (J) of a
“quantum”
h = Planck’s constant
(6.6162 x 10-34 JxS)
v = frequency (s-1 or Hz)
4. Analogy for Planck
a. Heating water
-seems continuous
-actually happening in infinitesimally
small steps while molecules absorb
“quanta” of energy
F. Albert Einstein (1905)
1. Back to particle!
2. Took Newton’s idea of “particles” of
light and developed it.
3. Proposal
a. light described as quanta of energy that
behaves as if particles
-photon
-Quantized by Planck
E=hxv
(variations in E are discontinuous! hv, 2hv, 3hv…)
4. Proposal proof
a. Photoelectric effect
-metals get excited when hit with light and
eject electrons
-BUT: only certain energy/frequency of
light will cause the effect
b. Classical physics
-no quantum concept
-light is energy, but waves
-All metals should eventually absorb enough
energy to eject electrons…not what is observed!
c. Einstein’s proposal
-threshold value of energy (hence frequency) below
which the effect does not occur.
- if: E = h x v
(h= Planck’s constant and black body effect)
-then: all “photons” in a certain color of light have
the same energy so same frequency.
-if not at threshold then photoelectric
effect will not happen.
c. photoelectric effect and photons
1. energy unit of one photon is small But one mol is measurable so
bring in Avagadro’s number…
E = Nahν
● 2. summary
● -light phenomena needs classical physics and electromagnetic
waves
● -black body radiation, photoelectric effect, discontinuous spectra
require photon/particle view
● Hence: light has a wave/particle duality!
Example: Given: violet light photon
ν = 6.15 x 10 14 s-1
Find: E = ?
Solve: E = hν
E = (6.626 x 10-34 Js/photon)(6.15 x 10 14 s-1 )
E = 4.07 x 10 -19 J/photon
Example: Given: uv light λ = 235nm
Find: E = ? J/photon
Solve: c = λν
ν = c/λ = (3.00 x 108m/s)/(235nm)(109nm/1m)
E = hν
E = (6.626 x 10-34 Js/photon)[(3.00 x
108m/s)/(235nm)(109nm/1m)]
E = 8.46 x 10-19 J/photon
Example: Given: red light λ = 632.8nm, 1mol photons
Find: E = ?? kJ
Solve: 1) find ν of red light using c = λν
ν = 4.738 x 1014s-1
2) E = NAhν
E = (6.022 x 1023 photons/mol)(6.626 x 10-34 JS/photon)(4.738 x 1014 s-1)
E = 189.1kJ/mol
G. Explanation of Emission Spectra
1. Hydrogen atom
a. One e-unexcited “ground state”
-absorbs “quanta” of energy
-ex: electricity
-jumps to certain, higher energy levels
-come back to “ground state” they release light
b. En
-energy of this electron is quantized
-all allowable values for an energy level given by…
En = -B/n2
B = 2.179 x 10-18 J (p276)
n = interger
Note: as long as e- remains in a given energy level it can’t emit energy
as electromagnetic radiation…this prevents it from spiraling into the
nucleus! (classical waves have it spiraling into the nucleus!)
Example
Calculate the energy of an e- in the 2nd energy level of a H atom.
En = -B/n2
E2 = - B/22
E2 = -2.179 x 10-18 J/4 = -5.448 x 10-19 J
2. line spectra
-to explain this discontinuous spectra Bohr came up with the
following:ΔElevel = Ef – Ei where En = -B/n2 so…
ΔElevel = B(1/ni2 – 1/nf2)
nf >ni, e- absorbs quantum and moves farther from nucleus ΔElevel = (+)
ni>nf, e- absorbs quantum and moves closer from nucleus ΔElevel = (-)
Example: Calculate the energy changes in joules when an e- falls from
ni = 5 to nf = 3 in a H atom.
ΔElevel = B(1/ni2 – 1/nf2)
= 2.179 x10-18(1/52 – 1/32) = -1.550 x10-19 J
Notes: negative sign means energy emitted as photons
Example: Next from above problem would be to determine ν and λ
when e- drops.
Find: ν = ? λ = ?
Solve: ΔElevel = hν
● ν = E/h = 1.550x10-19 J/6.626x10-34 Js = 2.339 x 1014s-1
● λ = c/ν = 3.00 x 108 m/s / 2.339 x 1014s-1 = ___________________
● Note: this λ is in the infrared part of the spectrum
H. Louis de Broglie (1892-1987)
1. Question:
a. Light behaves as particles and waves…
does matter do the same?
