Electrons in Atoms S P d f YouTube - The Atom Song.

Electrons in Atoms

S

P

d

f

YouTube - The Atom Song

Early models of the Atom

J.J.Thomson’s Plum Pudding Model

Dalton’s Model

Rutherford’s model of the Atom

If all atoms looked the same then,,,,,

they could not explain why different elements had different properties

Problems with Rutherford:Rutherford theory does not explain what the electrons are doing

Does not explain the experimental phenomenon of atomic spectra

Atomic Spectraatomic spectra- The range of

characteristic frequencies of electromagnetic radiation that are readily absorbed and emitted by an atom.

Neon lights…..

http://www.upscale.utoronto.ca/PVB/Harrison/Spectra/Spectra.html

Rydberg

Could mathematically explain atomic spectra – but no theory to apply it to….

Rydberg

Rydberg could mathematically explain the atomic spectra,(by treating particles as waves WHAT?- we will get to that later) but it wasn’t attached to any theory or model

Rydberg Equation

Rydberg developed this equation to explain Hydrogen’s atomic spectra

RH=109677.57 cm-1

This behavior could only be explained if the atoms gave off light in specific amounts of energy………we needed a better model of the atom to explain this…….

Idea!

Bohr Model of Atom1913

Protons and Neutrons in Nucleus Electrons were arranged in Energy Levels at

fixed positions around the Nucleus He thought- different elements had

different electron arrangements and therefore different properties

Atomic Structure Review Nucleus

Protons ( made of 3 quarks) Neutrons (made of 3 quarks)

Electron Cloud Area where electrons are located Electrons are arranged in energy levels

They found out that electrons don’t stay in one energy level, they can move from one energy level to another depending on how much energy they have

Ground state – where the electron usually is.

Excited state – when the electron has absorbed a quantum of energy is in a higher energy level

Amount of energy needed to move an electron from one energy level to another is called a QUANTUM of ENERGY

Electrons make QUANTUM LEAPS or SHIFTS

energy levels are not equally spaced.

different amounts of energy are needed to make the shift

The higher the energy level, the less energy it takes to make the shift to a higher level.

Once they shift up….Electrons do not hold energyEnergy quickly released (same

amount as absorbed) as packet of energy called a photon

Electron falls back down to Ground State

Einstein and Planck

• these transitions(shifts) usually involve the absorption or emission of a quantum of light.

• They named the packets of light energy called photons.

Golf ball demonstration

The frequency of light emitted by an element separates into discrete lines to give an atomic emission spectra

The emission spectra of an element is like a fingerprint-used to identify elements in stars

The Atomic Emission Spectrum: an explanation

http://chemistry.bd.psu.edu/jircitano/periodic4.html

Niels Bohr, was born 7 October 1885 in Copenhagen, Denmark and died in 1962, was the Danish physicist whose investigations of atomic structure earned him the 1922 Nobel Prize for physics. Bohr's work helped solve the problems classical physics could not explain about the nuclear model of the atom. He postulated that electrons moved in fixed orbits around the atom's nucleus, and he explained how they emitted or absorbed energy. Bohr attended the University of Copenhagen (1903-11), then studied for a time under ernest Rutherford in Manchester, england. By 1916 he was back at the University of Copenhagen as a professor of physics, and in 1920 he became the first director of the Institute of Theoretical Physics. Bohr's Institute became a gathering place for the world's top physicists, and he is considered one of the foremost scientists of modern physics, along with Albert Einstein, erwin Schrödinger and enrico Fermi. During World War II Bohr avoided Adolf Hitler's army and left Denmark in 1943; he ended up in the United States and was sent to Los Alamos, New Mexico to join Robert Oppenheimer and others working to develop the atomic bomb. After the war he returned to Denmark and spent the rest of his career advocating the peaceful uses of atomic energy.

What we learn from the Bohr atom

Electrons Travel in Fixed energy levels-energy levels are not equal-ladder- can only go from rung to rung, can’t stand in between…..energy is QUANTIZED (separate into levels)

electrons usually in the ground state- lowest energy state

Remember the mint? Nitrogen atoms in the air absorb the energy emitting UV light. Wintergreen molecules absorb the energy and release as bluish light

When an atom absorbs energy, electrons move into higher energy levels when they return to lower energy levels they lose energy and emit light in the form of photons at specific frequencies

But what is a photon? To be continued….

