AtomicStructure3

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Electromagnetic

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Electromagnetic Radiation

• Most subatomic particles behave as _{Most subatomic particles behave as}

PARTICLES and obey the physics of waves.

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wavelength

Visible light

wavelength

Ultaviolet radiation

Amplitude

Node

Electromagnetic

Radiation

Electromagnetic

Radiation

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• _{Waves have a frequency}

_{Waves have a frequency}

• _{Use the Greek letter “nu”,}_{Use the Greek letter “nu”,}



_{, for frequency,}_{, for frequency,}

and units are “cycles per sec”

• _{All radiation:}_{All radiation:}



_•_•



= c

where c = velocity of light = 3.00 x 10

where c = velocity of light = 3.00 x 1088 m/sec_m/sec

Electromagnetic Radiation

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Electromagnetic Spectrum

Long wavelength --> small frequency

Short wavelength --> high frequency

increasing

frequency

increasing

wavelength

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Electromagnetic

Spectrum

Electromagnetic

Spectrum

In increasing energy

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Excited Gases

& Atomic

Structure

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Atomic Line Emission

Spectra and Niels Bohr

Atomic Line Emission

Spectra and Niels Bohr

Bohr’s greatest contribution to

science was in building a simple

model of the atom. It was based

on an understanding of the

on an understanding of the LINE LINE EMISSION SPECTRA

EMISSION SPECTRA of excited of excited atoms.

atoms.

•Problem is that the model only _{Problem is that the model only}

works for H

works for H Niels Bohr

Niels Bohr

(1885-1962)

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Spectrum of White Light

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Line Emission Spectra

of Excited Atoms

Line Emission Spectra

of Excited Atoms

• Excited atoms emit light of only

certain wavelengths

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Spectrum of

Excited Hydrogen Gas

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Line Spectra of Other Elements

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Atomic Spectra

+

Electron orbit

One view of atomic structure in early 20th

century was that an electron (e-) traveled

about the nucleus in an orbit.

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Atomic Spectra and Bohr

Bohr said classical view is wrong.

Need a new theory — now called

QUANTUM

or

WAVE MECHANICS

.

e- can only exist in certain discrete

orbits

e- is restricted to

QUANTIZED

energy

state (quanta = bundles of energy)

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Schrodinger applied idea of e-

behaving as a wave to the

problem of electrons in atoms.

He developed the

WAVE

EQUATION

Solution gives set of math

expressions called

expressions called WAVE _WAVE

FUNCTIONS,



Each describes an allowed energy

state of an

state of an e-E. Schrodinger

E. Schrodinger

1887-1961

Quantum or Wave Mechanics

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Heisenberg Uncertainty

Principle

• The problem of defining nature of electrons in atoms solved by W. Heisenberg.

• He observed that on cannot simultaneously define the position and momentum (= m•v) of an electron.

• If we define the energy exactly of an electron precisely we

must accept limitation that we do not know exact position.

• The problem of defining nature of electrons in atoms solved by W. Heisenberg.

• He observed that on cannot simultaneously define the position and momentum (= m•v) of an electron.

• If we define the energy exactly of an electron precisely we

must accept limitation that we do not know exact position.

W. Heisenberg

1901-1976

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Arrangement of

Electrons in Atoms

Arrangement of

Electrons in Atoms

Electrons in atoms are arranged as

LEVELS

(n) (n)

SUBLEVELS

(l) (l)

ORBITALS

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QUANTUM NUMBERS

The

The shape, size, and energyshape, size, and energy of each orbital is a function of each orbital is a function of 3 quantum numbers which describe the location of

of 3 quantum numbers which describe the location of

an electron within an atom or ion

n

(principal)

(principal) ---> energy level---> energy level

l

(orbital) (orbital) ---> shape of orbital---> shape of orbital

m

_l_l (magnetic)(magnetic) ---> designates a particular ---> designates a particular suborbital

suborbital

The fourth quantum number is not derived from the

wave function

s

(spin)(spin) ---> spin of the electron ---> spin of the electron

(clockwise or counterclockwise: ½ or – ½)

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QUANTUM NUMBERS

So… if two electrons are in the same place at

the same time, they must be repelling, so at

least the spin quantum number is different!

