electron arrangement in atoms chapter 4. development of atomic model(s)

Electron Arrangement in Atoms

Chapter 4

Development of Atomic Model(s)

• Light is thought to be a wave– Since the 1900’s

• Newton thought light consists of particles.– Interestingly, the current view of atoms, the

quantum mechanical model of the atom, came about from the study of light.

Electromagnetic Spectrum: all the forms of light - acting like waves

Particle: would cause 2 linesWaves: interference causes multiple lines

Light Waves

• Wavelength ()– distance between the crests. (unit =nm)

• Amplitude– wave’s height from zero to the crest.

• Frequency ()– Is the number of waves to pass a given point per unit of

time.– The unit of frequency is cycles/second which is called the

hertz (Hz)

• C= speed of light, a constant

• c = wavelegth x frequency or c = • The product of frequency and wavelength

always equals a CONSTANT (c), the speed of light.

As wavelength increases, frequency decreases

• ALL EMS waves travel at c.

Which wave has greatest wavelength Greatest amplitude? How do the frequencies compare?

A

B

C

D

E

Electromagnetic Spectrum• As wavelength decreases, frequency increases

– High frequency waves = gamma– Low frequency waves = radio

• As frequency increases, energy increasesQuantum: min amount of energy that can be lost by an electron

Visible light is a small part of the ES• Parts of visible light: ROYGBIV• When visible light passes through prism, it can be

separated

Infrared waves: Low frequency/high wavelengthUltraviolet waves: High frequency/low wavelength

So…infrared is LOWER in energy than UV

atomic emission spectrum: the discrete lines created when an element emits distinct frequencies of light- unique to each element (like a fingerprint) - light is emitted when electrons move energy levels

from an exited state to a ground state • When atoms absorb energy, e- move into higher energy levels…

but they don’t stay there…these e- then lose energy (a quanta) and emit light when they return to their lower energy levels.

Photon: quantum “particle” of energyQuantum: min amount of energy that can be lost by an electron

Check point• How does the electromagnetic spectrum differ

from the atomic emission spectrum?

– Shows spectrum of all light vs distinct light given off my certain atoms?

The Dual Nature of Light• Is light a PARTICLE or a WAVE?• Quantum of light are called photons and

behave like PARTICLES but have WAVE properties– Light has BOTH qualities (Thx Einstein)– Light is a particle that travels like a wave!– Not all scientists believed

this at first…now it is accepted as FACT

De Broglie won the Nobel Prize for his studies on the wave nature of matter, changing the way the world saw quantum physics.

Classical Physics: Deals with average size objects traveling at average speeds…

Quantum Mechanics/ Physics: deals with atomic size objects traveling at incredibly high frequencies (speeds)

• Watch this physicist trying to explain what De Broglie opened up with this theory– https://www.youtube.com/watch?v=JIGI-eXK0tg#t=15 start

at 1:12

Bohr Model• Each possible e- has a fixed energy called an

energy level …like steps on a ladder.– Switching levels requires a gain or loss of energy in

the form of photons.• Gain energy to go up a level (farther from nucleus)• Loss energy to go down a level

– High energy levels are furthest from the nucleus.

Drawing an atom with Bohr model:• Can abbreviate what’s in the nucleus• First orbital holds only 2 e-

• all other orbitals hold 8 e-

Draw model for boron:

how electrons can be excited and move to excited energy states which is higher than their original ground state

The Heisenberg Uncertainty Principle• it is impossible to know exactly both the velocity and the position of a particle at the same time.

• If you determine one, it affect and changes the other, so you can’t ever calculate both at the same time. Ex: measuring position changes its velocity making its location uncertain.

• Based on Schrödinger equation: the same Schrödinger who came up with the dead and alive cat paradox to explain the Copenhagen interpretation of quantum mechanics

https://www.youtube.com/watch?v=QisnPsu7_Uk

• Erwin Schrödinger (1926) used theoretical mathematical calculations to describe e- motion.

• This model used mathematical equations describing the behavior of the e- in a hydrogen atoms.

• Modern quantum mechanics deals with the laws of motion which govern the behavior of atomic and subatomic particles.

Quantum Mechanical Model• Location of e- shown by a cloud.• mathematical theory that predicts the probability of

electron in certain area

Check point

• How does the quantum mechanical model differ from the Bohr model of an atom?

Quantum Mechanical Model cont. Includes 4 quantum #s describing the position of electrons • The quantum mechanical model of the atom does NOT

involve an exact path the e- takes around the nucleus (like the Bohr model does).

• Rather, it is based on the likelihood or probability of where the e- will be in various locations around the nucleus.