Electrons continue to “shift up” and “shift down”, giving off a steady flow of photons

These photons form a light wave.

n=1

n=2

n=3

n=4

Spectrum

UV

IR

Vi s ible

Ground State

Excited State

Excited StateExcited State unstable and drops back down

•Energy released as a photon

•Frequency proportional to energy drop

Excited State

But only as far as n = 2 this time

SummaryElectron normally in Ground State

Energy supplied [ as heat or

electricity]

Electron jumps to higher energy level

Now in Excited State

Unstable

Drops back to a lower level

Electromagnetic radiation

********Light acts as ********Light acts as both particles both particles and wavesand waves

Parts of a Wave

Use the Greek letter “nu” Use the Greek letter “nu” ( ( ) ) = wavelength= wavelength

Amplitude measures intensity amount of energy in the waveIn Sound = loudness

In light = brightness

Frequency

Number of waves per second ( )

Determines type of light produced measured in Hertz (Hz) (cycles per sec)

Electromagnetic Spectrum – a chart of all the light waves arranged by frequency

Electromagnetic SpectrumElectromagnetic Spectrum

Wavelength and frequency are inversely proportional

Long wavelength Long wavelength small frequency small frequencylow low energyenergyShort wavelength Short wavelength high frequencyhigh frequencyhigh high energyenergy

Frequency, Wavelength and Amplitude

Same frequency- different amplitude

Same amplitude- different frequency

If the light wave given off has the frequency visible light, color is seen

Photons have different amounts of energy

different energies= different frequencies= different wavelengths=

different colors of light

• = cc

c c = velocity of light = 3.00 x 10= velocity of light = 3.00 x 1088 m/sec m/sec

Wavelength x frequency = velocity of light

Velocity (speed) of light

What is the wavelength ( of a light wave that has a frequency of 3 x 1018 meters?

What type of electromagnetic wave is it?(see pg.139) *c c = velocity of light = 3.00 x 10= velocity of light = 3.00 x 1088 m/sec m/sec

• = cc

Ans. 1 x 10-10m Ans. X-rays

What is the frequency ( ) of a wave that has a wavelength of 1 x 10-4 m?What type of wave is it? • = cc

Ans. 3 x 10Ans. 3 x 1012 12 m/secm/sec Ans. microwaveAns. microwave

Electrons and LightCurrent model of atom comes from study of light

Physicists would heat up different elements until they glowed, and then direct the light through a prism...

White light passing through a prism

Continuous spectrum of white light

White Light made of all the various wavelengths and

frequencies of visible light ROYGBIV Red – Low energy Violet-High energy Called a continuous spectrum

Excited Gases Excited Gases & Atomic & Atomic StructureStructure

When physicists heated gases and observed the light they produced…….

didn't see the whole rainbow. only bright lines of certain colors

?

Line Emission Spectra Line Emission Spectra of Excited Atomsof Excited AtomsLine Emission Spectra Line Emission Spectra of Excited Atomsof Excited Atoms

Excited atoms emit light of only certain wavelengths

The wavelengths of emitted light depend on the arrangement of the electrons in the atom.

Spectrum is not continuousA unique line emission spectrum is formed

Spectrum of Spectrum of Excited Hydrogen GasExcited Hydrogen Gas

Spectral Lines Each element has its own unique number of electrons

and produces own set of spectral lines Since electrons absorb a different quanta of energy

as they jump up energy levels, they release photons of varying energies, producing different frequencies of light

They are the “fingerprints” of an element.

Line Spectra of Other ElementsLine Spectra of Other Elements

Let’s look at the spectral lines produced by various gases…..

http://jersey.uoregon.edu/vlab/elements/Elements.html

Atomic Line Emission Atomic Line Emission Spectra and Niels BohrSpectra and Niels BohrAtomic Line Emission Atomic Line Emission Spectra and Niels BohrSpectra and Niels Bohr

Bohr’s model of the Bohr’s model of the atom was based on atom was based on an understanding of an understanding of thethe LINE EMISSION LINE EMISSION SPECTRASPECTRA of excited of excited atoms.atoms.

