Unit 3: Electrons in the Atom

Copyright © Houghton Mifflin Company

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The Rutherford atom model.

A positive nucleus surrounded by electrons like our solar system.

However, this model did not properly explain chemical reactivity and certain light phenomena.

In 1665 Sir Isaac Newton noticed that white (sun) light could be split into a multicolored band of light just like a rainbow. The multicolored band of light is called a color spectrum.

Brief History of Light

Light as a wave• In the 19th century the works of Michael Faraday

and (later) James Maxwell showed that electricity and magnetism are simply two parts of a single phenomenon, electromagnetism.

• This phenomenon would produce waves which travel at the speed of light, having light waves other than those that produced the light that we could see. We now refer to this collection of different waves of electromagnetic radiation (light) as the electromagnetic spectrum (EMS)

Waves• There are 2 types of waves, transverse

and longitudinal.

Waves are characterized by three properties:Waves are characterized by three properties:1. wavelength, 2. 1. wavelength, 2. frequency, 3.frequency, 3. amplitude amplitude

Wavelength and Frequency

• An important feature of wave motion is the inverse relationship between wavelength and frequency. That is, as one increases the other decreases.

• They are related in the following way: c = λ ν• Speed of light (c) = wavelength (λ) x frequency (ν)

• c has a constant value of 3.00 x 108 m/s.• Wavelength is in meters (m).• Frequency is in cycles/s, in s-1, in 1/s, or in

Hertz (Hz).

Example

• Calculate the frequency of light with a wavelength of 5.22 x 10-10 m.

• c = λ ν

• 3.00 x 108 m/s = 5.22 x 10-10 m x ν

so ν = (3.00 x 108 m/s) / (5.22 x 10-10 m) = 5.75 x 1017/s

= 5.75 x 1017Hz

Practice

• Calculate the wavelength of a radio station signal with a frequency of 99.7 MHz.

ElectroElectromagneticmagnetic SpectrumSpectrum

In increasing frequency,In increasing frequency, RROOYY GG BBIIVV

Electromagnetic Electromagnetic SpectrumSpectrum

Electromagnetic Electromagnetic SpectrumSpectrum

Long wavelength --> small frequencyLong wavelength --> small frequency

Short wavelength --> high frequencyShort wavelength --> high frequency

increasing increasing frequencyfrequency

increasing increasing wavelengthwavelength

Light as particles

• In 1900 German scientist Max Planck found that light is given off in discrete units (quanta).

• He also found light energy (E) is proportional to its frequency (ν).

• The relationship is: E = hν, where h = Planck’s constant, 6.63 x 10-34 J . S.

Wave particle nature• In 1905 Albert

Einstein confirmed Planck’s findings and he called the quanta “photons” (packets of energy).

Example• Calculate the energy of a photon of light with a

frequency of 5.45 x 1014 Hz.• E = hν = 6.63 x 10-34 Js x 5.45 x 1014 s-1

= 3.61 x 10-19 J.

You try: Calculate the energy of a photon having a wavelength of 4.5 x 10-7 m.

Bohr’s observations• In 1913 Niels Bohr used the observations

of Planck to explain the specific lines observed in the hydrogen emission spectrum. These lines resulted from some whole number transition.

Bohr’s model• Bohr suggested that a transition corresponded

to an electron jumping from one possible orbit to another and emitting a photon of light energy.

• In Bohr’s model of the atom, the electron can only exist in these specific orbits, known as energy levels, in an atom.

• Normally the electron would be in its lowest available energy level, this is called its ground state.

If the atom is exposed to an energy source the electron can absorb a quantum of energy (photon) and the electron will make a quantum leap to a higher energy level. Then the electron will drop

back down to a lower energy level, thereby emitting a photon of light. The energy of this photon would correspond exactly to the energy difference between the two levels.

Light emitted produces a unique emission spectrum.

n=1

n=2

n=3

n=4

Spectrum

UV

IR

Vi s ible

Ground State

Excited State

Excited StateExcited State unstable and drops back down

•Energy released as a photon

•Frequency proportional to energy drop

Excited State

But only as far as n = 2 this time

Emission Spectrum Animation

Line Spectra of Other Line Spectra of Other ElementsElements

Wave – particle duality

• Louis De Broglie (1924) proposed that ALL matter has wave and particle properties, not just electrons.

