Electronic Structure of Atoms

Chapter 6Chemistry 100

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

If the principal quantum number is n, the shell can hold up to 2n2 electrons

s Orbitals are Spherical

p Orbitals are Dumbbell Shaped

d Orbitals are Complex

Aufbau Principle

1s 2s2p 3s 3p 4s3d 4p 5s 4d 5p 6s4f 5d 6p 7s 5f 6d6f

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.

Electronic Configuration & Periodic Table

I’m in a spin!!!

Nitrogen has Atomic Number 7 Electronic Configuration: 1s22s22p3

Let’s draw an orbital diagram: