…and, in conclusion…

You will need to know the contributions of…

• Bohr

• Planck

• Einstein

• Heisenberg

• de Broglie

Q: How is light produced?


• A: An excited electron…

•



• loses energy…

•



• loses energy…

• as it falls…



• loses energy…

• as it falls…

• from a higher to a lower energy level,…

•



• loses energy…

• as it falls…


• emitting that energy…

•



• loses energy…

• as it falls…


• emitting that energy…

• as a photon of light

An electron that gains energy is excited:

moves to a higher energy level,

physically moves away from nucleus

An electron in its normal position is in its ground state

*

(gains energy)

Ground state

Excited *

• Light is produced*

As it loses energy

Ground state

Excited *

We use a * to mark an excited electron or atom.

• An atom can be excited by:

•

•

•

We use a * to mark an excited electron or atom.

• An atom can be excited by:

Heat (light bulbs, stars, sparks, flames)

Electricity (sparks, fluorescent bulbs) or

Chem. rxns (lightning bugs, glow sticks)

Let’s do the wave.

The wave equation

Where

c is the speed of light, 3.00 x 108 m/s

is the wavelength (the symbol, lambda, is the Greek “l” for length) and

is the frequency (the symbol, nu, is the Greek “n”)

c=

Please notice:

• Wavelength and frequency are inversely related.

• When wavelength increases, frequency decreases.

• Q: What is the frequency of light that has a wavelength of 570 nm(5.7x10-7m)?


• c=


• c==c/


• c==c/=(3.00x108m/s) / (5.7x10-7m)


• c==c/=(3.00x108m/s) / (5.7x10-7m)

• 5.3 x 1014 /s =5.3 x 1014 Hz

Next: Planck’s equation

• The energy of a photon is directly related to its frequency

E=h• where

• E is the energy (in Joules),

• is the frequency (in waves/s, or Hz), and

• h is the conversion factor, Planck’s constant.

h=6.63 x 10-34 Js

The electromagnetic spectrum

For any pair--

Which has greater , , E, v?

The electromagnetic spectrumIncreasing wavelength

Increasing frequency

Increasing energy

and, c is the velocity of light (c for constant!)

Visible light400 nm

700 nm

Short wavelength High energy High frequency

Long wavelength Low energy Low frequency

Fill in the missing information

Type of photon

c=

3x108m/s

(m) (s-1) E (per photon)

