electron configuration! of orbitals as sort of a "border” for ... has an s orbital. ......

Electron Configuration!Chapter 5

DO NOW - Finish coloring your periodic tables! (5 min)

State at Room Temperature

Appearance Conductivity Malleability and Ductility

Metals - solid except for mercury (a liquid)

- shiny lustre - good conductors of heat and electricity

- malleable- ductile

Non-Metals - some gases- some solids- only bromine is a liquid

- not very shiny - poor conductors of heat and electricity

- brittle- not ductile

Metalloids - solids - can be shiny OR dull

- may conduct electricity- poor conductors of heat

-brittle - not ductile

Helpful Videos

https://www.youtube.com/watch?v=i77ISmwTJ8M

https://www.youtube.com/watch?v=J-DjEIlynjE

https://www.youtube.com/watch?v=Aoi4j8es4gQ

https://www.youtube.com/results?search_query=firework+colors+chemistry

Quantum Mechanics

● Better than any previous model, quantum mechanics does explain how the atom behaves.

● Quantum mechanics treats electrons as particles that act like waves (like light waves) which can gain or lose energy.

● But they can’t gain or lose just any amount of energy. They gain or lose a “quantum” of energy.

A quantum is just an amount of energy that the electron needs to gain (or lose) to move to the next energy level.In this case it is losing the energy and dropping a level.

Atomic Orbitals

● Much like the Bohr model, the energy levels in quantum mechanics describe locations where you are likely to find an electron.

● Remember that orbitals are “geometric shapes” around the nucleus where electrons are found.

● Quantum mechanics calculates the probabilities where you are “likely” to find electrons.

Atomic Orbitals

● Of course, you could find an electron anywhere if you looked hard enough.

● So scientists agreed to limit these calculations to locations where there was at least a 90% chance of finding an electron.

Think of orbitals  as sort of a "border” for spaces around the nucleus inside which electrons are allowed. No more than 2 electrons can ever be in 1 orbital. The

orbital just defines an “area” where you can find an electron.

Energy Levels

Quantum mechanics has a principal quantum number. It is represented by a little n.

It represents the “energy level” ● n=1 describes the first energy level● n=2 describes the second energy level… etc.

Each energy level represents a period (row) on the periodic table. Red n = 1

Orange n = 2Yellow n = 3Green n = 4Blue n = 5Indigo n = 6Violet n = 7

Sub-levels = Specific Atomic Orbitals

Each energy level has 1 or more “sub-levels” which describe the specific “atomic orbitals” for that level.

● n = 1 has 1 sub-level (the s orbital)● n = 2 has 2 sub-levels (s & p)● n = 3 has 3 sub-levels (s, p, & d)● n = 4 has 4 sub-levels (s, p, d & f)

There are 4 types of atomic orbitals:● s, p, d and f● Each of these sub-levels represent the

blocks on the periodic table.

Blue = s blockYellow = p blockRed = d blockGreen = f block

Orbitals

● In the s block, electrons are going into s orbitals.

● In the p block, the s orbitals are full. New electrons are going into the p orbitals.

● In the d block, the s and p orbitals are full. New electrons are going into the d orbitals.

s p

d

● Complete the chart in your notes as we discuss this.

● The first level (n=1) has an s orbital. It has only 1. There are no other orbitals in the first energy level.

● We call this orbital the 1s orbital.

Energy Level

Sub-levels Total Orbitals Total Electrons

Total Electrons per Level

n = 1 s 1 (1s orbital) 2 2

n = 2 sp

1 (2s orbital)3 (2p orbitals)

26

8

n = 3 spd

1 (3s orbital)3 (3p orbitals)5 (3d orbitals)

2610

18

n = 4 spdf

1 (4s orbital)3 (4p orbitals)5 (4d orbitals)7 (4f orbitals)

261014

32

Where are these Orbitals?

1s

2s

3s

4s

5s

6s

7s

3d

7p

6p

5p

4p

3p

2p

5f

4f

6d

5d

4d

Electron Configurations

What do I mean by “electron configuration?”

●The electron configuration is the specific way in which the atomic orbitals are filled.

●Think of it as being similar to your address. The electron configuration tells me where all the electrons “live.”

