Ms. ClearyChem 11

A modelA representation or explanation of a reality

that is so accurate and complete that it allows the model builder to predict events.

Scientific Method leads to model buildingGather data, develop a model, formulate a

hypothesis, test and modify the model.

Bohr’s Model Why don’t the electrons fall into the nucleus?Move like planets around the sun.

In circular orbits at different levels.

Amounts of energy separate one level from another.

Bohr’s ModelNucleus

Electron

Orbit

Energy Levels

Nucleus

Electron

Orbit

Energy Levels

Bohr postulated that: Fixed energy related to the orbit Electrons cannot exist between orbits The higher the energy level, the further it is

away from the nucleus An atom with maximum number of

electrons in the outermost orbital energy level is stable (unreactive)

How did he develop his theory?He used mathematics to explain

the visible spectrum of hydrogen gas

http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/linesp16.swf

Radiowaves

Microwaves

Infrared .

Ultra-violet

X-Rays

GammaRays

Low energy

High energy

Low Frequency

High Frequency

Long Wavelength (700 nm)

Short Wavelength (400 nm)

Visible Light

Energy and Visible Light

The line spectrum electricity passed

through a gaseous element emits light at a certain wavelength

Can be seen when passed through a prism

Every gas has a unique pattern (color)

Line spectrum of various elements

Bohr’s Triumph His theory helped to explain periodic lawHalogens are so reactive because it has one

e- less than a full outer orbitalAlkali metals are also reactive because they

have only one e- in outer orbital

DrawbackBohr’s theory did

not explain or show the shape or the path traveled by the electrons.

His theory could only explain hydrogen and not the more complex atoms

Further away from the nucleus means more energy.

There is no “in between” energy

Energy Levels

First

Second

Third

Fourth

Fifth

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Complete Bohr Diagrams for the Following:MgLiNeF

The Quantum Mechanical Model Energy is quantized. It comes in chunks.A quanta is the amount of energy needed

to move from one energy level to another.Since the energy of an atom is never “in

between” there must be a quantum leap in energy.

Schrödinger derived an equation that described the energy and position of the electrons in an atom

Atomic OrbitalsPrincipal Quantum Number (n) = the

energy level of the electron.Within each energy level the complex

math of Schrödinger's equation describes several shapes.

These are called atomic orbitalsRegions where there is a high

probability of finding an electron

OrbitalsElectrons spin around the nucleus creating

an electron cloud.The electron clouds come in 4 different

shapes, called orbitals.The four orbitals are called s, p, d, and f.

Each orbital is capable of holding different numbers of electrons:

Orbital # of Electrons

s 2

p 6

d 10

f 14

S orbitals 1 s orbital forevery energy level

1s 2s 3s

Spherical shapedEach s orbital can hold 2 electronsCalled the 1s, 2s, 3s, etc.. orbitals

P orbitals Start at the second energy level 3 different directions3 different shapesEach orbital can hold 2 electrons

The p Sublevel has 3 p orbitals

The D sublevel contains 5 D orbitals The D sublevel starts in the 3rd energy level 5 different shapes (orbitals)Each orbital can hold 2 electrons

The F sublevel has 7 F orbitals The F sublevel starts in the fourth energy levelThe F sublevel has seven different shapes

(orbitals)2 electrons per orbital

Summary

s

p

d

f

# of shapes (orbitals)

Max # of electrons

1 2 1

3 6 2

5 10 3

7 14 4

Sublevel

Starts at energy level

Electron Configurations The way electrons are arranged in atoms.Aufbau principle- electrons enter the lowest

energy first.This causes difficulties because of the overlap

of orbitals of different energies.Pauli Exclusion Principle- at most 2 electrons

per orbital - different spins

Electron Configurations First Energy Levelonly s sublevel (1 s orbital)only 2 electrons1s2

Second Energy Levels and p sublevels (s and p orbitals are

available)2 in s, 6 in p2s22p6

8 total electrons

Levels Third energy level• s, p, and d orbitals• 2 in s, 6 in p, and 10 in d• 3s23p63d10

• 18 total electronsFourth energy level• s,p,d, and f orbitals• 2 in s, 6 in p, 10 in d, and 14 in f• 4s24p64d104f14

• 32 total electrons

Electron ConfigurationsElectron configurations are a shorthand

for writing exactly what was in the energy level diagrams.

Electron configuration for O is: 1s22s22p4

period orbital

# of electrons

Electron configuration for Ar is:1s22s22p63s23p6

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nerg

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1s

2s

3s

4s

5s6s

7s

2p

3p

4p

5p

6p

3d

4d

5d

7p 6d

4f

5f

Electron Configuration Hund’s Rule- When electrons occupy orbitals

of equal energy they don’t pair up until they have to .

The first to electrons go into the 1s orbital

Notice the opposite spins

only 13 more

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y

1s

2s

3s

4s

5s6s

7s

2p

3p

4p

5p

6p

3d

4d

5d

7p 6d

4f

5f

The next electrons go into the 2s orbital

only 11 more

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nerg

y

1s

2s

3s

4s

5s6s

7s

2p

3p

4p

5p

6p

3d

4d

5d

7p 6d

4f

5f

• The next electrons go into the 2p orbital

• only 5 more

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nerg

y

1s

2s

3s

4s

5s6s

7s

2p

3p

4p

5p

6p

3d

4d

5d

7p 6d

4f

5f

• The next electrons go into the 3s orbital

• only 3 more

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nerg

y

1s

2s

3s

4s

5s6s

7s

2p

3p

4p

5p

6p

3d

4d

5d

7p 6d

4f

5f

Incr

easi

ng e

nerg

y

1s

2s

3s

4s

5s6s

7s

2p

3p

4p

5p

6p

3d

4d

5d

7p 6d

4f

5f

• The last three electrons go into the 3p orbitals.

• They each go into separate shapes

• 3 unpaired electrons

• 1s22s22p63s23p3

Orbitals fill in order Lowest energy to higher energy.Adding electrons can change the energy of

the orbital.Half filled orbitals have a lower energy.Makes them more stable.Changes the filling order

Write these electron configurations

Titanium - 22 electrons

1s22s22p63s23p64s23d2

Vanadium - 23 electrons

1s22s22p63s23p64s23d3

Chromium - 24 electrons

1s22s22p63s23p64s23d4 is expectedBut this is wrong!!

Chromium is actually 1s22s22p63s23p64s13d5

Why?This gives us two half filled orbitals.Slightly lower in energy.The same principal applies to copper.

Copper’s electron configuration

Copper has 29 electrons so we expect1s22s22p63s23p64s23d9

But the actual configuration is1s22s22p63s23p64s13d10

This gives one filled orbital and one half filled orbital.

Remember these exceptions

Electron Configuration and the Periodic TableGroups 1 and 2 represent the s orbitalGroups 13-18 represent the p orbitalGroups 3-12 represent the d orbitalLanthanides and Actinides represent f

orbital

Practice1. Time to practice: Draw the following energy

level diagrams on your own filling up electron configurations: H, He, Be, N, Na, Ni, Br,

2. Do electron configurations for the elements listed in #1.