Electrons and Quantum Mechanics

Unit 5

Electrons

• Rutherford described the dense center of the atom called the nucleus.

• But the Electrons spin around the outside of that nucleus.– Provide the chemical

properties of the atoms.– Responsible for color and

reactivity.

Energy

• Energy is transmitted from one place to another.– Light carries this energy.– Converted into heat.

• Light is called Electromagnetic Radiation.

Electromagnetic Spectrum

• Radio• Infrared• Visible Light– ROY G BIV

• Ultraviolet• X Rays• Gamma Rays

Light

• Light travels as a wave.• Wave Properties– Wavelength (λ) =

distance between two waves (m)

– Frequency (f) = number of peaks per second (Hz)

– Speed of Light (c) = how fast light moves.

Light

• Light Equation

c= ƒλ• Speed of light is a

constant = 3 x 108 m/s• Nothing travel faster

than the speed of light!– Maybe?!?!?!?!?!?!?!?!?

Light

• The Dual Nature of Light– Light carries energy

through space like a wave.

– Light also behaves like a particle?!?• A beam of light is made of

tiny packets of energy called PHOTONS!

• Which travel in waves!?!

Light• The Energy of a photon

depends on its frequency.– So is the color of light!!!

E = hĥ ELECTRONS are like

photons!– Act as waves and particles.– Orbit the nucleus in a

wave-like motion.

Blackbody Radiation

• Rutherford could never explain why objects change colors when they are heated.

• As the object heats, it must give off electrons of certain frequencies and energies.

Photoelectric Effect

• Similarly, light on a metal object can knock off electrons.– Shine different colors on

a metal.– Measure the number of

electrons knocked off.– Found that no electrons

were knocked off below a certain frequency.

The Bohr Model

• Proposed the electrons orbit the nucleus with fixed energies.– Called Energy Levels– Much like the rungs of a

ladder.• Quantum describes the

amount of energy required to move an electron from one level to another.

The Bohr Model

• Ground State– Lowest possible energy

of an electron.– Normal location

• Excited State– If electron absorbs

energy, it moves up an energy level (absorption)

– If an electron gives off energy, it moves down an energy level (emission).

The Bohr Model

Atomic Spectra

• Hydrogen Atom Line Emission Spectrum– Expected continuous

spectrum of light– But only specific

frequencies were given off.• Red (656.6 nm)• Blue-green (486.1 nm)• Violet (434.1 nm)• Violet (419.2 nm)

Atomic Spectra• Shine a light on an Atom

– When atoms absorb energy, electrons move to higher energy levels.

– When atoms release the energy, electrons return to the lower energy level.

• Atomic Spectra– Frequencies of light emitted

by a certain element.– No two elements have the

same spectrum.

http://student.fizika.org/~nnctc/spectra.htm

Flame Tests

• Because no two atoms produce the same spectrum, elements can be identified by the colors they emit.

• Spectral Analysis uses this properties to identify elements.

Quantum Mechanics

• Max Planck (1900)– Founder of Quantum

Mechanics

E = hf• Albert Einstein (1905)– Wave-Particle Duality– Electrons are small

particles that move like waves.

Quantum Mechanics

• Neils Bohr (1922)– Electrons orbit in distinct

energy levels.• Louis de Brogelie (1923)– Wave Mechanics says

that ALL MATTER behaves like waves.

mv/λ = h

Quantum Mechanics

• Werner Heisenberg (1927)– Principle of Indeterminacy– You can’t know both the

position and the velocity of an electron.

• Erwin Schrödinger (1930)– Used wave mechanics to

show the PROBABLE location of an electron.

– Electrons exist in 3D clouds of probability!!!

Quantum Mechanical Model

• Uses Schrodinger’s equation to predict the probable location of an electron.– Determines the energies

an electron is allowed to have.

– Determines how likely it is to find the electron in various locations around the nucleus.

Quantum Numbers

• Describes the location and behavior of an electron– Like an electron’s

address– No two electrons can

have the same quantum numbers.

• Four Numbers

Quantum Numbers• Principle (1st) Quantum

Number (n)– The Energy Level– Describes the size of the

cloud and the distance of the cloud from the nucleus.

– Shows the number of electrons

n = 1 = 2 electronsn = 2 = 8 e-

n = 3 = 18 e-

n = 4 = 32 e-

Quantum Numbers

• 2nd Quantum Number (l)– Each energy level has

sublevels.– The number of sublevels

equals n.– Sublevels are called:

s = sphericalp = peanut-shapedd = daisy-shaped

f = unknown?

Quantum Numbers

• 3rd Quantum Number (ml)– Divides sublevels into orbitals.– Tells the shape the electron

moves in.– Number of orbitals = n2

– Examples

s = 1 orbitalp = 3 orbitalsd = 5 orbitalsf = 7 orbitals

Quantum Numbers

• 4th Quantum Number (ms)– Describes the electron’s

spin.– Only two electrons fit in an

orbital.– Their charges repel causing

them to spin in opposite directions (+½ or –½)

– Use up and down arrows.

Quantum Numbers

• Pauli Exclusion Principle– No two electrons can

have the same set of 4 quantum numbers.

– The electrons repel each other.

• Hund’s Rule– Every orbital must get

one electron before doubling up.

Quantum Numbers

• Diagonal Rule– Electrons fill orbitals in

predictable patterns– Some People Do Forget– Electrons dill the lowest

energy level possible.

1s2s 2p3s 3p

3d4s 4p4d 4f5s 5p5d 5f

Orbital Notation

• Draw out the locations of each electron in an atom with arrows.

Electron Configuration

• Write out the configurations of electrons using superscripts.

• Examples:– H = 1s1

– He = 1s2

Electron Configurations

• Noble Gas Shorthand– Write the Noble Gas just

before the element.– Add the remainder of

the configuration.

Lewis Dot Diagrams• A way to show the number

and position of the valence electrons.– Outermost energy level– Look at the column number

to get this number.• Use the chemical symbol

and number of valence electrons.– All four sides must have a

dot before you double up.

X

p1

p3

p2 s

p orbitals s orbital