450 nm

93.3 MHz

353 m

9.47 x 10-21 J

4.03 x 10-19 J


Type of photon

c=

3x108m/s


3.00 x 108 m/s

450 nm 6.67 x 1014 Hz

4.42 x 10-19 J

93.3 MHz

353 m

9.47 x 10-21 J

4.03 x 10-19 J


Type of photon

c=

3x108m/s


3.00 x 108 m/s

450 nm 6.67 x 1014 Hz

4.42 x 10-19 J

3.00 x 108 m/s

3.22 m 93.3 MHz 6.19 x 10-26 J

353 m

9.47 x 10-21 J

4.03 x 10-19 J


Type of photon

c=

3x108m/s


3.00 x 108 m/s

450 nm 6.67 x 1014 Hz

4.42 x 10-19 J

3.00 x 108 m/s

3.22 m 93.3 MHz 6.19 x 10-26 J

3.00 x 108 m/s

353 m 8.50 x 105 Hz

5.63 x 10-28 J

9.47 x 10-21 J

4.03 x 10-19 J


Type of photon

c=

3x108m/s


3.00 x 108 m/s

450 nm 6.67 x 1014 Hz

4.42 x 10-19 J

3.00 x 108 m/s

3.22 m 93.3 MHz 6.19 x 10-26 J

3.00 x 108 m/s

353 m 8.50 x 105 Hz

5.63 x 10-28 J

3.00 x 108 m/s

2.10 x 10-5 m

1.43 x 1013 Hz

9.47 x 10-21 J

4.03 x 10-19 J


Type of photon

c=

3x108m/s


3.00 x 108 m/s

450 nm 6.67 x 1014 Hz

4.42 x 10-19 J

3.00 x 108 m/s

3.22 m 93.3 MHz 6.19 x 10-26 J

3.00 x 108 m/s

353 m 8.50 x 105 Hz

5.63 x 10-28 J

3.00 x 108 m/s

2.10 x 10-5 m

1.43 x 1013 Hz

9.47 x 10-21 J

3.00 x 108 m/s

4.94 x 10-7 m

6.08 x 1014 Hz

4.03 x 10-19 J


Type of photon

c=

3x108m/s


Blue or indigo light

3.00 x 108 m/s

450 nm 6.67 x 1014 Hz

4.42 x 10-19 J

FM radio 3.00 x 108 m/s

3.22 m 93.3 MHz 6.19 x 10-26 J

AM radio 3.00 x 108 m/s

353 m 8.50 x 105 Hz

5.63 x 10-28 J

IR 3.00 x 108 m/s

2.10 x 10-5 m

1.43 x 1013 Hz

9.47 x 10-21 J

Blue or Green light

3.00 x 108 m/s

4.94 x 10-7 m

6.08 x 1014 Hz

4.03 x 10-19 J

The Bohr Model of the atom• Bohr’s solar system model

-- shows why the H gives off only 4 wavelengths of visible light.

• He calculated the energy that the electrons did give off, and the differences in energy.

Bohr drew a picture like this

Electron loses energy

Electron gains energy

• These 4. No more. There is nothing between the levels– no “half-transitions”

Visible wavelengths produced by

hydrogen atoms

In the hydrogen spectrum

• The wavelengths are:– 410 nm (violet)– 434 nm (blue)– 486 nm (green)– 656 nm (red)

4.85x10-19J

4.58x10-19J

4.09x10-19 J

3.03x10-19J


4.85x10-19J

4.58x10-19J

4.09x10-19 J

3.03x10-19J


What is the energy difference between these two levels?

4.85x10-19J

4.58x10-19J

4.09x10-19 J

3.03x10-19J


What is the energy difference between these two levels?

.49 x 10-19 J

• These energies represent the differences in the energy levels…

…and that’s how we know where the energy levels are.

Other phenomena that teach us about electrons and light

The photoelectric effect

– Described by Einstein for his Nobel prize– Light knocks electrons off a metal– Indicates the particle nature of

light

One photon excites one electron


• Heisenberg’s Uncertainty Principle

– “You cannot determine both the location and momentum of a particle exactly.”

– If you measure one, you change the other unpredictably

– Leads to the wave and particle natures of everything

Gratuitous joke

Heisenberg was pulled over on the highway. The officer asks,

“Do you know how fast you were going?”

Heisenberg replies,

Gratuitous joke

Heisenberg was pulled over on the highway. The officer asks,

“Do you know how fast you were going?”

Heisenberg replies,

“No, but I do know where I am!”


• DeBroglie’s wavelength

– Describes the wave nature of particles– Considers the uncertainty in position as a

wavelength

Electron Configurations

…and now, the rest of the story

Fe (2,8,14,2)

• --An electron configuration (EC) shows the location of all electrons in an atom or ion.

• --In an atom, number of electrons = number of protons = atomic number

• --Electrons are found around the nucleus of an atom in specific energy levels.

Levels have sublevels!

Levels have sublevels!

Sublevels have orbitals!

1st energy level

• -has one sublevel, 1s

• -An s sublevel (spherical) has one orbital,

• -An orbital can hold two electronsZ

XY

• So the electron configuration of hydrogen and helium are more properly written

• 1H 1s1

• 2He 1s2

• 1H 1s1

• 2He 1s2

• The 1 refers to the energy level

• 1H 1s1

• 2He 1s2

• The s refers to the sublevel

• 1H 1s1

• 2He 1s2

• The superscripts are the number of electrons in this sublevel.

2nd energy level

• -has 2s (bigger, still spherical) and 2p (has 3 bi-lobed orbitals in x, y, and z directions)

• -holds up to 8 electrons total

(2 in the s and 3 x 2 in the p)

Z

XY

Z

XY

Z

XY

• 3Li 1s22s1

• 4Be 1s22s2

• 5B 1s22s22p1

• 6C 1s22s22p2

• 7N 1s22s22p3 …

Third energy level

has three sublevels, 3s, 3p, and 3d.

The 3s and 3p sublevels are similar in structure (but bigger) than the s and p sublevels seen before.

The 3d sublevel

• (and any d) has five orbitals of varying shapes.

• These orbitals can hold two electrons each for a total of ten electrons.

• The 3d sublevel is the highest energy sublevel of energy level 3, so high, in fact, that energy level 4 begins to fill (the 4s sublevel fills) before the 3d sublevel

See text

Z

XY

Z

XY

Z

XY

Z

XY

Z

XY

See handout

The fourth energy level has four sublevels (notice the trend?) They are called 4s, 4p, 4d, and 4f. The s, p, and d sublevels are structured as before.