The

The Pauli Exclusion PrinciplePauli Exclusion Principle says that says that no two no two electrons within an atom (or ion) can have the

electrons within an atom (or ion) can have the

same four quantum numbers.

If two electrons are in the same energy level,

the same sublevel, and the same orbital, they

must repel.

Think of the 4 quantum numbers as the address

of an electron… Country > State > City >

Street

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Energy Levels

• Each energy level has a number

_{Each energy level has a number}

called the

PRINCIPAL

_PRINCIPAL

QUANTUM NUMBER, n

• Currently n can be 1 thru 7,

_{Currently n can be 1 thru 7,}

because there are 7 periods on

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Energy Levels

n =

₁

1 n = 2

n = 2

n = 3

n = 4

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Types of Orbitals

• The most probable area to find

_{The most probable area to find}

these electrons takes on a shape

• So far, we have 4 shapes. They

_{So far, we have 4 shapes. They}

are named s, p, d, and f.

• No more than 2 e

_{No more than 2 e}

--

can be assigned

_{can be assigned}

to any single orbital – one spins

clockwise, one spins

counterclockwise

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Types of Orbitals (

l

)

s orbital

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p Orbitals

The

p sublevel

has has

3 orbitals

They are designated as p

They are designated as p_x_x, p, p_y_y, , and p

and p_z_z

The

p sublevel

has has

3 orbitals

They are designated as p

They are designated as p_x_x, p, p_y_y, , and p

and p_z_z planar node

Typical p orbital

There is a

There is a PLANAR PLANAR

NODE

NODE thru the thru the

nucleus, which is

an area of zero

probability of

finding an electron

3p

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p Orbitals

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d Orbitals

• d sublevel has 5

_{d sublevel has 5}

orbitals

typical d orbital

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f Orbitals

For l = 3,

---> f sublevel with 7 orbitals

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Diagonal Rule

• Must be able to write it for the test! This will be question #1 ! Without it, you will not get correct answers !

• 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

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Diagonal Rule

s

s 3p 3d

s

2p

s 4p 4d 4f

s 5p 5d 5f

5g?

s 6p 6d

6f 6g? 6h?

s 7p

7d 7f 7g? 7h? 7i?

1

2

3

4

5

6

7

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.

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!

By this point, we are past

the current periodic table

so we can stop.

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Why are d and f orbitals always

in lower energy levels?

• d and f orbitals require LARGE amounts of energy

• It requires less in energy to skip a

sublevel that requires a large amount of energy such as d and f orbtials for one in a higher level but lower energy such as a s or p orbital

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s orbitals

s orbitals d orbitalsd orbitals

Number of Number of orbitals orbitals Number of Number of electrons electrons p orbitals

p orbitals f orbitalsf orbitals

How many electrons can be in a sublevel?

Remember: A maximum of two electrons can

be placed in an orbital.

1 3 5 7

2

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Electron Configurations

A list of all the electrons in an atom (or ion)

• Must go in order (Aufbau principle)

• 2 electrons per orbital, maximum

• We need electron configurations so that we can

determine the number of electrons in the outermost energy level. These are called valence electrons.

• The number of valence electrons determines how

many and what this atom (or ion) can bond to in order to make a molecule

1s

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Electron Configurations

2p

4

Energy

Energy LevelLevel

Sublevel

Number of

1010

5p

_5p

66

6s

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Let’s Try It!

• Write the electron configuration for

the following elements:

H

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An excited lithium atom emitting a

photon of red light to drop to a

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Orbitals and the Periodic

Table

• Orbitals grouped in s, p, d, and f orbitals _{Orbitals grouped in s, p, d, and f orbitals}

(sharp, proximal, diffuse, and fundamental)

s orbitals

p orbitals

d orbitals

f orbitals

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Shorthand Notation

• A way of abbreviating long electron

configurations

• Since we are only concerned about the outermost electrons, we can skip to

places we know are completely full, i.e. the noble gases , and then finish the

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Shorthand Notation

• Step 1: Find the closest noble gas to the atom (or ion), WITHOUT GOING OVER the number of electrons in the atom (or ion). It will be at the end of the period above the element that you are working with Write the noble gas in brackets [ ].