Niels BohrNiels Bohr

(1885-1962)(1885-1962)

HowStuffWorks Videos "100 Greatest Discoveries: Signature Light of Elements"

Models of the Atom From Dalton to Schroedinger

Quantum Mechanical Model of the Atom

Probability of finding an electron in a certain position is based on its allowed energies

Energy levels divided into sublevels

Sublevels divided into orbitals

Electrons

electrons usually in the ground state- lowest energy state

But where are they ?

Electron ConfigurationElectron configuration gives us directions to where an electron is most likely located

Schroedinger Equation A mathematical formula in which there

is more than one solution Leads to a probability of finding an

electron in a particular place – firefly These solutions are known as quantum

numbers – which give probable locations of the electron

Arrangement of Arrangement of Electrons in AtomsElectrons in AtomsArrangement of Arrangement of Electrons in AtomsElectrons in Atoms

Electrons in atoms are arranged asElectrons in atoms are arranged as

LEVELSLEVELS (n) (n)

SUBLEVELSSUBLEVELS (l) (l)

ORBITALSORBITALS (m (mll))

ENERGY LEVELS

MAIN LOCATION OF ELECTRONS AROUND ATOMS

THE HIGHER THE ENERGY, THE GREATER THE DISTANCE FROM THE NUCLEUS

ENERGY LEVELS

Also called energy states or energy shells.

They are referred to by the principal quantum numbers 1, 2, 3, 4, 5, 6 and 7

ENERGY SUBLEVELS

An energy level is made up of one or more sublevels (subshells)

An energy level can have s, p, d and f sublevels.

Sublevels are shown as “blocks” on the Periodic Table

Current maximum electron numbers for each energy level

• Energy level 1 =(s) 2 electrons• Energy level 2 = (s,p) 8 electrons• Energy level 3 = (s,p,d) 18 electrons• Energy level 4 = (s,p,d,f) 32 electrons• Energy level 5 = (s,p,d,f) 32 electrons• Energy level 6 = (s,p,d) 18 electrons• Energy level 7 =(s) 2 electrons

ELECTRON ORBITALS

Each sublevel is made up of one or more electron orbitals

Orbitals are 3-Dimensional areasThe orbitals in sublevels are referred

to as s orbitals,p orbitals, d orbitals and f orbitals

Each orbital can only hold 2 electronsWithin an orbital the electrons spin

in opposite directions so they don’t repel each other

Each orbital (area) has a specific shape. (remember they are 3- dimensional)

Orbitals and Orbitals and their Shapestheir Shapes

Types of Orbitals Types of Orbitals

s orbitals orbital p orbitalp orbital d orbitald orbital

S orbital

p Orbitalsp Orbitalsp Orbitalsp Orbitals

The three p orbitals lie 90The three p orbitals lie 90oo apart in spaceapart in space

P orbital

YouTube - s and p orbitals, orbitais s e p

The shapes and labels of the five 3d orbitals.

“D” Orbitals

YouTube - atomic d orbitals

f Orbitalsf Orbitalsf Orbitalsf Orbitals

f sublevel with 7 orbitalsf sublevel with 7 orbitals

F orbitals

YouTube - Electron Orbitals - s,p & d

s orbitalss orbitals d orbitalsd orbitals

Number ofNumber oforbitalsorbitals

Number of Number of electronselectrons

11 33 55

22 66 1010

p orbitalsp orbitals f orbitalsf orbitals

77

1414

How many electrons can be in a sublevel?How many electrons can be in a sublevel?

Remember: A maximum of two electrons can Remember: A maximum of two electrons can be placed in an orbital.be placed in an orbital.

Orbital Diagrams

• Rules to Remember: – Aufbau principle: fill lower

energy levels first– Pauli Exclusion: two

electrons in each orbital with opposite spin.