-yes
λ = h/mv
λ = wavelength
h = Planck’s constant
m = mass of particle
v = velocity of particle
Example: Given: Ve = 2.74 x 106 m/s
me = 9.11 x 10-21 kg
1J = 1kgm2s-2
Find: λ = ?? m and nm
Solve: λ = h/mv
= (6.626 x 10-34 kgm2s-2 ) / (9.11 x 10-21 kg)(2.74 x 106 m/s)
= 2.65 x 10-10 m or 0.265 nm
I.
Werner Heisenberg (1901-1976)
A. Heisenberg Uncertainty Principle
1. impossible to know both
velocity and position
of a particle at the same time
IV. Quantum Mechanical Model
A. Erwin Schrodinger (1887-1961)
1. Used math to describe location and energy of electron in
Hydrogen
2. developed a wave equation to describe the H atom
a. solutions to his equation called wave functions (ψ)
3. Model is based on Schrodinger’s solution to de Broglie’s
equation…
a. All math…no physical model!
b. Probability of finding an electron in a certain position.
4. ψ = wave function (energy state of atom)
a. Allows to find the probability of an e- in a certain region.
Note: can’t know position and velocity
so get geometry probabilities!
B. Atomic Orbitals
1. e- is everywhere at once in an energy “orbital”
a. Cloud: -e found here 90% of time
2. quantum numbers
a. n = principle quantum number
-whole #, + integers… 1,2,3,4…
-higher #, higher energy
-higher #, larger distance from nucleus
b. = orbital angular momentum quantum number
-shape
-whole #, + integers with a value of n-1
-s, p, d, f
0, 1, 2, 3
-n value tells how many orbitals there are!
orbitron
V. Electron organization
A. Principal energy levels (n)
a. 1, 2, 3, 4…
b. # tells # of sublevels
B. Sublevels
a. One “s” sublevel (orbital)
-hold 2 electrons
b. Three “p” sublevels
-hold 2 electrons each
c. Five “d” sublevels
-hold 2 electrons each
VI. Electron Configuration
A. Electron “address”
B. Three rules
1. Aufbau Principle
a. e- enter orbitals of lowest
energy levels first
- 1234
- spdf
b. however not always this
simple in multielectron atoms!
2. Pauli Exclusion Principle
a. an atomic orbital holds 2 eb. e-have spin to them
c. 2e- in same orbital must have opposite spins
3. Hund’s Rule
a. One e- will enter each of the
three p orbitals before pairing up. (parallel)
C. Types
1. spdf notation (expanded)
2. orbital diagrams
3. noble gas configuration
D. Magnetic properties
1. diamagnetism
a. weak repulsion of paired e-s to an external magnetic field
2. paramagnetism
b. stronger attraction of unpaired e-s to an external magnetic field
Write electron configurations for the
following:
Argon
Ar 1s12s22p63s23p63d1
Sodium
Na 1s12s22p63s2
Phosphorus
P 1s12s22p63s23p4
Write arrow diagrams for the above.
E. Exceptions to Aufbau Filling
1. Filled or half filled “d”sublevels are more
stable than alternatives.
Ex: Cr 1s12s22p63s23p63d44s2
Cr 1s12s22p63s23p63d54s1
Ex: Cu 1s12s22p63s23p63d94s2
Cu 1s12s22p63s23p63d104s1
VII. Periodic Trends
A. Periodic Law
B. Atomic properties (trends)
1. based on…
a. e- shielding (period and group trend?)
b. nuclear pull (period and group trend?)
c. effective nuclear charge (Zeff)
Zeff = nuclear charge – core e- count
Fe: 26-24 = 2 (124pm), Co: 27-25 = 2 (124pm), Ni: 28-26 = 2 (125pm)
2. Types
a. Atomic radii (picometers (10-12) or angstrom (10-10))
b. Ionic radii
-cations lose e- so smaller than parent
*nuclear charge > than e- # so pulls in!
-anions gain e- so larger than parent
*more e- to repel one another so it spreads!
c. ionization energy
1. energy needed to remove an e2. period trend? Group trend?
d. Electron affinity (EA)
1. adding an electron
2. the more (-) the value the higher the affinity
period trend?
group trend?
e. Metallic character
1. the easier it is to remove e- the more metallic
period trend?
group trend?