I. Waves & Particles Ch. 4 - Electrons in Atoms. A. Waves zWavelength ( ) - length of one complete wave zFrequency ( ) - # of waves that pass a point.

I. Waves & Particles

Ch. 4 - Electrons in Atoms

A. Waves

Wavelength () - length of one complete wave

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 or crest

A. Waves

Agreater

amplitude

(intensity)

greater frequency

(color)

crest

origin

trough

A

B. EM Spectrum

LOW

ENERGY

HIGH

ENERGY

B. EM Spectrum

LOW

ENERGY

HIGH

ENERGY

R O Y G. B I V

red orange yellow green blue indigo violet

B. EM Spectrum

Frequency & wavelength are inversely proportional

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

B. EM Spectrum

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

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

C. Quantum Theory

Planck (1900)

Observed - emission of light from hot objects

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

Quantum - minimum amount of energy change

C. Quantum Theory

Planck (1900)

vs.

Classical Theory Quantum Theory

C. Quantum Theory

Einstein (1905)

Observed - photoelectric effect

C. Quantum Theory

Einstein (1905)

Concluded - light has properties of both waves and particles

“wave-particle duality”

Photon - particle of light that carries a quantum of energy

C. Quantum Theory

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

E = h

The energy of a photon is proportional to its frequency.

C. Quantum Theory

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

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

II. Bohr Model of the Atom

Ch. 4 - Electrons in Atoms

A. Line-Emission Spectrum

ground state

excited state

ENERGY IN PHOTON OUT

B. Bohr Model

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

B. Bohr Model

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. Other Elements

Each element has a unique bright-line emission spectrum.

“Atomic Fingerprint”

Helium

Bohr’s calculations only worked for hydrogen!

III. Quantum Model

of the Atom

Ch. 4 - Electrons in Atoms

A. Electrons as Waves

Louis de Broglie (1924)

Applied wave-particle theory to e-

e- exhibit wave properties

QUANTIZED WAVELENGTHS

A. Electrons as Waves

QUANTIZED WAVELENGTHS

A. Electrons as Waves

EVIDENCE: DIFFRACTION PATTERNS

ELECTRONSVISIBLE LIGHT

B. Quantum Mechanics

Heisenberg Uncertainty Principle

Impossible to know both the velocity and position of an electron at the same time

B. Quantum Mechanics

σ3/2 Zπ

11s 0

eΨ a

Schrödinger Wave Equation (1926)

finite # of solutions quantized energy levels

defines probability of finding an e-

B. Quantum Mechanics

Radial Distribution CurveOrbital

Orbital (“electron cloud”)

Region in space where there is 90% probability of finding an e-

C. Quantum Numbers

UPPER LEVEL

Four Quantum Numbers:

Specify the “address” of each electron in an atom

C. Quantum Numbers

1. Principal Quantum Number ( n )

Energy level

Size of the orbital

n2 = # of orbitals in the energy level

C. Quantum Numbers

s p d f

2. Angular Momentum Quantum # ( l )

Energy sublevel

Shape of the orbital

C. Quantum Numbers

n = # of sublevels per level

n2 = # of orbitals per level

Sublevel sets: 1 s, 3 p, 5 d, 7 f

C. Quantum Numbers

3. Magnetic Quantum Number ( ml )

Orientation of orbital

Specifies the exact orbitalwithin each sublevel

C. Quantum Numbers

px py pz

C. Quantum Numbers

Orbitals combine to form a spherical

shape.

2s

2pz2py

2px

C. Quantum Numbers

4. Spin Quantum Number ( ms )

Electron spin +½ or -½

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

C. Quantum Numbers

1. Principal #

2. Ang. Mom. #

3. Magnetic #

4. Spin #

energy level

sublevel (s,p,d,f)

orbital

electron

Pauli Exclusion Principle

No two electrons in an atom can have the same 4 quantum numbers.

Each e- has a unique “address”:

Feeling overwhelmed?

Read Section 4-2!

IV. Electron Configuration(p. 105 - 116,

128 - 139)

Ch. 4 - Electrons in Atoms

A. General Rules

Pauli Exclusion Principle

Each orbital can hold TWO electrons

with opposite spins.