– Hund’s Rule: One electron in One electron in each orbital of an energy each orbital of an energy level, then you can pair them level, then you can pair them upup

Orbital DiagramsGraphical representation of an

electron configurationEach arrow represents one

electronShows spin(opposite arrow)

and which orbital the electron is in within a sublevel

Must put 1 electron in each orbital of a sublevel before pairing them up

Exceptions- d4 and d9, two electrons in each orbital, etc.

LithiumLithiumLithiumLithium

Group 1AGroup 1AAtomic number = 3Atomic number = 31s1s222s2s11 ---> 3 total electrons ---> 3 total electrons

1s

2s

3s3p

2p

CarbonCarbonCarbonCarbonGroup 4AGroup 4AAtomic number = 6Atomic number = 61s1s2 2 2s2s2 2 2p2p22 ---> ---> 6 total electrons6 total electrons

Here we see for the first time Here we see for the first time

HUND’S RULEHUND’S RULE. When . When placing electrons in a set of placing electrons in a set of orbitals having the same orbitals having the same energy, we place them singly energy, we place them singly as long as possible.as long as possible.1s

2s

3s3p

2p

Lanthanide Element Lanthanide Element ConfigurationsConfigurations

4f orbitals used for Ce - Lu and 5f for Th - Lr

4f orbitals used for Ce - Lu and 5f for Th - Lr

1s

2s2p

3s3p

3d 4s4p

energy

Oxygen (O)Chromium (Cr)Krypton (Kr)

Draw these orbital diagrams!

Why do “d” and “f” orbitals fill in out of order? d and f orbitals require LARGE amounts

of energy So…. the electrons fill in the next “s”

sublevel before they fill in the previous “d” or “f” sublevels because it’s a lower energy sublevel.

Electron Configurations

2p4

Energy LevelEnergy Level

SublevelSublevel

Number of Number of electrons in electrons in the sublevelthe sublevel

1s1s22 2s 2s22 2p 2p66 3s 3s22 3p 3p66 4s 4s22 3d 3d1010 4p 4p66 5s 5s22 4d 4d1010 5p 5p66 6s6s22 4f 4f1414…… etc.etc.

Electron Configurations Show how the electrons are arranged Periodic Table can be used as a “road

map” to map out electron configurations

Parts of the Periodic Table

Columns = Groups or Families-Have same # of valence electrons-Have similar properties

Rows = Periods or Series-Have same number of energy levels

Electrons and the Periodic Table You can use the periodic table to

determine how the electrons are arranged (electron configuration)

The Periodic Table shows sublevels and orbitals

Orbitals and the Orbitals and the Periodic TablePeriodic Table

s orbitals orbital(1orbital)(1orbital) p orbitalsp orbitals

(3 orbitals)(3 orbitals)

d orbitalsd orbitals(5 orbitals)(5 orbitals)

f orbitalsf orbitals(7 orbitals)(7 orbitals)

1

2

3

4

5

6

7

6

7

group # = # valence (outside) e-

d p

f

sRow=

# shells

1

2

3

4

5

6

7

6

7

peri

od

# =

# e

- sh

ells

1A2A

3B4B5B6B7B8B8B8B1B2B

3A4A5A6A7A8Agroup # = # valence e-

d

f

3d

4d

5d

6d

4f

5f

Subshells d and f are “special”

Let’s Try It!Write the electron configuration for

the following elements:HLiNNeKZnPb

Principal Quantum number (energy level)

sublevels Orbitals per sublevel

Orbitals per energy level

Number of electrons per sublevel

Total Electrons in energy level

1 s 1 1 2 2

2 s

p

1

3

4

2

6

8

3 s

p

d

1

3

5

9

2

6

10

18

4 s

p

d

f

1

3

5

7

16

2

6

10

14

32

Ws. Electron Configurations

Diagonal Rule The diagonal rule is a memory device that

helps you remember the order of the filling of the orbitals from lowest energy to highest energy

Order of Electron Subshell Filling:It does not go “in order”

1s2

2s2 2p6

3p6

4p6

5p6

6p6

7p6

3s2

4s2

5s2

6s2

7s2

3d10

4d10

5d10

6d10

4f14

5f14

1s2 2s22p6 3p63s2 4s2 4p65s23d10 5p66s24d10 6p67s25d104f14 7p66d105f14

Diagonal Ruless

s 3p 3ds 3p 3d

s 2ps 2p

s 4p 4d 4fs 4p 4d 4f

s 5p 5d 5f 5g?s 5p 5d 5f 5g?

s 6p 6d 6f 6g? 6h?s 6p 6d 6f 6g? 6h?

s 7p 7d 7f 7g? 7h? 7i?s 7p 7d 7f 7g? 7h? 7i?