Rules for Electron Configurations

3 rules govern electron configurations:● Aufbau Principle● Pauli Exclusion Principle● Hund’s Rule

Using the orbital filling diagram at the right will help you figure out HOW to write them

● Start with the 1s orbital. Fill each orbital completely and then go to the next one, until all of the elements have been accounted for.

Fill Lower Energy Orbitals FIRST

● The Aufbau Principle states that electrons enter the lowest energy orbitals first.

● The lower the principal quantum number (n) the lower the energy.

● Within an energy level, s orbitals are the lowest energy, followed by p, d and then f. F orbitals are the highest energy for that level.

Each line represents an orbital.

1 (s), 3 (p), 5 (d), 7 (f)

Low Energy

High Energy

Write the electron config. below for each of these elements:

Sodium:

Iron:

Bromine:

Barium:

Write the electron config. below for each of these elements:

Sodium: 1s22s22p63s1

Iron: 1s22s22p63s23p64s23d6

Bromine: 1s22s22p63s23p64s23d104p5

Barium: 1s22s22p63s23p64s23d104p65s24d105p66s2

No more than 2 Electrons in Any Orbital…ever

● The next rule is the Pauli Exclusion Principal.

● The Pauli Exclusion Principle states that an atomic orbital may have up to 2 electrons and then it is full.

● The spins have to be paired.

● We usually represent this with an up arrow and a down arrow.

● Since there is only 1 s orbital per energy level, only 2 electrons fill that orbital.

Quantum numbers describe an electrons position, and no 2 electrons can have the exact same quantum numbers. Because of that, electrons must have opposite spins from each other in order to “share” the same orbital.

Hund’s Rule

● Hund’s Rule states that when you get to degenerate orbitals, you fill them all half way first, and then you start pairing up the electrons.

● What are degenerate orbitals?● Degenerate means they have the

same energy.

Don’t pair up the 2p electrons until all 3 orbitals are half full.

DO NOW

1. Please take your packet

out for me to check

2. Go over packet

3. Start lab activity

Noble Gas Notation

Let’s try it out!

● NOW that we know the rules, we can try to write some electron configurations.

● Lets write some electron configurations for the first few elements, and let’s start with hydrogen.

Electron Configurations

Element Configuration Element ConfigurationH Z=1 1s1 He Z=2 1s2

Li Z=3 1s22s1 Be Z=4 1s22s2

B Z=5 1s22s22p1 C Z=6 1s22s22p2

N Z=7 1s22s22p3 O Z=8 1s22s22p4

F Z=9 1s22s22p5 Ne Z=10 1s22s22p6

(2p is now full)Na Z=11 1s22s22p63s1 Cl Z=17 1s22s22p63s23p5

K Z=19 1s22s22p63s23p64s1 Sc Z=21 1s22s22p63s23p64s23d1

Fe Z=26 1s22s22p63s23p64s23d6 Br Z=35 1s22s22p63s23p64s23d104p5

Note that all the numbers in the electron configuration add up to the atomic number for that element. Ex: for Ne (Z=10), 2+2+6 = 10

● One last thing. Look at the previous slide and look at just hydrogen, lithium, sodium and potassium.

● Notice their electron configurations. Do you see any similarities?

● Since H and Li and Na and K are all in Group 1A, they all have a similar ending. (s1)

Electron Configurations

Element ConfigurationH Z=1 1s1

Li Z=3 1s22s1

Na Z=11 1s22s22p63s1

K Z=19 1s22s22p63s23p64s1

●This similar configuration causes them to behave the same chemically.●It’s for that reason they are in the same family or group on the periodic table.●Each group will have the same ending configuration, in this case something that ends in s1.

Quantized Energy

Electromagnetic Radiation

● Visible light is a type of electromagnetic radiation - A form of energy that exhibits wavelike behavior through space.

● Electrons are particles that act like waves.○ Electrons that transition from excited to ground

state has a different amount of energy and a different wavelength of light.

Electromagnetic Wave Relationship

Amplitude is the wave’s height from the origin to a crest

Electromagnetic Spectrum

The electromagnetic spectrum (EM spectrum) includes all forms of electromagnetic radiation, with

the only difference in types of radiation being frequencies and wavelengths.

Electromagnetic Wave Relationship

Frequency and wavelength are inversely proportional

c: speed of light (3.00 x 108 m/s) v: frequency (Hz)

: wavelength (m, nm, etc.)

c = x v

Lets try a problem...

