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Atomic Emission SpectraAtomic Emission Spectra

Read p. 54-59

What is light?

White light: reflection of all colorsBlack light: absorption of all colorsColors are each a different

wavelength (λ: lamda) of light

Colors Different wavelengths of light are seen

as different colors.

Question of the Day

Where are the electrons around the nucleus?

Study the light emitted (produced) by atoms and ions to deduce (find out) the structure of atoms.

When an atom is “excited” its electrons gain energy and move to a higher energy level. To return to a lower energy level, electrons must lose energy. They do this by giving off light.

Continuous spectrum:Continuous spectrum: all wavelengths of visible light contained in white light.

Light emitted by an atom can be separated into a line spectrumline spectrum that shows exactly what frequencies of light are present.

Because the light emitted from atoms is a line spectrumline spectrum (not a continuous spectrum) we determine that:

There are “discrete” (separate) energy energy levelslevels for each atom that can only produce light of certain wavelengths (this is NOT ordinary white light!).

c=fc=fλ λ (velocity of light = frequency x wavelength)

the greater the frequency the shorter the wavelength

ΔE = hfΔE = hf (energy lost by the electron = h(constant) x frequency

Frequency (and thus, color) of the light depends on the amount of energy lost by the electron.

Wavelength (λ) increases

Frequency (f) increases

ΔE = hfΔE = hf c=f c=fλ λ

Increasing frequency (f) (increasing energy)

Increasing frequency (f) (increasing energy)

c=fc=fλλ

When atoms are “exited” (energy is added) they produce light.

Not white or all-colored light, but one color at a time.

Different colors indicate (show) different energy levels.

Hydrogen Emission Spectrum

Only certain energy levels can occur (not a continuous spectrum)

Energy Level Diagram

The larger the difference in energy, the greater the frequency (thus, the more purple the light).

Incre

asin

g fre

qu

en

cy

Visible

Visible

Incre

asin

g fre

qu

en

cy

Closer to nucleus

Further from nucleus

Increasin

g frequ

ency

convergence:convergence: the lines in a spectrum converge (get closer together) as frequency increases. related to how much energy is required

to remove the electron from the atom (ionize)

Stop

If you are missing a notebook (lab or class/homework), please come to my desk at 5:10 pm today.

Extra help: Tuesdays 8:15-9:00 pm Thursdays 7:00-8:00 pmChem is try!!

When electrons are excited, they emit colors in a line spectrum. Only certain wavelengths of light are produced. Since c=flamda only certain frequencies are produced. Since ChangeE=hf, only certain changes in energy occur. Thus, electrons can not be making all changes in Energy, but only changes between discrete, separate energy levels.

Electronic StructureElectronic Structure

Energy LevelsShells

Electronic Structure and the Periodic Table

1st energy level = 22nd energy level = 8Electronic structure: number of

electrons in each energy levelAfter second level – more complicated

they don’t fill in order – more later

2 8 5

2 2

2 8

H=1O=2,6 (two electrons in the first energy

level, six in the second)Al=2,8,3S=Cl=Ar=

Different isotopes have the same electronic structure and the same chemical properties!

Electron BehaviorElectron BehaviorValence shell: outer energy level of

an atom determine the physical and chemical

properties of an atom

Valence Shell

12 3 4 5 6 7

8

Valence electrons

Period

Gro

up

How many electrons in valence shell? Al Ne Li Ca

Stop here

Electronic Structure of AtomsElectronic Structure of AtomsZumdahl2: p. 307-312

Electronic Structure

Energy levelsSub-levels

OrbitalSpin (electrons)

Energy Levels

Major shells (layers) around the nucleus filled before higher levels are filled 1st: 2 electrons 2nd: 8 electrons

Sub-levelsDifferent shapess – sphere

one orbitalp – figure eight

three orbitalsd –

five orbitalsf –

seven orbitals

p Sub-levelp sub-level has three orbitalspx, py, pz

d and f sub-levels have very complex shapes

Orbitals

Each orbital can hold two electrons.