See handout

The f sublevel has seven orbitals, and can hold up to fourteen electrons.

The 5s sublevel is filled before the 4d, and the 5p and 6s sublevels precede the 4f.

Let’s look at a picture instead

You must remember this…

• s=1 orbital, 2 electrons

• p=3 orbitals, 6 electrons

• d=5 orbitals, 10 electrons

• f=7 orbitals, 14 electrons

The Aufbau diagram

This structure is shown below. Boxes are orbitals, each can hold two electrons

1s

2s

3s

4s

5s

6s

7s

7p

6p

5p

4p

3p

2p

3d

4d

5d

6d 5f

4f

1s

2s

3s

4s

5s

6s

7s

7p

6p

5p

4p

3p

2p

3d

4d

5d

6d 5f

4f

Rules, rules, rules.

• Orbitals are filled according to three rules:– Aufbau (building up) principle—lower energy

sublevels are filled first– Pauli exclusion principle—electrons sharing

an orbital must have opposite spins– Hund’s Rule—when a sublevel has several

orbitals, electrons will distribute to separate orbitals with parallel spins, before sharing orbitals with opposite spins

1s

2s

3s

4s

5s

6s

7s

7p

6p

5p

4p

3p

2p

3d

4d

5d

6d 5f

4f

Levels have sublevels

• All Electron configurations are some subset of the order shown below. Only the last sublevel might be incomplete

• 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10

5p6 6s2 4f14 5d10 6p6 7s2 5f14 6d10 7p6…

Watch out for two things

1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d10 6p6 7s2 5f14 6d10 7p6 … gets old.

Ex:

87Fr 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d10 6p6 7s1

87Fr [Rn] 7s1

(Radon is a noble gas, and accounts for the first 86 electrons)

Look for the last octet and use a noble gas core

• Practice:

• Write the full EC for Titanium (element 22) and write the EC with a noble gas core

• Practice:

• Write the full EC for Titanium (element 22) and write the EC with a noble gas core

• A) Ti 1s2 2s2 2p6 3s2 3p6 4s2 3d2

• and Ti [Ar] 4s2 3d2

…and watch out for Cu and Cr!• Chromium, copper and a few others

rearrange electrons to become more stable.

• Sublevel energies go down when a sublevel is full or half full.

• Cr 1s2 2s2 2p6 3s2 3p6 4s2 3d4 becomes• Cr 1s2 2s2 2p6 3s2 3p6 4s1 3d5 and

• Cu 1s2 2s2 2p6 3s2 3p6 4s2 3d9 becomes• Cu 1s2 2s2 2p6 3s2 3p6 4s1 3d10

Problems:

1. Write the electron configurations for phosphorus and molybdenum. Then draw the aufbau diagrams for these elements.

2. Write the complete electron configurations for magnesium, sulfur, and potassium. Then write their electron configurations using the symbols for the noble gases.

3. What element is represented by [Ne]3s23p6?

4. Determine the electron configuration for the last SUBLEVEL of the following elements: S, Pt, Sr, K, and Al.

5. The an unknown element has an electron configuration of 1s22s22p63s23p4.

A. What is the element?

B. What does the superscript 6 refer to?

C. What does the letter s refer to?

D. What does the coefficient 3 refer to?

6. Write the electron configuration for calcium, a nutrient essential to healthy bone growth and development.

7. Write the electron configuration for copper, which is used in pennies.

8. Use the symbols for the noble gases to write the electron configurations for the following elements:

A. Zr B. U C. Rn

9. Write the electron configuration and draw the orbital diagrams for the following elements:

A. Carbon B. Silver C. Aluminum

VOCABULARY

• Electron configuration• Atomic number• Energy level• Valence level• Sublevel• s,p,d,f• Orbital• Spin

• Proton• Electron• Spherical• Two-lobed• Dumbell-shaped• Aufbau principle• Pauli’s exclusion

principle• Hund’s rule

Be able to:

• Describe the levels, sublevels, and orbitals.• Recreate the aufbau order from the periodic

chart• Write a complete EC and EC with a noble gas

core for any element• Determine the last sublevel, and number of

electrons there from a position on a periodic chart

• Identify a position on the chart and the element from an EC

• Fill out an aufbau diagram for any element

Trends in the Periodic Chart

…since the position on the chart indicates electrons, and electrons are responsible for physical and chemical

properties…

• …the position on the periodic chart indicates the physical and chemical properties of elements!