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 Electrons

Core Electrons

S 16e- [Ne] 3s2 3p4

1s2 2s2 2p6 3s2 3p4

B. Notation

Longhand Configuration

© 1998 by Harcourt Brace & Company

sp

d (n-1)

f (n-2)

1234567

67

C. Periodic Patterns

C. 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

1s11st column of s-block

C. Periodic Patterns

Example - Hydrogen

1

2

3

4

5

6

7

C. 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] 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

Ch. 5 - The Periodic Table

I. History

A. Mendeleev

Dmitri Mendeleev (1869, Russian) Organized elements

by increasing atomic mass.

Elements with similar properties were grouped together.

There were some discrepancies.

A. Mendeleev

Dmitri Mendeleev (1869, Russian) Predicted properties of undiscovered

elements.

B. Moseley

Henry Mosely (1913, British)

Organized elements by increasing atomic number.

Resolved discrepancies in Mendeleev’s arrangement.

II. Organization of theElements

Ch. 5 - The Periodic Table

MetalsNonmetalsMetalloids

A. Metallic Character

Main Group ElementsTransition MetalsInner Transition Metals

B. Blocks

III. Periodic Trends

Ch. 5 - The Periodic Table

0

50

100

150

200

250

0 5 10 15 20Atomic Number

Ato

mic

Ra

diu

s (

pm

)

A. Periodic Law

When elements are arranged in order of

increasing atomic #, elements with similar

properties appear at regular intervals.

0

50

100

150

200

250

0 5 10 15 20Atomic Number

Ato

mic

Ra

diu

s (

pm

)

B. Chemical Reactivity

Families Similar valence e- within a group result in

similar chemical properties

B. Chemical Reactivity

Alkali MetalsAlkaline Earth MetalsTransition MetalsHalogensNoble Gases

Atomic Radius size of atom

© 1998 LOGALFirst Ionization Energy

Energy required to remove one e- from a neutral atom.

© 1998 LOGAL

Melting/Boiling Point

C. Other Properties

Atomic Radius

0

50

100

150

200

250

0 5 10 15 20Atomic Number

Ato

mic

Ra

diu

s (

pm

)

D. Atomic Radius

Li

ArNe

KNa

1

2

3

4 5

6

7

Atomic Radius Increases to the LEFT and DOWN

D. Atomic Radius

Why larger going down?

Higher energy levels have larger orbitals

Shielding - core e- block the attraction between the nucleus and the valence e-

Why smaller to the right?

Increased nuclear charge without additional shielding pulls e- in tighter

D. Atomic Radius

First Ionization Energy

0

500

1000

1500

2000

2500

0 5 10 15 20Atomic Number

1s

t Io

niz

ati

on

En

erg

y (k

J)

E. Ionization Energy

KNaLi

Ar

NeHe

1

2

3

4 5

6

7

First Ionization Energy Increases UP and to the RIGHT

E. Ionization Energy

Why opposite of atomic radius?

In small atoms, e- are close to the nucleus where the attraction is stronger

Why small jumps within each group?

Stable e- configurations don’t want to lose e-

E. Ionization Energy

Successive Ionization Energies

Mg 1st I.E. 736 kJ

2nd I.E. 1,445 kJ

Core e- 3rd I.E. 7,730 kJ

Large jump in I.E. occurs when a CORE e- is removed.

E. Ionization Energy

Al 1st I.E. 577 kJ

2nd I.E. 1,815 kJ

3rd I.E. 2,740 kJ

Core e- 4th I.E. 11,600 kJ

Successive Ionization Energies

Large jump in I.E. occurs when a CORE e- is removed.

E. Ionization Energy

1

2

3

4 5

6

7

Melting/Boiling Point Highest in the middle of a period.

F. Melting/Boiling Point

Ionic Radius

Cations (+)

lose e-

smaller

© 2002 Prentice-Hall, Inc.

Anions (–)

gain e-

larger

G. Ionic Radius

Which atom has the larger radius?

Be or Ba

Ca or Br

Ba

Ca

Examples

Which atom has the higher 1st I.E.?

N or Bi

Ba or Ne

N

Ne

Examples

Which atom has the higher melting/boiling point?

Li or C

Cr or Kr

C

Cr

Examples

Which particle has the larger radius?

S or S2-

Al or Al3+

S2-

Al

Examples

Periodic Trends Summary