Electrons spin in opposite directions

Energy Level

Types of sub-

levels

Total orbitals

Electron capacity

1 s 1 2

2 s, p 1+3=4

8

3 s, p, d 1+3+5=9

18

4 s, p, d, f 1+3+5+7=16

32

Energies of sub-levels

Energy States

Depending on where an electron was around the nucleus, it had different energy states.

Ground state: An orbital as close to the nucleus as possible: not very exited at all.

Excited State: An orbital farther awayfrom the nucleus: much morepotential for giving off energy.

Electron Configuration

The arrangement of electrons in an atom Each element’s atoms are different Arrangement with the lowest energy=

ground state electron configurationHow do we figure out what the

ground state electron configuration looks like?

How do we figure out where the electrons are?

1. Figure out the energy levels of the orbitals

2. Add electrons to the orbitals according to three rules

Three Rules for Electron Configuration

1: Aufbau Principle1: Aufbau Principle

An electron goes to the lowest-energy orbital that can take it.

2: Pauli Exclusion Principle

No two electrons can have exactly the same configuration (location). Can have the same orbital, but must

have opposite spins.

3. Hund’s RuleOrbitals of equal energy are each

occupied by one electron before any orbital is occupied by a second electron

All electrons that are by themselves in an orbital must have the same spin.

Electron configuration: the arrangement of electrons in an atom.

So…what do we do with this information?!

How do we write electron configurations?

Orbital NotationElectron Configuration NotationNoble Gas Notation (shorthand)

All ways to communicate where the electrons are in the ground state of any atom.

s – sphere one orbital – 2 electrons

p – figure eight three orbitals – 6 electrons

d – five orbitals – 10 electrons

f – seven orbitals – 14 electrons

Orbital Notation

Electron configuration goes below a line or box

Arrows representing electrons go on the lines or in the boxes

1s 2s 2p 3s 3p 4s

Fluorine? 9 electrons

Magnesium? 12 electrons1s 2s 2p 3s 3p 4s

1s 2s 2p 3s 3p 4s

Write the orbital notationHomework notebook!

1. carbon2. phosphorus3. potassium4. krypton5. rubidium6. chromium7. copper8. gold

Electron Configuration NotationMain energy level Sub-levelElectrons

Carbon (6 electrons) 1s22s22p2

Aluminum (13 electrons) 1s22s22p63s23p1

OxygenArgonCopperZirconium

Orbital Notation

Boron1s22s22p1

Atomic Number?How many electrons?Orbital notation

Aufbau Principle: start with 1s and work up in energy level

Pauli Exclusion Principle: No two elements can have the same arrangement of electrons

Hund’s Rule: Fill in one electron per orbital first, then go back.

Orbital Notation

Nitrogen1s22s22p3

Atomic Number?Number of Electrons?Orbital Notation?

Homework:

Write the orbital notation for:FluorineAluminumCarbonSilverBromineHeliumNickel

Noble Gas Notation (Shorthand)Noble gases have completely filled valence

shellsIf Ne (a noble gas) is 1s22s22p6

Na = [Ne]3s1.Sodium has one more electron than Neon,

so its Noble Gas Notation is Neon plus one electron in the s sublevel of the third energy level.

Mg = [Ne]3s2.

Practice

Fe Electron Configuration Notation Noble Gas Notation

K Electron Configuration Notation Noble Gas Notation

LiBe

Homework

Noble gas notation:CsSrMnMo

SL Chemistry Vocabblockchemical (adj.)physicalpropertymetal (metallic)halogennobleperiodicity

successiveionisationelectronegativityradiusattract (attraction)repel (repulsion)react (reactivity)

HL Chemistry VocabSL Vocab plus:quantumnatureprincipalazimuthalmagneticendothermicmagnituderepulsion

effectivesuccessivecounteracted

Special HL Class optional

Monday, April 77-9 pmRoom 409

s, p, d and f blocks of the periodic table

groups

periods

“Long form” periodic table

• Electronic structures are related to the position of the elements on the periodic table

• s-block: s orbitals are filled

• p-block: p orbitals are filled

• etc.

Zumdahl2 p. 322