• Opposite charges attract– electrons are attracted by the protons in the nucleus

Simple

• Valence electrons

– The representative (tall) columns represent the number of valence electrons

– Transition elements have only two valence electrons, but have a part-filled d sublevel

Simple

• Ionization pattern

• Metals

- have only a few valence electrons

-lose these electrons, empty their valence level

-form positive ions

Simple

• Ionization pattern

• Non-metals have

-more valence electrons

-gain electrons to fill valence level

-form negative ions

Medium

• Atomic size

• In a column, a lower row indicates an extra energy level

• Outer energy levels are larger

• The largest atom in a column is the bottom element

Medium

• Ionization energy —the energy required to remove an electron

X + IE X+ + e-

• Outer energy levels are farther from the nucleus• It is easier to remove an electron from a larger

energy level• The lowest ionization energy is at the bottom of the

column

Medium

• Electronegativity —the attraction an atom has for a shared pair of electrons

• See IE—the highest electronegativity is at the top

Medium

• Electron affinity —attraction an atom has for an electron from the outside

X + e- X- + EA

• See e-neg—the highest electron affinity is at the top

One moment…

As you go across a row, you get more protons in the nucleus

• They attract the electrons better

• Each energy level gets smaller

Hard

• Atoms get smaller as you go across a row

• Ionization energy gets larger

• Electronegativity gets larger

• Electron affinity gets larger

--All because there are more protons--

Recap

• Atomic size

• INCREASES as you go down

and left

Recap

• Electronegativity, ionization energy, and electron affinity

• INCREASE as you go up

and right

Pop Quiz

• Which element on the entire periodic chart is the largest?

• Smallest?

• Which element on the entire periodic chart has the largest IE?

• Smallest?

Pop Quiz

• Which element on the entire periodic chart has the largest e-negativity?

• Smallest?

• Which element on the entire periodic chart has the largest EA?

• Smallest?

Pop Quiz

• Who’s your favorite pop star?

• Which has the greatest / least: Size, IE, EA, e-neg?

C

Sn I

F

The diagonal effect

• Of the previous four elements, the ones in the same row and column are easy, right?

• What can you say about carbon and iodine?

• They might be just about the same!

The metal/nonmetal line

Diagonal effect!

(due to electronegativity or EA)

Hard (cont’d)

• Ionic radius

Negative ions are (way!) larger than their atom

Positive ions are (way!) smaller than their atom

• Which has the largest / smallest ion?

Li

K Ca

Be

• Which has the largest / smallest ion?

S

Te I

Cl

• Which has the smallest / largest ion?

Mg

Sr Te

S

• Which has the smallest / largest ion?

+2 -2

+2 -2

Mg

Sr Te

S

Hard (cont’d)

• Second and third ionizations

• If you ionize an atom, you make a (+) ion

• It’s harder to ionize it again

• It gets way harder after you empty the valence level

Hard (cont’d)

• First, second, and third ionization energies

X + IEX+1 + e-

X+1 + IE2X+2 + e-

X+2 + IE3X+3 + e-

Shielding

• Shielding –weakening of attraction due to electrons interfering with attraction of the nucleus

• Shielding increases a little as you go across a period (not as much as attraction)

• Shielding jumps tremendously as you start a new energy level

Remem-mem-mem…

• Ionic radius



Remem-mem-mem…

• Ionic radius


…due to shielding by the extra electrons


…because they have lost shielding or shielded electrons

Disclaimer

• Noble gasses have no electronegativity—they don’t share electrons

• Noble gasses have no electron affinity—they don’t gain electrons

• Most metals have no electron affinity—they don’t gain electrons

• Ionization Energies in kJ/mol• 1 2 3 4 5H 1312He 23725250Li 520 7297 11810Be 899 1757 14845 21000B 800 2426 3659 25020 32820C 1086 2352 4619 6221 37820N 1402 2855 4576 7473 9442O 1314 3388 5296 7467 10987F 1680 3375 6045 8408 11020Ne 2080 3963 6130 9361 12180Na 496 4563 6913 9541 13350Mg 737 1450 7731 10545 13627

Size of metal atoms

Size of Atoms (radius, nm)

Size of Anions

A positive charge is

worth about three energy

levels!

Size of Cations

Electron Affinity*, Electronegativity*, Ionization energy

Atomic number, shielding, diagonal effect

Th-th-that’s all, folks.

…and, in conclusion…

Documents

frequency of light

energy q

excited electron q

speed of light

higher energy level

velocity of light c

lower energy level

plancks equationthe