Calculate the wavelength of a wave that has the frequency of 3.44 x 109 Hz.

Calculate the wavelength of a wave that has the frequency of 3.44 x 109 Hz.

c = x v → = c/v

= 3.00 x 108 m/s

3.44 x 109 Hz

Hz = S -1

Calculate the wavelength of a wave that has the frequency of 3.44 x 109 Hz.

= 3.00 x 108 m/s

3.44 x 109 s -1

= 8.72 x 10-2 m

Energy of Quantum

A quantum is the minimum amount of energy that can be gained or lost by an atom. Max Planck studied light by heating objects. These

different colors correspond to different frequencies and wavelengths.

E: Energy (Joules or kJ) v: frequency (Hz or s-1)h: Planck’s constant 6.626 x 10-34 J x S

Equantum = h x v

What is a photon?

A massless particle that carries a quantum of energy.

They can be absorbed or released by electrons!

Photoelectric Effect

When a photon hits an electron on a metal surface, the electron can be emitted. The emitted electrons are

called photoelectrons.

Lets try a problem...

What is the energy of a photon from the violet portion of the Sun’s light if it has a frequency of 7.230 x 1014 s-1?

Lets try a problem...

What is the energy of a photon from the violet portion of the Sun’s light if it has a frequency of 7.230 x 1014 s-1?

Equantum = h x v

*Unknown is E

E = (6.626 x 10-34 J.s) (7.230 x 1014 s-1)

E = 4.791 x 10-19 J

Atomic Emission Spectrum

Flame Test

Atomic Emission Spectra

A set of frequencies of the electromagnetic waves emitted by atoms of the element. Each element is unique!

Periodic Trends

Periodic Trends

● More than 20 properties change in predictable way based location of elements on PT

● Some properties: ○ Density○ Melting point/boiling point○ Atomic radius○ Ionization energy○ Electronegativity

Going down group 1 →

2-8-18-32-18-8-1Fr72-8-18-18-8-1Cs62-8-18-8-1Rb52-8-8-1K42-8-1Na32-1Li21H1

ConfigurationElementPeriod

Increasing # of energy levels ...

Increasing number of energy levels

Atomic Radius

● How large the radius is of an atom is… ○ what do you think influences that?

Incr

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nerg

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Increasing Atomic Radius

next

Li: Group 1 Period 2 Cs: Group 1 Period 6

Cs has more energy levels, so it’s bigger

2-8NeVIIIA or 182-7FVIIA or 172-6OVIA or 162-5NVA or 152-4CIVA or 142-3BIIIA or 132-2BeIIA or 22-1LiIA or 1

ConfigurationElementFamily

As we go across, elements gain electrons, but they are getting smaller!

Incr

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Incr

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Decreasing Atomic Radius

Why does this happen?

• As you go from left to right, you again more protons (the atomic number increases)

• You have greater “proton pulling power” – Remember the nucleus is + and the electrons

are - so they get pulled towards the nucleus• The more protons your have, the more

Proton Pulling Power

Ionization Energy

• Amount energy required to remove a valence electron from an atom in gas phase

• 1st ionization energy = energy required to

remove the most loosely held valence electron (e- farthest from nucleus)

•Cs valence electron lot farther away from nucleus than Li •electrostatic attraction much weaker so easier to steal electron away

from Cs•THEREFORE, Li has a higher Ionization energy than Cs

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Incr

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Incr

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Decreasing Atomic Radius

Incr

ease

d El

ectro

n Sh

ield

ing

Decreased Ionization Energy (easier to remove an electron)

Increased Ionization Energy (harder to remove an electron)

Electronegativity

• ability of atom to attract electrons in bond

• noble gases tend not to form bonds, so don’t have electronegativity values

• Fluorine: most electronegative element = 4.0 Paulings

Incr

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umbe

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nerg

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Incr

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Decreasing Atomic Radius

Incr

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Dec

reas

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niza

tion

Ener

gy (e

asie

r to

rem

ove

an e

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ron)

Increased Ionization Energy (harder to remove an electron)

Decreased Electronegativity

Increased Electronegativity

Electronegativity

• ability of atom to attract electrons in bond

• noble gases tend not to form bonds, so don’t have electronegativity values

• Fluorine: most electronegative element = 4.0 Paulings