Electrons in Atoms

Wavelength () - length of one complete wave measured in m, cm, or nm

In light it tells us which color it is

Frequency () - # of waves that pass a point during a certain time period, hertz (Hz) = 1/s

Amplitude (A) - distance from the origin to the trough (dip) or crest (peak Amount of energy the wave is carrying - height of wave. It

is measured in meters. In SOUND it tells us how LOUD it is. In LIGHT it tells how

BRIGHT it is.

A

greater amplitude

greater frequency

crest

origin

trough

A

Understanding electronic structure of atoms

Must understand light

Emitted and absorbed by substances. Visible light - type of Electromagnetic Radiation (EM)

Carries (radiant) energy through space Travels at speed of lightExhibits wavelike behavior.

Think of light as particlehelp understand how EM radiation and atomsinteract

LOW

ENERGY

HIGH

ENERGY

Move through a vacuum at the

‘speed of light’ 3.00 x 108 m/s

Behaves like waves that move through water

Result of energy transferred to the water (from

a stone)

Expressed as up and down movement

Both electric and magnetic properties

Wave Speed = (distance between wave peaks) x (frequency)

= (wavelength)

x (frequency)

EM radiation moves through a vacuum at the “speed of light”

3.00 x 108 m/s also called c. A lower energy wave (infrared and red) has a longer

wavelength() and lower frequency(f) 

A higher energy wave (blue - violet) has a shorter

wavelength() and higher frequency(f).

Frequency & wavelength are inversely proportional

c = c: speed of light (3.00 108 m/s): wavelength (m, nm, etc.): frequency (Hz)

EX: Find the frequency of a photon with a wavelength of 434 nm.

GIVEN:

= ?

= 434 nm = 4.34 10-7 m

c = 3.00 108 m/s

WORK: = c

= 3.00 108 m/s 4.34 10-7 m

= 6.91 1014 Hz

Planck (1900)

Observed - emission of light from hot objects

Concluded - energy is emitted (absorbed or released) in small, specific amounts (quanta)

Quantum - smallest energy packet that can be emitted or absorbed as EM radiation by an atom.

E: energy (J, joules)h: Planck’s constant (6.6262 10-34 J·s): frequency (Hz)

E = h

Planck proposed that the energy, E, of a single quantum energy packet equals a constant (h) times its frequency

The energy of a photon is proportional to its frequency.

EX: Find the energy of a red photon with a frequency of 4.57 1014 Hz.

GIVEN:

E = ? = 4.57 1014 Hzh = 6.6262 10-34 J·s

WORK:

E = h

E = (6.6262 10-34 J·s)(4.57 1014 Hz)

E = 3.03 10-19 J

Energy is always emitted or absorbed in whole number multiples of hv, such as hv, 2 hv, 3 hv, 4hv, ….

The allowed energies are quantized values are restricted to certain quantities.  

The notion of quantized rather than continuous energies is strange.

Ramp vs Staircase Ramp - vary the length your steps and energy used on

the walk up. Stairs - must exert exactly the specific amount of

energy needed to reach the next step. Your steps on steps are quantized, you cannot step

between them.

Planck (1900)

vs.

Classical Theory Quantum Theory

Einstein (1905)

Observed – photoelectric effect

Dispersed light falls on metal samples, the different frequencies produce different energetic photoelectrons

Einstein (1905)

Concluded - light has properties of both waves and particles (photons)

“wave-particle duality”

Photon - particle of light that carries a quantum of energy

Used planck’s quantum theory to deduced that: Ephoton = hv

Electrons in Atoms

ground state

excited state

ENERGY IN PHOTON OUT

Elements’ atoms absorb electrical energy

e- get excited, become unstable, and release energy

Energy is in form of light

Set of frequencies of EM waves

e- exist only in orbits with specific amounts of energy called energy levels

Therefore…

e- can only gain or lose certain amounts of energy

only certain photons are produced

Ground state = lowest allowable atomic electron energy state

Excited state = any higher energy state

Energy of photon depends on the difference in energy levels

Bohr’s calculated energies matched the IR, visible, and UV lines for the H atom

1

23

456

Each element has a unique bright-line emission spectrum. “Atomic Fingerprint” Helium