Valence e-: electrons in outermost orbital• have the most

potential energy

Quantum Numbers SummaryPrinciple: n

Energy level

1-7 periods

Angular: l

Orbital type

s, p, d, f,

Magnetic: m

Orbital orientation

--------------

Spin Electron spin direction

+1/2 or -1/2

4 quantum numbers:1. Principle Quantum #: n

- Describes energy level occupied by e-

- represented by the periods 1-7 (horizontal rows)

the larger the number, the farther the valence e- from the nucleus, the more energy (potential)

Valence e-: electrons in outermost orbital• have the most potential energy

4 quantum numbers:2. Angular quantum #: l

- Indicates shape of orbital: s, p, d, f

4 quantum numbers:3. Magnetic Quantum #: m

- orientation of orbital around nucleus

4 quantum numbers:4. Spin Quantum #: +1/2 or -1/2

- Spin of an electron

Quantum Numbers SummaryPrinciple: n

Energy level

1-7 periods

Angular: l

Orbital type

s, p, d, f,

Magnetic: m

Orbital orientation

--------------

Spin Electron spin direction

+1/2 or -1/2

https://www.youtube.com/watch?v=accyCUzasa0

Electron configurations• Electron configurations are similar to postal

“zipcodes”.– Written distribution of electrons in their orbitals

• ExamplesHydrogen has 1 electron: 1s1

He has 2 electrons: 1s2

Li has 3 electrons: 1s2 2s1

Be has 4 electrons: 1s2 2s2

B has 5 electrons: 1s22s22p1

Sublevels: s, p, d, f • Within each energy level, there are sublevels

– s sublevel- consist of e- from groups 1 and 2– p sublevel- consist of e- from groups 13-18– d sublevel- consists of e- from groups 3-12– f sublevel- consists of e- from the inner transition

metals S

pd

f

Orbitals: hold a maximum of 2 electrons.• Each sublevel can be broken down into

orbitals:– s sublevel: 1 orbital : 2 e- total– p sublevel: 3 orbitals: 6 e- total– d sublevel: 5 orbitals: 10 e- total– f sublevel: 7 orbitals: 14 e- total

Rules for writing electron configurations1.The number of electrons in an atom is equal to

the atomic number; unless told otherwise. 2.Electrons fill the lowest energy level/sublevel

before moving to a higher energy level/sublevel. Order: s, p, d, f

3.All orbital must have one electron before any getting a second electron

4.An orbital cannot take more than 2 electrons, and they must have opposite spins.

P: 1s22s22p63s23p3

Orbital Diagram

Electron Configuration

Ex P: 1s22s22p63s23p3

RIGHTWRONG

General Rules• Hund’s Rule = “empty bus seat rule”

– Within a sublevel, place one electron per orbital before pairing them.

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S: 1 orbital : 2 e- totalp: 3 orbitals: 6 e- totald: 5 orbitals: 10 e- totalf: 7 orbitals: 14 e- total

O 8e-

• Orbital Diagram

• Electron Configuration

1s2 2s2 2p4

Notation

1s 2s 2p

Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

O15.9994

8

S: 1 orbital : 2 e- totalp: 3 orbitals: 6 e- totald: 5 orbitals: 10 e- totalf: 7 orbitals: 14 e- total

Filling Rules for Electron Orbitals come from these principles:

Aufbau Principle: Electrons are added one at a time to the lowest energy orbitals available until all the electrons of the atom have been accounted for.

Pauli Exclusion Principle: An orbital can hold a maximum of two electrons.To occupy the same orbital, two electrons must spin in opposite directions.

Hund’s Rule: Electrons occupy equal-energy orbitals so that a maximum number of unpaired electrons results.

*Aufbau is German for “building up”

Spin Quantum Number (ms): +1/2, -1/2

Practice:

1.Be

2.N

3.Ni

4.He

5.S

1.Be2.N3.Ni

4.He5.S

Practice: S

p

d

f

3d456

4f5f

Practice:

1.Be

2.N

3.Ni

4.He

5.S

1.1s2 2s2

2.1s2 2s2 2p3

3.1s2 2s2 2p6 3s2 3p6 4s2 3d8

(or 1s2 2s2 2p6 3s2 3p6 3d8

4s2)

4.1s2

5.1s2 2s2 2p6 3s2 3p4

Orbital Filling

Element 1s 2s 2px 2py 2pz 3s Configuration

Electron ConfigurationsElectron

H

He

Li

C

N

O

F

Ne

Na

1s1

1s22s22p63s1

1s22s22p6

1s22s22p5

1s22s22p4

1s22s22p3

1s22s22p2

1s22s1

1s2

• Shorthand Configuration

S 16e-

Valence ElectronsCore Electrons

S 16e- [Ne] 3s2 3p4

1s2 2s2 2p6 3s2 3p4

Notation

• Longhand Configuration

Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

S32.066

16

Abbreviate configuration using noble gas notation: Ex: 1. Zn 2. Sn

1.Be2.N3.Ni

4.He5.S

Practice: S

p

d

f

3d456

4f5f

Practice:

