Electron Configuration and the Periodic Table

Mallard Creek Chemistry - Rines

Electromagnetic Radiation

Wave Nature of Light

Property of Waves Frequency

▪ No. of waves per second

Wave Length▪ Distance between

corresponding points in a wave

Amplitude▪ Size of the wave

peak

Electromagnetic Radiation

Mathematical Relations

C = speed of light = 3.0 x 108 m/s

λ (lamda) = wavelength (m)

f= frequency (Hz or s-1)This is how we know what color light is emitted!

C = λ f

Frequency is inversely proportional to Wavelength

If λ increases f decreasesIf f increases λ decreases

Speed of the wave is always constant at 3.0 x 108

m/s

Bohr ModelNucleus: Neutrons and

Protons Orbits: Electrons

We know both specific energy and location of each electron

Electrons orbit the nucleus in certain fixed energy levels (or shells)Nucleus

Energy Levels

Bohr Model Bohr’s Atomic

Model of Hydrogen

Bohr - electrons exist in energy levels AND defined orbits around the nucleus.

Each orbit corresponds to a different energy level.

The further out the orbit, the higher the energy level

Bohr’s Model The Photoelectric

Effect Light releases electrons Not all colors work

Atomic Emission Spectra Hydrogen gas emitted

specific bands of light Bohr’s calculated

energies matched the IR, visible, and UV lines for the H atom

12

3456

Electromagnetic Radiation Photoelectric Effect – There is a

minimum frequency to eject the electron

Electromagnetic Radiation Photoelectric

Effect Only explained by “energy

packets” of light called a quantum

Quantum - minimum amount of energy that can be gained or lost by an atom

Photons are massless particles of light of a certain quantum of energy Based on the frequency and

wavelength of the photon

Bohr’s Model Excited electrons

Energy added to atom – electrons “jump” up energy levels

When the atom relaxes - electron “falls” to lower energy levels and emits photon

Bohr Model of hydrogen

Reference Sheets!!!!!

Electromagnetic Radiation

Atomic Line Spectra

Electrons in an atom add energy to go to an “excited state”.

When they relax back to the ground state, they emit energy in specific energy quanta

Electromagnetic Radiation

These observations suggested that electrons must exist in defined energy levels

Next, the excited electron relaxes to a lower excited state or ground state

First, the electron absorbs energy and jumps from the ground state to an excited state

5 ______

4 ______

3 ______

2 ______

1 ______

5 ______

4 ______

3 ______

2 ______

1 ______

hv

5 ______

4 ______

3 ______

2 ______

1 ______

hv

Electromagnetic Radiation

Particle Nature of Light

Wave nature could not explain all observations (Plank & Einstein)Photoelectric Effect

When light strikes a metal electrons are ejected

Atomic Line Spectra▪ When elements are

heated, they emit a unique set of frequencies of visible and non-visible light.

E = hf

Other Scientists Contributions

De Broglie

Heisenburg

Modeled electrons as waves

Heisenberg Uncertainty Principle: states one cannot know the position and energy of an electron

Electrons exist in orbital’s of probability

Orbital - the area in space around the nucleus where there is a 90% probability of finding an electron

Other Scientists Contributions

Schrödinger Schrödinger Wave Equation - mathematical solution of an electron’s energy in an atom

Quantum Mechanical Model of the atom – current model of the atom treating electrons as waves.

Quantum Mechanical Model

Nucleus: Neutrons and protons

Orbitals: region in space surrounding the nucleus where there is a 95% probability of finding an electron.

We know either energy or location of each electron.

Solutions to the Wave Equation

Quantum Numbers

Wave Equation generates 4 variable solutions n - size l - shape m - orientation s – spin

Address of an electron

Quantum Numbers n – Primary

Quantum Number Describes the size and energy of the orbital

n is any positive # n = 1,2,3,4,…. Found on the periodic

table Like the “state” you

live in

Quantum Numbers

l – Orbital Quantum Number

Sub-level of energy

Describes the shape of the orbital

l = 0,1,2,3,4,….(n-1)