Examples: Iron

Now, we can calculate for all elements and their electrons – next section

Louis de Broglie (1924)

Proposed eˉ in their orbits behave like a waveEVIDENCE: DIFFRACTION PATTERNS

ELECTRONSVISIBLE LIGHT

Heisenberg Uncertainty Principle Impossible to know velocity and position of an

electron at the same time

Trying to observe an electron’s position changes its momentum

Trying to observe an electron’s momentum changes its position

Electrons cannot be locked into well-defined circular orbits around the nucleus.

Orbital (“electron cloud”) a specific distribution of electron density in space. Each orbital has a characteristic energy and shape.

Orbital

Specify the “address” of each electron in an atom

UPPER LEVEL

1. Principal Quantum Number (n = 1, 2, 3, …) (see periodic table left column)

Indicates the relative size and energy of atomic orbitals

As (n) increases, the orbital becomes larger, the electron spends more time farther from the nucleus

Each major energy level is called a principle energy level Ex: lowest level = 1 ground state, highest level = 7 excited state

2. Energy Sublevel Defines the shape of the orbital (s, p, d, f)

# of orbital related to each sublevel is always an odd # s = 1, p = 3, d = 5, f = 7

Each orbital can contain at most 2 electrons

s p d f

Subscripts x, y, z designates orientation Specifies the exact orbital within each

sublevel

px py pz

4. Spin Quantum Number ( ms ) Electron spin +½ or -½

An orbital can hold 2 electrons that spin in opposite directions.

Pauli Exclusion Principle A maximum of 2 electrons can occupy a single

atomic orbital

Only if they have opposite spins

1. Principal #

2. Energy sublevel 3. Orientation

4. Spin #

energy level

(s,p,d,f)

x, y, z

exact electron

Electron Configuration

Electrons in Atoms

A. General Rules

Aufbau Principle

Electrons fill the lowest energy orbitals first.

“Lazy Tenant Rule”

RIGHTWRONG

A. General Rules

Hund’s Rule

Within a sublevel, place one e- per orbital before pairing them.

“Empty Bus Seat Rule”

O

8e-

Orbital Diagram

Electron Configuration

1s2 2s2 2p4

B. Notation

1s 2s 2p

Shorthand Configuration

S 16e-

Valence ElectronsCore Electrons

S 16e- [Ne] 3s2 3p4

1s2 2s2 2p6 3s2 3p4

B. Notation

Longhand Configuration

Valence electrons: determine chemical properties of that

element & are the electrons in the atoms outermost orbital

© 1998 by Harcourt Brace & Company

sp

d (n-1)

f(n-2)

1234567

67

Notation

1

2

3

4

5

6

7

Shorthand Configuration Core e-: Go up one row and over to the

Noble Gas. Valence e-: On the next row, fill in the #

of e- in each sublevel.

Shorthand Notation

[Ar] 4s2 3d10 4p2

C. Periodic Patterns

Example - Germanium

Full energy level

1

2

3

4 5

6

7

Full sublevel (s, p, d, f)Half-full sublevel

D. Stability

Electron Configuration Exceptions

Copper

EXPECT: [Ar] 4s2 3d9

ACTUALLY: [Ar] 4s1 3d10

Copper gains stability with a full d-sublevel.

D. Stability

Electron Configuration Exceptions

Chromium

EXPECT: [Ar] 4s2 3d4

ACTUALLY: [Ar] 4s1 3d5

Chromium gains stability with a half-full d-sublevel.

D. Stability

D. Stability

Ion Formation Atoms gain or lose electrons to become

more stable. Isoelectronic with the Noble Gases.

O2- 10e- [He] 2s2 2p6

D. Stability

Ion Electron Configuration

Write the e- config for the closest Noble Gas

EX: Oxygen ion O2- Ne