1.Be

2.N

3.Ni

4.He

5.S

1.1s2 2s2

2.1s2 2s2 2p3 = [He] 2s2 2p3

3.1s2 2s2 2p6 3s2 3p6 4s2 3d8

[Ar] 4s2 3d8

4.1s2

5.1s2 2s2 2p6 3s2 3p4

abbreviated notation

Bozeman orbitals: start at 6:49 https://www.youtube.com/watch?v=2AFPfg0Como

This fills the valenceshell and tends to givethe atom the stabilityof the inert gasses.

The Octet RuleAtoms tend to gain, lose, or share electrons until they have eight valence electrons.

8

ONLY s- and p-orbitals are valence electrons.

• The groupA number equals the number of electrons in its outermost energy level

(this will be more important later on…)

• So all Group 1A elements have 1 electron in the outer shell

Stability: to gain stability, atoms will gain or lose electrons, forming an ion with the same # of e- as the closest noble gas

Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

• Electron Configuration Exceptions to gain stability:

– Copper

EXPECT: [Ar] 4s2 3d9

ACTUALLY: [Ar] 4s1 3d10

– Copper gains stability with a full d-sublevel.

1.Be2.N3.Ni

4.He5.S

Practice: S

p

d

f

3d456

4f5f

1

2

3

4 5

6

7

Stability• Ion Formation

– Atoms gain or lose electrons to become more stable to be isoelectronic with the Noble Gases.

Group 1-2 lose elections = cations Group 3A to 6A gain electrons= for anions (except He!)

Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

form cations! form anions!

METALS form CATIONS

• Heavy metal cats…ROCK ON!!

Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

• Ion Electron Configuration– Write the e- configuration for the closest Noble

Gas• EX: Oxygen ion O2- Ne

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

Gilbert N. Lewis• Lewis suggested that atoms with

fewer than 8 valence e- bond together to share e- and complete their valence shells.

• He noticed that many elements are more stable when they have eight electrons in their outer shell.

Lewis Dot Structure• Simple dot diagram of valence electrons

Drawing Lewis Structures

• Practice writing Lewis Dot Structures for the following elements:

1. K

2. Mg

3. Ga

4. C5. P6. Se7. Br8. Ar

1. K2. Mg3. Ga

Practice: S

p

d

f

4.C5.P

6.Se7.Br

8.Ar

Ch 5: Periodic Trends• Valence electron rules, electron configuration,

and cation and anion formation all are part of the trends that govern the periodic table!

Electrons are the Key!

Valence electrons: group elements have same number of valence electrons

• The reason elements in the same family(group) have the same chemical and physical properties is because they have the same number of electrons in their outer shell

Ch 5

Ch 5 Summary of Trends

Increasing Electronegativty

Increasing Ionization Energy

Decreasing Atomic RadiusD

ecre

asin

g I

on

izat

ion

En

erg

y

Dec

reas

ing

Ele

ctro

neg

ativ

ity

Incr

easi

ng

Ato

mic

Rad

ius

Atomic Radius (Size)• decrease as you go L to

R across period.– Same energy level but:– Add p+ and e-, increase

nuclear charge, pulls in orbitals closer to the nucleus

• increase as you go down a group– Electrons being added to

outer orbital (increasing principle energy level)

Ion Size

• Cations are smaller in size than the neutral element.

• Anions are larger in size than the neutral element.

Ionization Energy

• Ionization Energy (IE): The energy required to remove and electron from a gaseous atom.– Remove 1st electron = 1st IE– Remove 2nd electron = 2nd IE

Ionization Energy cont.• IE generally increases as you move L to

R across the period.– Harder to remove an electron as you go L to

R because of greater attraction to nucleus =Shielding effect

• IE generally decreases as you go down a group. (first IE only)– Atom gets bigger, outermost e- farthest from

nucleus, easy to be removed.

Electronegativity• Electronegativity: The tendency for

atoms to attract electrons when they are chemically combined. – Do they share electrons equally? (remember

polar water bonds1?)

Electronegativity Trends• increases as you go L to R across the period

– Elements want to be like noble gases!

• decreases as you go down a group.

The most electronegative element is Fluorine.

– Noble gases have no electronegativity as they already have a full valence shell.

Summary of Trends

Increasing Electronegativty

Increasing Ionization Energy

Decreasing Atomic RadiusD

ecre

asin

g I

on

izat

ion

En

erg

y

Dec

reas

ing

Ele

ctro

neg

ativ

ity

Incr

easi

ng

Ato

mic

Rad

ius

Is that it!