• Heisenberg (1927) said that because of size and speed it is impossible to know both exact position and momentum of an electron at the same time.– This is referred to as “Heisenberg Uncertainty

Principle”

Quantum mechanical model• Schroedinger (1887-1961) developed the

“quantum mechanical model” of the atom. • He calculated the probability where to find

electrons, thereby creating “electron clouds”: areas with a great chance (90 %) to find electrons.

• The region in space in which there is a high probability of finding an electron is now known as an “orbital”.

http://physicalworld.org/restless_universe/figs/fig_1_30.gif

Orbitals

• Every element has discrete energy levels called principal energy levels (given with letter n).

• The principal levels are divided into sublevels.• Sublevels contain spaces for the electron called

orbitals.

Orbital types

• s-orbital = spherical shape, only 1 of them• p-orbital = gumdrop or dumbell shape, 3 of them

– one on each axis (x,y,z)• d-orbital = donut shape, 5 of them• f-orbital = cigar shape, 7 of them• Each orbital contains a max of 2 electrons

s and p orbitals

d orbitals

Filling orbitals with electrons• We have 3 general rules for “distributing” these

electrons.

1.Pauli Exclusion Principal: Orbitals contain no more than two electrons. – Each electron has a spin: up (↑) or down (↓)– Two electrons must have opposite spins to occupy an

orbital

2. Hund Rule: When filling orbitals, assign one electron to each orbital (of that type)

before doubling up with two electrons per orbital.

3. Aufbau: Electrons fill lowest orbitals first, then proceed to higher energy levels.

First 4 energy levels

Filling orbitalsEnergy level

Orbital type

# orbitals

# of types # electrons

n = 1 s 1 1 s 2

n = 2 s, p 4 1 s, 3 p 8

n = 3 s, p, d 9 1 s, 3 p, 5 d 18

n = 4 s, p, d, f 16 1 s, 3 p, 5 d, 7 f

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Energy level = the number of orbital typesTotal number of orbitals in an energy level = n2

Total number of electrons in any energy level = 2n2

Electron configuration

• The electron configuration of an atom is a shorthand method of writing the location of electrons by sublevel.

• The sublevel is written followed by a superscript with the number of electrons in the sublevel.

Electron configuration

• H 1s1 He 1s2

• Li 1s22s1 Be 1s2 2s2

• B 1s2 2s2 2p1 C 1s2 2s2 2p2

• N 1s2 2s2 2p3 O 1s2 2s2 2p4

• F 1s2 2s2 2p5 Ne 1s2 2s2 2p6

Filling Diagram for Sublevels

Order of filling orbitals

• 1s (with 2 electrons)

• 2s (2), 2p (6)

• 3s (2), 3p (6)

• 4s (2), 3d (10), 4p (6)

• 5s (2), 4d (10), 5p (6)

• 6s (2), 4f (14), 5d (10), 6p (6)

Practice: Give the electron configuration for:

P

1s2 2s2 2p6 3s2 3p3

Mn

1s2 2s2 2p6 3s2 3p6 4s2 3d5

Br

Al

Electron configuration and the Periodic Table

Abbreviated notation• When an energy level is completely filled we

often use an abbreviated notation with the noble gas configuration of the last filled period representing the inner electrons.

• Example: Na

• 1s22s22p63s1 or [Ne]3s1

Practice

• Give the abbreviated electron configuration of the following elements:

• S

• Co

• I

Electron configuration of Cu

• Cu: 1s2 2s2 2p6 3s2 3p6 4s2 3d9

• However, Cu is 1s2 2s2 2p6 3s2 3p6 4s1 3d10

• It is energetically slightly favorable for Cu to completely fill the 3d orbital, so one electron is moved from the 4s to the 3d orbital.

Electron configuration of Cr

• Cr shows a similar electron configuration effect as Cu.

• Cr is 1s2 2s2 2p6 3s2 3p6 4s1 3d5 rather than

1s2 2s2 2p6 3s2 3p6 4s2 3d4

• Due to the fact that a half-filled 3d orbital is energetically favorable over a filled 4s orbital.

Valence electrons

• These electrons are in the outermost principal energy level of an atom: the s and p electrons beyond the noble gas core.

• These electrons are involved in forming bonds with other atoms

• Inner electrons (core electrons) are NOT involved in bonding

Electron dot structure

• Elements (except helium) have the same # of valence electrons as their group #.

• Electron dot structures are used to show valence electrons.

• We use one dot for each valence electron.• Consider phosphorus, P, which has 5 valence

electrons. Here is the method for writing the electron dot formula.