“City” you live in

n = 3l = 0,1,2n = 2

l = 0,1n =

1l = 0

Quantum Numbers

l – Orbital Quantum Number

# level = # sublevels 1st level – 1 sublevel 2nd level – 2 sublevels 4th level = 4 sublevels

Quantum Numbers

s l = 0 Spherical in shape

p l = 1 Dumbbell in shape

d l = 2

f l = 3

s p d f

Sublevels are named for their shape

Quantum Numbers

m – Magnetic Quantum Number

Describes the orientation of the orbital in space

Also denotes how many orbital's are in each sublevel

For each sublevel there are 2l +1 orbital's

“Street” you live on

Quantum Numbers

Look at Orbital's as Quantum Numbers

l = 0 m = 0Can only be one s orbital

l = 1 m = -1, 0, +1For each p sublevel there are 3 possible orientations, so three

3 orbital's

Orbital DesignationsOrbital

Designation

n l M2l+1

No. of Orbita

l

No. of Electro

n3d 3 2 -2,-1,0,+1,+2 5 103p 3 1 -1,0,+1 3 63s 3 0 0 1 22p 2 1 -1,0,+1 3 62s 2 0 0 1 21s 1 0 0 1 2

Orbital RulesEnergy Level

Possible sub-

levels

Number of Sub-levels

n

No. of Orbitals

n2

No. of Electron

s2n2

4 s, p, d, f 4 16 32

3 s, p, d 3 9 18

2 s, p 2 4 8

1 s 1 1 2

Reflection How is the Bohr model different from

the earlier models of the atom? Who contributed to the modern model

of the atom? How is it different from Bohr’s?

Why do atoms give unique atomic line spectra?

What are ground and excited states? Is 2d possible? 4f ? 2s ? 6p? 1p? How many total orbital's in the 2nd

level? 4th level.

Aufbau Principle Aufbau

Principal Lowest energy orbital available fills first

“Lazy Tenant Rule”

Pauli’s Exclusion Principle

Pauli Exclusion Principle

No two electrons have the same quantum #’s

Maximum electrons in any orbital is

two ()

Hund’s Rule Hund’s

Rule When filling degenerate

orbital's, electrons will fill an empty orbital before pairing up with another electron.

Empty room rule

RIGHT WRONG

Periodic Table & Electron Configuration

Periodic Table & Electron Configuration

Using the periodic table for the filling order of orbitals, by going in atomic number sequence until

you use all the needed electrons in the element

Orbital Energy Diagram

d ______ ______ ______ ______ ______

p ______ ______ ______

3 s ______

p ______ ______ ______

2 s ______

1 s ______

An energy diagram for the first 3 main energy

levels

Level (n)

Sub-level (l)

Orbitals (m)

Incr

easi

ng E

nerg

y

Orbital Energy Diagram and Electron Configuration

p ______ ______ ______

3 s ______

p ______ ______ ______

2 s ______

1 s ______

An energy diagram for Neon

Incr

easi

ng E

nerg

y

Electron Spin

1s2

2s2 2px22py

22pz2

2p61s2

2s2

Electron Configuration Notation

Orbital Notation Orbital Notation shows each orbital O (atomic number 8)

____ ____ ____ ____ ____ ____ 1s 2s 2px

2py 2pz

3s

1s22s22p4 electron configuration!

Orbital Notation Orbital Notation shows each orbital O (atomic number 8)

____ ____ ____ ____ ____ ____ 1s 2s 2px

2py 2pz

3s

!

Orbital Notation Write the orbital notation for S S (atomic number 16)

___ __ __ __ __ __ __ __ __ 1s 2s 2p 3s 3p

1s22s22p63s23p4

How many unpaired electrons does sulfur have?2 unpaired electrons!

Valence Electrons Valence Electrons

As (atomic number 33) 1s22s22p63s23p64s23d104p3

The electrons in the outermost energy level.

s and p electrons in last shell

5 valence electrons

Shorthand Configuration

S 16e-

Valence Electrons

Core Electrons

S 16e- [Ne] 3s2 3p4

1s22s22p63s23p4

Valence Electrons Longhand Configuration

[Ar]

1 2 3 4 5 6 7

4s23d104p2

Noble Gas Configuration Example - Germanium

X X X X X X X X X X X X X

Electron ConfigurationLet’s Practice

P (atomic number 15) 1s22s22p63s23p3

Ca (atomic number 20) 1s22s22p63s23p64s2

As (atomic number 33) 1s22s22p63s23p64s23d104p3

W (atomic number 74) 1s22s22p63s23p64s23d104p65s24d105p66s24f145d4

Noble Gas Configuration[Ne] 3s23p3

[Ar] 4s2

[Ar] 4s23d104p3

[Xe] 6s24f145d4

Electron Configuration

Your Turn N (atomic number 7)