11

22

33

44

55

66

77

Steps:Steps:

1.1. Write the energy levels top to bottom.Write the energy levels top to bottom.

2.2. Write the orbitals in s, p, d, f order. Write Write the orbitals in s, p, d, f order. Write the same number of orbitals as the energy the same number of orbitals as the energy level.level.

3.3. Draw diagonal lines from the top right to the Draw diagonal lines from the top right to the bottom left.bottom left.

4.4. To get the correct order, To get the correct order,

follow the arrows!follow the arrows!

By this point, we are past By this point, we are past the current periodic table the current periodic table so we can stop.so we can stop.

Shorthand NotationA way to abbreviating electron configurations

Since we are only concerned about the outermost electrons, we can skip to places we know are completely full (noble gases), and then finish the configuration

Shorthand NotationStep 1: Find the closest noble gas to the atom (or ion), WITHOUT GOING OVER the number of electrons in the atom (or ion). Write the noble gas in brackets [ ].

Step 2: Finding the next energy level after noble gas

Step 3: Resume the configuration until it’s finished.

Shorthand Notation Chlorine

Longhand is 1s2 2s2 2p6 3s2 3p5

You can abbreviate the first 10 electrons with a noble gas, Neon. [Ne] replaces 1s2 2s2 2p6

The next energy level after Neon is 3So you start at level 3 on the diagonal rule (all levels start with s) and finish the configuration by adding 7 more electrons to bring the total to 17

[Ne] 3s2 3p5

Practice Shorthand NotationWrite the shorthand notation

for each of the following atoms:

ClKCaIBi

Electron Configurations

A list of all the electrons in an atom (or ion) Used to determine valence electrons.valence electrons. Valence electrons are the electrons in the Valence electrons are the electrons in the

outermost energy level.outermost energy level. The number of valence electrons determines

bonding capabilities Valence electrons are the electrons in the last last

“s” and “p” sublevels.“s” and “p” sublevels.

1s1s22 2s 2s22 2p 2p66 3s 3s22 3p 3p66 4s 4s22 3d 3d1010 4p 4p66 5s 5s22 4d 4d1010 5p 5p66 6s 6s22 4f 4f1414…… etc.etc.

Valence ElectronsValence ElectronsValence ElectronsValence ElectronsElectrons are divided between Electrons are divided between core core and and valence valence

electronselectronsB 1sB 1s22 2s 2s22 2p 2p11

Core = 1sCore = 1s22 , , valence = 2svalence = 2s22 2p 2p11

Br [Ar] Br [Ar] 4s4s22 3d 3d1010 4p 4p55

Core = [Ar] 3dCore = [Ar] 3d1010 , , valence = 4svalence = 4s22 4p 4p55

Electron Dot DiagramsThe The core electronscore electrons are represented are represented by the by the symbol symbol of the atom.of the atom.

The The valence electronsvalence electrons are are represented with a represented with a dotdot..

Rules of the GameRules of the GameRules of the GameRules of the Game

Atoms like to either empty or fill their outermost Atoms like to either empty or fill their outermost level.level. Since the outer level contains two s Since the outer level contains two s electrons and six p electrons (d & f are always in electrons and six p electrons (d & f are always in lower levels), the optimum number of electrons lower levels), the optimum number of electrons is eight. This is called theis eight. This is called the octet rule.octet rule.

Let’s Review….Let’s Review…. S sublevels have 1 orbital P sublevels have 3 orbitals D sublevels have 5 orbitals F sublevels have 7 orbitals

EACH ORBITAL CAN ONLY HOLD 2 ELECTRONS (in opposite spin)

The End !!!!!!!!!!!!!!!!!!!