• Step 2: Find where to resume by finding the next energy level.

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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 3

So 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

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Practice Shorthand Notation

• Write the shorthand notation for

each of the following atoms:

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Valence Electrons

Electrons are divided between core and

valence electrons

B

B 1s1s22 2s2s22 2p 2p11

Core = [He]

Core = [He] , , valence = 2svalence = 2s22 2p 2p11

Br

Br [Ar]_[Ar] 4s_4s22 3d3d1010 4p4p55

Core = [Ar] 3d

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Rules of the Game

No. of valence electrons of a main group

atom = Group number (for A groups)

Atoms like to either empty or fill their outermost

level. Since the outer level contains two s

electrons and six p electrons (d & f are always in

lower levels), the optimum number of electrons

is eight. This is called the

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Keep an Eye On Those Ions!

• _{Electrons can be lost or gained by}

atoms to form ions

• _{negative ions have gained}

electrons, positive ions have lost

electrons

• _{The electrons that are lost or}

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Forming Ions!

• Tin

Try Some Ions!

• Write the longhand notation for these:

F

-Li

+

Mg

+2

• Write the shorthand notation for these:

Br

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Exceptions to the Aufbau

Principle

• _{Remember d and f orbitals require LARGE}

amounts of energy

• _{If we can’t fill these sublevels, then the next}

best thing is to be HALF full (one electron in each orbital in the sublevel)

• _{There are many exceptions, but the most}

common ones are d4 and d9

For the purposes of this class, we are going to assume that ALL atoms (or ions) that end in d4

or d9 are exceptions to the rule. This may or

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Exceptions to the Aufbau Principle

• d4 is one electron short of being HALF full

• In order to become more stable (require less energy), one of the closest s electrons will

actually go into the d, making it d5 instead of d4.

• For example: Cr would be [Ar] 4s2 3d4, but

since this ends exactly with a d4 it is an

exception to the rule. Thus, Cr should be [Ar] 4s1 3d5.

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Exceptions to the Aufbau Principle

• s orbitals can also be subject to exceptions to the Aufbau process, especially in the d block elements

• Remember, half full is OK. When an s orbital loses 1 electron, it too becomes half full!

• Having the s half full and the d half full is

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Exceptions to the Aufbau Principle

• d9 is one electron short of being full

• Just like d4, one of the closest s electrons will

go into the d, this time making it d10 instead of

d9.

• Example: Au would be [Xe] 6s2 4f14 5d9, but

since this ends exactly with a d9 it is an

exception to the rule. Thus, Au should be [Xe]

6s1 4f14 5d10.

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Try These!

• _{Write the shorthand}

notation for:

Cu

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Orbital Diagrams

• Graphical representation of an

electron configuration

• One arrow represents one electron

• Shows spin and which orbital

within a sublevel

• Same rules as before (Aufbau

principle, d

4

and d

9

exceptions, two

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Orbital Diagrams

• _{One additional rule:}_Hund’s

Rule

– In orbitals of EQUAL ENERGY

(p, d, and f), place one electron in each orbital before making any pairs

– All single electrons must spin

the same way

• Think of this rule as the

“Monopoly Rule”

• In Monopoly, you have to build

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Lithium

Group 1A

Atomic number = 3

1s

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Carbon

Group 4A

Atomic number = 6

1s

2 2

2s

2 2

2p

22

--->

6 total electrons

Here we see for the first time

HUND’S RULE

.. When When

placing electrons in a set of

orbitals having the same

energy, we place them singly

as long as possible.

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Lanthanide Element

Configurations

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

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Draw these orbital diagrams!

• Oxygen (O)

• Chromium (Cr)

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Ion Configurations

To form anions from elements, add 1 or more

e- from the highest sublevel.

P [Ne] 3s

P [Ne] 3s22 3p 3p33 + 3e- ---> P + 3e- ---> P3-3- [Ne] 3s [Ne] 3s22 3p 3p66 or [Ar] or [Ar]

1s 2s

3s 3p 2p

1s 2s

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