How did we discover electron arrangement

in an atom?

ELECTROMAGNETIC

RADIATION ! ! !

Waves

Repeated disturbance through a medium (air, liquid) from origin to distant points.

Medium does not move

Ex. Ocean waves, sound waves

Characteristics of Waves

Wavelength Distance between 2 points within a wave cycle 2 peaks

Frequency # of wave cycles passing a point for a

particular time unit Usually seconds.

Wavelength and frequency are inversely proportional.

c = νλ

c = speed of light, 3.0 x 10 8 m/s Constant

ν= frequency (s-1 or Hz)

λ= wavelength (m)

Example 1:

Find the frequency of a green light that has a wavelength of 545 nm.

Electromagnetic Waves

Produced from electric charge movement

Changes within electric and magnetic fields carried over a distance

No medium needed

Electromagnetic Spectrum

Contains full range of wavelengths and frequencies found with electromagnetic radiation Wavelength/frequency changes cause color

changes

Mostly invisible, visible range (390 nm -760 nm)

Different materials absorb/transmit the spectrum differently.

Types of Spectra What is a spectra?

Spectrum– white light/radiation split into different wavelengths and frequencies by a prism

Continuous spectrum No breaks in spectrum Colors together

Line spectrum Line pattern emitted by light from excited atoms of

a particular element Aided in determining atomic structure

Line Spectrum

Pattern emitted by light from excited atoms of an element

Specific for each element Only certain wavelengths of visible spectrum

present

Used for element identification

Flame Tests

Some atoms of elements produce visible light if heated

Each element has a specific flame color

Examples: Li, Na, Cs, Ca

A Bit of Quantum Theory……

Max Planck 1900

Related energy and radiation E = hν h= 6.626 x 10 -34 Js (Planck’s constant) E = energy per photon (J)

Quantum---smallest amount of energy

Atoms can only absorb/emit specific quanta

Albert Einstein 1905

Added to Planck’s concept

Photons— Bundles of light energy Same energy as quantum E = hν (energy of photon)

Photons release energy and electrons gain energy Threshold frequency– minimum amount of energy

needed by photon to extract electron

THEREFORE ………

Light is in the form of electromagnetic waves

Photons can resemble particles

Gave raise to the possibility of thinking about wave AND particle qualities of subatomic particles (electron)

Example 1

Calculate the energy found in a photon of red light with a wavelength of 700.0 nm

Example 2

How much energy (in joules) is found in the radiation of the hydrogen atom emission spectrum with a 656.3 nm wavelength?

Example 3:

A sodium atom emits yellow light with a wavelength of 589 nm when it is excited. Find the energy per photon of this light.

Quantum Theory/Electron

Orbitals

Coulomb’s Law

Describes the attractive force between negative electrons and positive nucleus. Force is directly related to the charge of electron and

nucleus Force is inversely related to distance between particles

F = qe x qp

r2

(IE…an electron’s energy is dependent on distance from nucleus)

Early Models of the Atom Bohr

1913—hydrogen atom structure

Physics + quantum theory

Electrons move in definite orbits around the positively charged nucleus—planetary model

Does not apply as atoms increase in electron number

Electrons orbit nucleus in different energy levels

Lower energy levels, closest to nucleus (n = 1)

Higher energy levels increase electron’s distance from nucleus

Electrons can “transition” or jump between energy levels through photons Gain/absorb photon—higher energy level Lose/emit photon—lower energy level

Bohr Model

Energy States in an Atom

Atoms can gain or loss energy.

Specific energy states within an atom. Can be counted Ground State = lowest energy state Excited State = higher energy level

than ground, gained energy

So, where does the Bohr Model fit in?

Electrons orbit around the nucleus at different energy levels/orbits.

Electron’s energy level = orbit level where electron is located.

Light absorption = electron moves from a state of low energy to high energy. “becomes excited”

Light Emitted = electron falls from an “excited” state of energy to a lower energy level.

Ex. Li

Erwin Schrödinger

Quantum mechanics

1926---wave equation

Electrons behave more like waves than particles

Heisenberg’s Uncertainty Principle Electron’s location and direction cannot be

known simultaneously

Electron as cloud of negative charge

Modern Model of the AtomThe electron cloud

Sometimes called the wave model

Electron as cloud of negative charge

Spherical cloud of varying density

Varying density shows where an electron is more or less likely to be

Treats electron’s location as wave property

Defined by quantum numbers

Quantum numbers Provide information about size, shape, and

orientation of atomic orbitals Define atomic orbitals from general to specific

Quantum Theory

Determines orbital size and electron energy

Same as “n” value/orbital in Bohr model

Positive whole number, NOT 0

Shells – orbitals with same value

n = 1, 2, 3, 4, etc.

Principal Quantum Number (n)

Defines orbital shape for a particular region of atom

Think as “subshell”

l = n-1

# of orbitals/subshells = principal quantum #

Orbital Angular Momentum Quantum

Number (l)

l Orbital/Subshell

0 s

1 p

2 d

3 f

Describes orbital orientation within an atom

Range from –l to +l, 0 is possible

ml = 0, ± 1, ± 2, etc.

ml = 2l + 1 (number of orientations)

Magnetic Quantum Number (ml)

2p

4f

How do you specify orbitals?

s orbital 1 possible orbital orientation, spherical shape n value determines size Charge cloud found near center, likely electron location

p orbital 3 possible orbital orientations, dumbbell shape pX, py, pz

Orbital Shapes

P

What does atomic structure REALLY look

like?

Describes the motion of an electron, spinning

As electron moves, magnetic field induced

Electrons with opposite spins, cancel magnetic field of other

Values: +1/2, -1/2

Electron Spin

Read lab procedure

Read pp. 267-289 (for Friday)

Homework