Atomic Structure Part II

Electrons in Atoms

Radiant energy travels in

the form of waves that

have both electrical and

magnetic properties.

These electromagnetic

waves can travel through

empty space, as you know

from the fact that radiant

energy from the sun

travels to Earth every day.

Electrons in Atoms

I. Properties of Waves

1. Definition:

Energythat exhibits wave-like(or

oscillating) behavior as it travels

through space

Electrons in Atoms

2. Wavelength (λ λλλ) distance from peak

to peak, leng

th of

one com

plete

wave

3. Frequency (ν ννν)

a. number of peaks that pass at a given

point each sec

b. can be called cycles per

second (peak/sec)

c. cps now called 1 Hertz (Hz)

Electrons in Atoms-Cont.

4. Velocity (C = speed of light)

a. distance a given peak moves in

a unit of time

b. velocity (m/s) = frequency x

wavelength

c = ν ννν

x λ λλλ

II. Behavior of Light

A.Newton (1600) thought light consisted of particles

(beam of light is a stream of particles)

B. Maxwell (1864) thought light was a wave

phenomenon.

•Calcu

late

d the velocity of th

epr

opag

ation

of anelectro

magn

etic wave

and fou

nd it

was the sam

e for

light

II. Behavior of Light

1. some say light is like waves, some say

its like particles

2. modern theory says that it behaves as

both "wave/particle duality"

II. Behavior of Light

3. Max Planck (early 1900's) said:

a. light is made up of bundles of energy called

photons (or quanta)

b. the energy of each photon is proportional to

the frequency of the light

(Quantum Theory)

Quantum is the minimum amount of

energy that can be gained or lost by

an atom.

•example: CONTINUOUS SPECTRUM

*** when white light is passed through

a prism, it is separated into a band of

colors from red �violet. It's called a

continuous spectrum

c. the work of Planck & Einstein led to

E=energy,

ν= frequency,

h=planks constants (6.6262x10-34J /sec) J is

the symbol for joule the SI unit for energy

E = h x ν ννν

Energy of a quantum is related to the

frequency of the emitted radiation by

this equation

c. the work of Planck & Einstein led to

According to Planck’s theory, for a given

frequency matter can emit or absorb energy

only in whole-number mulitplesof hv, that is

1hv, 2hv, 3hv, and so on.

E = h x ν ννν

c. The photoelectric effect

In the photoelectric effect, electrons, called

photoelectrons, are emitted from a metal’s

surface when light of a certain frequency shines

on the surface. (exam

ple: solar calculator)

c. The photoelectric effect

Einstein said light can both wavelike and

particle like natures. That is, while a beam of

light has many wavelike characteristics, it

also can be thought of as a stream of tiny

particles, or bundles or energy, called

photons.

APhotonis a particle of electromagnetic

radiation with no mass that carries a

quantum of energy.

III. Bright line spectrum

•A. a spectrum that shows separate bright

lines, each with a specific wavelength

•B. bright-line spectra occur when an

element is heated and the colored light

given off is viewed through a

spectroscope. Each element has a unique

set of lines, characteristic of that element

(like a fingerprint)

Line-Emission Spectrum

ground state

excited state

ENERGY IN

PHOTON OUT

Fireworks? Hmmm…

IV. Electromagnetic Spectrum

•A. visible light (like the continuous spectrum)

is only one type of radiation. All other types

are not visible to the human eye.

H I G H E N E R G Y

L O W E N E R G Y

Electromagnetic Spectrum

L O W E N E R G Y

H I G H E N E R G Y

RO

YG.

BI

V

red

orange

yellow

green

blue

indigo

violet

Electromagnetic Spectrum

B. all forms of electromagnetic radiation

travels at the speed of light.

1. speed of light = 3.00 x 108meters/sec

2. use form

ula:

c = ν νννx λ λλλ

3. each line spectrum has a particular

frequency (ν ). If know wavelength (λ), we

can find ν

using c

as a constant.

C. The energy in a photon of light is directly

proportional to the frequency of the light.

•1.

frequency,

energy

•2. can find the energy of a single

quantum (photon) of radiation at

any given frequency.

C. The energy in a photon of light is directly

proportional to the frequency of the light.

•3. proportionality constant that

relates the two is called Planck's

constant (h).

•4. form

ula:

E = hx f

example: a spectral line has frequency of

3.5x1012hertz. What is the energyof a

photon of radiation of this frequency?

E = hx f

(h=6.6262x10-34J/sec)

E = (3.5x1012Hz) (6.6262x10-34J •sec)

E = (2.3x10-21J)

V. Electron energy levels in

Bohr's Model

A. There are certain different

orbits in which an electron can

travel around a nucleus.

1. each circular orbit (or shell) is at a

fixed distance from the nucleus

V. Electron energy levels in

Bohr's Model

2. the greater the radius of that shell,

the greater the energy of the

electron in that shell.

3. these electron orbits are known as

energy levels

B. When electrons absorb energy firm an

outside source, they jump from lower to

higher energy levels.

�when they fall back to their original levels ,

energy is

emitted (light); the same amount as was

absorbed.

