Bohr model and electron configuration

Mrs. A. Kay

Chem 11

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 Model

Nucleus

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

Short WavelengthVisible Light

The line spectrum

electricity passed through a gaseous element emits light at a certain wavelengthCan be seen when passed through a prismEvery gas has a unique pattern (color)

Line spectrum of various elements

Bohr’s Triumph

His theory helped to explain periodic law

Halogens are so reactive because it has one e- less than a full outer orbital

Alkali metals are also reactive because they have only one e- in outer orbital

Drawback

Bohr’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 LevelsFirst

Second

Third

Fourth

Fifth

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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 Orbitals

Principal 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 orbitals

Regions where there is a high probability of finding an electron

S orbitals

1 s orbital for

every energy level

1s 2s 3s

Spherical shaped

Each s orbital can hold 2 electrons

Called the 1s, 2s, 3s, etc.. orbitals

P orbitals

Start at the second energy level

3 different directions

3 different shapes

Each orbital can hold 2 electrons

The p Sublevel has 3 p orbitals

The D sublevel contains 5 D orbitalsThe 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 level

The 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

Third energy level

s, p, and d orbitals

2 in s, 6 in p, and 10 in d

3s23p63d10

18 total electrons

Fourth 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

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1s

2s

3s

4s

5s6s

7s

2p

3p

4p

5p

6p

3d

4d

5d

7p 6d

4f

5f

Electron ConfigurationHund’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 moreIncr

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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

Incr

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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

Incr

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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

Incr

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1s

2s

3s

4s

5s6s

7s

2p

3p

4p

5p

6p

3d

4d

5d

7p 6d

4f

5f

Incr

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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 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 expected

But this is wrong!!

Chromium is actually1s22s22p63s23p64s13d5

Why?

This gives us two half filled orbitals.

Slightly lower in energy.

The same principal applies to copper.

Copper’s electron configurationCopper has 29 electrons so we expect

1s22s22p63s23p64s23d9

But the actual configuration is

1s22s22p63s23p64s13d10

This gives one filled orbital and one half filled orbital.Remember these exceptions

Great site to practice and instantly see results for electron configuration.

Practice

Time to practice on your own filling up electron configurations.

Do electron configurations for the first 20 elements on the periodic table.