electronic structure of atoms chapter 6 chemistry 100
TRANSCRIPT
What is light?
Light is obviously real - it is part of our world. Darkness is the absence of light
Light is NOT a solid, a liquid, or even a gas. So what is it? It is a form of energy
A form of radiant energy because it carries energy through space
Electromagnetic radiation
Visible light is a type of electromagnetic radiation
Other types include: infra-red, ultra-violet, X-rays, radar waves, microwaves, radio and TV waves
Electromagnetic radiation has wave like properties
Waves
Wavelength (lamda) Frequency (nu) Speed c (see?) = c c = 3.00 108m/s for
all types of electromagnetic radiation.
So how is IR different from UV, for example?
Electromagnetic Spectrum
Electromagnetic radiation is characterized by a wave length () and a frequency ()
Frequency: number of cycles (vibrations) per second. Unit is second-1 or s-1 or the Hertz (SI
unit for frequency). Hence, 82,000 s-1 is the same as 82 kHz
(kiloHertz)
Units for wavelength
Unit Symbol Length (m) Type of Radiation
Angstrom Å 10-10
X-rayNanometre nm 10
-9UV & visible
Micrometre mm 10-6
IRMillimetre mm 10
-3IR
Centimetre cm 10-2
microwaveMetre m 1 TV, Radio
Max Planck and his constant h
Suggested that energy is quantized - comes in small chunks
E = h where n = 1, 2, 3
Compare the potential energy of a brick on a staircase to one on a slope
Can this be true?
We do not find that energy is quantized in everyday life - h is very small. Cannot see the difference between
200,000,000h and 200,000,001 h Einstein used Planck’s idea to explain the
photoelectric effect For electromagnetic radiation, E = h where is
the frequency of the radiation. High frequency more energy
What is light?
Examine how light behaves in experiments with lens, mirrors, etc., we are led to believe that light has
wave properties In the photoelectric effect, light
appears to consist of particles - which we call photons
Dual nature of electromagnetic radiation
Bohr’s Atom
Bohr said: if energy is quantized then the energy of an electron in an atom is quantized
Radius of its orbit cannot be any arbitrary value
Must obey the quantum theory. Only certain orbits are allowed
Allowed Orbitals in Bohr’s Atom
2
2
220
4
n
22
20
n
nZ
h8me
E
...3 ,2 ,1n where nZme
hr
The quantity n is a quantum number
Bohr’s Atom 1913
Electrons move in orbitals with specified radii
Each orbital is associated with a specific energy
This explains why atoms emit (or absorb) light of well-defined frequency. Examples: the yellow sodium street
light and the neon tube.
Wave Behaviour
Louis de Broglie (1892-1987) If light can have both wave and particle
behaviour, why not wave behaviour for all particles?
= h/m He talked about matter waves
Matter waves
Find for electron moving at 5.97 106 m/s
rays- Xof wavelength the toSimilar nm122.0m1022.1
kg1g10
J1sm.kg1
)s/m1097.5)(g1011.9(Js1063.6
mvh
10
322
629
34
Find for baseball moving at 100 km/h
meaningfulor - measurable be to small tooFar m1065.1
hrs3600
km1m1000
kg1g10
J1sm.kg1
)hr/km100)(g145(Js1063.6
mvh
34
32234
Heisenberg
Postulated that there is a limit to how precisely we can measure both position and momentum
The measurement effects the object being measured
Heisenberg’s Uncertainty Principle
Schrödinger’s wave equation
In 1926, Schrödinger put de Broglie’s and Heisenberg’s ideas together and came up with the wave equation
The quantity 2 provides information about the electron's position when it has
energy E!
!!equation!ugly truly A EVdxd
m8h
2
2
2
2
Quantum Numbers
Schrödinger's wave equation has three quantum numbers. Principal quantum number n. Has integer
values 1, 2, 3 Azimuthal quantum number, l. Allowed values
values of 0, 1... up to n - 1 Magnetic quantum number, ml. Allowed values
-l … 0 … +l There is also the Spin quantum number, ms. It
can have a value of -½ or +½
Atomic orbitals
The first shell n = 1 The shell nearest the nucleus l = 0 We call this the s subshell (l = 0) m = 0 There is one orbital in the subshell
s = -½ The orbital can hold two electronss = + ½ one with spin “up”, one “down”
No two electrons in an atom can have the same value for the four quantum numbers: Pauli’s Exclusion Principle
The second shell
n = 2 l = 0 or 1 There are two subshells
l = 1 The p subshell
m = -1, 0, +1 Three orbitals in the subshells = -½ or + ½ Each orbital can hold 2 electrons.
p subshell can hold 6 electrons
l = 0 The s subshell
m = 0 One orbital in the subshell
s = -½ or + ½ Subshell can hold two electrons
The second shell can hold 8 electrons:
2 in s orbitals and 6 in p orbitals
Let’s do Sodium, Z = 11
Aufbau Principle 1s 2s 2p 3s …. First 2 electrons 1s2 that’s
2 Next 2 electrons 2s2 that’s
4 Six this time 2p6 that’s
10 1 more to go 3s1 that’s
all, folksElectronic configuration of Na is 1s22s22p63s1
Hund’s Rule
The configuration with the maximum spin is more stable.
Shall we use1s 2s 2p() () ()()
Or, shall we use1s 2s 2p
() () ()()()
Shorthand configurations
The configuration of Neon is: 1s22s22p6
Na is 1s22s22p63s1, or in short form: [Ne]3s1 The configuration of Argon:1s22s22p63s23p6
K is: 1s22s22p63s23p64s1, which in short form becomes [Ar]4s1
Note the similarity of the two elements from the same group in the periodic table.The incomplete orbitals are 3s1 and 4s1.
Same group, similar configuration
Fluorine: [He]2s22p5
Chlorine: [Ne]3s23p5
Bromine: [Ar]3d104s24p5
Iodine: [Kr]4d105s25p5
The outer-shell configuration in each case is s2p5
We need not be concerned with the d electrons here because d10 is a filled subshell.