1

23

456

�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

C. In energy atom in its norm

al

state, all electrons are in the

lowest energy levels available

(energetically stable)

VI. Atoms and

Radiation

•A. When all of

the lowest

energy levels

are occupied,

the atom is in

the ground

state

(unexcited).

VI. Atoms and

Radiation

•B. When

electron moves

to higher

energy level,

atom is in the

excited state,

and is

energetically

unstable.

VI. Atoms and

Radiation

C. Bright line

spectrum of an

element

represents the

energy levels in

its atoms.

�problems with Bohr's Model:

� ���only explained some of the lines in the

bright line spectrum

� ���really only worked for hydrogen

� ���need sublevels and electron cloud

model to account for all of the lines.

VII. The Modern Model of the Atom

A. Mechanics

1. Classical Mechanics -Newton's Laws of

Motion (Newtonian Mechanics)

Describes the behavior of visible objects

traveling at ordinary velocities.Bohr’s basis

for his model, but couldn’t explain why

electrons would stay at on energy level or

another. When looking at H-spectral lines,

noticed more one (several closely spaced).

VII. The Modern Model of the Atom

2. Quantum Mechanics –(wave

mechanics)

Describes the behavior of extrem

ely

small particles traveling at velocities at

or near the speed of light

a. Louis de Broglie -particles could

have properties of waves

Planks quanta gave wave properties,

deBrogliesaid electron streams are like

waves of light and have properties of

both particles and waves (matter behaves

as waves)

•b. Schrodinger -described the

behavior of electrons in terms of

quantized energy changes

"quantum mechanics"

Describe a wave equation used to

determine the probability of finding an

electron in any give place or orbital

Schrodinger’sCat

Radial Distribution Curve

Orbital

c. Heisenberg -uncertainty principle

�Region of space where there is a

probability of finding an electron is

called an orbital

"The more

precisely

the POSITIO

Nis

determined,

the less precisely

the

MOMENTUM is

known"

B. Principal Energy Levels

1. Energy Levels

•Bohr -

High Energy

(outer level)

Low Energy

1 2 3 4

Principal Quantum Numbers (N)Number of electrons

2

8 18

32

Corresponds to

energy level

2. Sublevels

Principal Quantum Numbers (N) Sublevel

Present

11s

22s2p

33s3p3d

44s4p4d4f

•Orbital -

–Region of space where an

electron is probably found

•Electron spin

–An orbital can hold 2 electrons that spin in

opposite directions.

Electrons are

represented by arrows

Rules: 1. PauliExclusion Principle

–Each orbital can hold TWO

electrons with opposite

spins.

–No two electrons in an atom

can have the same 4

quantum numbers.

–Each e

-has a unique

“address”:

2. AufbauPrinciple

Electrons fill the

lowest energy

orbitalsfirst.

Electrons to be

added must be

placed in unfilled

orbitalsof lowest

energy for stable

configuration.

3. Hund’sRule

–Within a sublevel, place one e-

per orbital before pairing them.

–“Empty Bus Seat Rule” RIGHT

WRONG

Energy Level Diagram

orbital

Orbital -Place where electrons are probablyfound

Electrons have “up”and

“down”spin

c

Shapes of electron orbitals

The s orbital

The p orbitals

py

px

The p orbitals

pz

The d

The d orbitals

orbitals

The d

The d orbitals

orbitals

The f

The f orbitals

orbitals

f

Click here for orbital viewer

View the grand table

orbital views

C. Electron Configurations

1s2= Helium

1s22s1= Lithium

1s22s22p63s23p6

4s23d104p6

= Krypton

Energy Level

Sub Level # of Electrons

O 8e-

•Orbital Diagram

•Electron Configuration

1s2 2s22p4

Electron Configuration Notation

1s

2s

2p

•Shorthand Configuration

S 16e-

Valenc

e E

lectro

nsCor

e E

lectro

ns

S16e-[Ne]3s2 3p4

1s22s22p63s23p4

Notation

•Longhand Configuration

©1998 by Harcourt Brace & Company

sp

d (n-1)

f (n-2)

1 2 3 4 5 6 7

6 7

Periodic Patterns

Periodic Patterns

•Period #

–energy level (subtract for d & f)

•A/B Group #

–total # of valence e

-

•Column within sublevel block

–# of e-in sublevel

s-block

1st Period

1s1

1st colum

n of

s-block

1

2

3

4

5

6

7

Periodic Patterns

•Example -Hydrogen

1

2

3

4

5

6

7

Periodic Patterns

•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.

[Ar]1

2

3

4

5

6

7

4s23d104p2

Periodic Patterns

•Example -Germ

anium

•Full energy level

1

2

3

4

5

6

7

•Full sublevel (s, p, d, f)

•Half-full sublevel

Stability

1

2

3

4

5

6

7

Stability

•Ion Formation

–Atoms gain or lose electrons to become more

stable.

–Isoelectronicwith the Noble Gases.

1

2

3

4

5

6

7

Stability

•Ion Formation

–Atoms gain or lose electrons to become more

stable.

–Isoelectronicwith the Noble Gases.

Feeling overwhelmed?

Try a few!

Mg =

Fe =

Ru=

Ir=

Ca+2 =

Cl-1 =