1s22s22p3

Na (atomic number 11) 1s22s22p63s1

Sb (atomic number 51) 1s22s22p63s23p64s23d104p65s24d105p3

Cr (atomic number 24) 1s22s22p63s23p64s23d4

Noble GasConfiguration[He] 2s22p3

[Ne] 3s1

[Kr]5s24d105p3

[Ar] 4s23d4

Stability Full energy level Full sublevel Half full sublevel

1 2 3 4 5 6 7

Exceptions Copper

Expect: [Ar] 4s2 3d9

Actual: [Ar] 4s1 3d10

Silver Expect: [Kr] 5s2 4d9

Actual: [Kr] 5s1 4d10

Chromium Expect: [Ar] 4s2 3d4

Actual: [Ar] 4s1 3d5

Molybdenum Expect: [Kr] 5s2 4d4

Actual: [Kr] 5s1 4d5

Exceptions are explained, but not

predicted!

Atoms are more stable with half full

sublevel

Stability Atoms create stability by losing, gaining

or sharing electrons to obtain a full octet Isoelectronic with noble gases

1 2 3 4 5 6 7

+1 +

2-3 -2 -1

0+3

+4

Atoms take electron configuration of the closest noble gas

Stability Na (atomic number 11)

1s22s22p63s1

1s22s22p6 = [Ne]1 2 3 4 5 6 7

Na

1 Valence electronMetal = Loses

Ne

Try Some P-3 (atomic number 15)

1s22s22p63s23p6

Ca+2 (atomic number 20) 1s22s22p63s23p6

Zn+2 (atomic number 30) 1s22s22p63s23p63d10

Lost valence electrons (s and p)

Full Octet

Lewis Structures Shows valence electrons only!

s & p electrons

1. Write noble gas configuration for the element

2. Place valence electrons around element symbol in order

X87

6

5

43

21 s electrons

p electrons

Try Some Write the Lewis structures for: Oxygen (O)

– [He] 2s2 2p4

Iron (Fe)– [Ar] 4s2 3d6

Bromine (Br)– [Ar] 4s2 3d10 4p5

O

Fe

Br

Valence electrons

• • • • •

•

• •

• •

• • • • •

What Do I Need to Know?

How the periodic table is arranged Be able to identify subcategories of

the periodic table How the elements within a group are

similar How the elements within a period are

similar Be able to compare and contrast the

electronegativities, ionization energies, and radii of metals and non-metals

Periodic TableDmitri Mendeleev – Father of the Periodic Table

What He Did Put elements in

rows by increasing atomic weight

Put elements in columns by similar properties

Some ProblemsHe left blank spaces for

what he said were undiscovered elements (he was right!)

He broke the pattern of increasing atomic weight to keep similar reacting elements together

MosleyArranged by Atomic # Columns = Groups

Rows = Periods

Periodic Table Organization

MetalloidsMetalsNon-Metals

Periodic Table Organization

Representative ElementsTransition MetalsInner Transition Metals

Metals and Non-metals

Metals Shiny Malleable Ductile (pulled into

wires) Conduct heat and

electricity Low specific heat High melting points Solids Lose electrons

Non-metals Dull Brittle Poor conductors Low melting/boiling

points Varied properties Varied phases

Atomic Radius Atomic Radius = ½ the distance between adjacent

nuclei Increases towards Francium

Ionic RadiusCations

Positive Ion Metals Lose electrons

Radius gets smaller!

Anions Negative Ion Non-metals Gain electrons

Radius gets larger!

K K+ Cl Cl-

Ionization Energy Energy required to remove an electron from

an atom Why are there peaks in this trend?

Ionization EnergyNoble gases have the highest first

Ionization Energy

Electronegativity Pull of electrons in a covalent bond “Attraction” of atoms towards an electron Fluorine is “the man”

Electronegativity Increases

Periodic TrendsNuclear Charge increasesAtomic radius decreasesIonization energy increasesElectronegativity increases

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Elec

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What Do I Need To Know?

How are electrons arranged in an atom

The two natures of electromagnetic radiation: Particles vs. Waves

How to use the periodic table to list the configuration or orbital diagram

What quantum numbers are and how they are related to electron configuration.

How the periodic table is arranged The basic periodic trends