Atomic Structure Hl

7/27/2019 Atomic Structure Hl

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ElectronConfiguration

IB Chemistry Power Points

Topic 12

Atom ic Structure

www.pedagogics.ca


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HL Topic 12.1 Electron Configuration

Ionization energy is defined as the amount ofenergy required to remove a mole of electrons froma mole of gaseous atoms of a particular element.

E(g)

E+

(g) + e-


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Trends in ionization energy occur in the Periodic Table.Ionization energy decreases down a group and increasesacross a period.


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Trends in ionization energy occur in the Periodic Table.Ionization energy decreases down a group and increasesacross a period.

WHY?


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Effective nuclear charge is the net positive charge felt by anelectron in an atom.

The basics of electrostatics imply that each and every electron is attracted to each andevery proton in the nucleus and repelled by every other electron. However . . .

Electrons between the valence electron and the nucleusprovide a shielding effectweakening the electrostatic force

on the valence electron.


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Across a Period:

shielding remains constant

atomic number increases so effective nuclear charge increasesionization energy increases


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Down a Group:

shielding increases AND atomic number increaseseffective nuclear charge does not change significantlyvalence electrons further from nucleusso weaker electrostatic force and lower ionization energy


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Lithium (Z=3)

Sodium (Z=11)

Hydrogen (Z=1)+ e

-

+

+

+

e-

8e-

2e-

++++

e-

+

e-

H+

+

++ e

-

2e-

Li+

e-

8e-

2e-

++++

Na+

2e-


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This diagram shows how 1stionization energy decreases

down a group and that trends in ionization energy also occuracross a period. Look at the H, Li, Na, K, Rb, Cs values!

.


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Looking at just the trend across the 1stperiod, what does the

graph imply?The theory is . . .

Across a period, number ofp

+increases so effective

nuclear charge increases.

As a result, the valenceelectrons are morestrongly held, and arecloser to the nucleus(radius decreases)

This does not explain thedrop in ionization energy(decreased stability)

observed between Be and Band between N and O.


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NEW IDEA suborbitals (or subshells)

Within a given energy level(shell), there are differentsubshells that electrons canoccupy that have slightlydifferent energy levels


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Old Idea expanded: 2,8,8,2 configuration with energy shells1,2,3,4 . . . but each shell subdivided

Electron arrangement in atoms can be described by terms

called quantum numbers no two electrons can have thesame quantum number!

1

st

Term: Shell(n)- principle energy level

n = 1

n = 2

n = 3

lone electronof Hydrogen

2

nd

Term: subshell- designated by s, p,d,f

1s

n = 2

n = 3

The first energy shell (1) has one subshell (s).


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2nd

Term: subshell- designated by s, p, d, f- designates the sub-energy level

within the shell.- refers to the shape(s) of the

volume of space in which theelectron can be located.

1s

n = 2

n = 3

The first shell (1) has one subshell (s).

The ssubshell has 1 spherical shaped orbital

orbitals are volumes of space where theprobability of finding an electron is high


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The Electronic Configuration of Hydrogen

1s

Hydrogen has one electron located in the firstshell (1). (Aufbau principle)

The first shell has only one subshell (s). The ssubshell contains 1 spherical orbital.

1s1

shell

subshell

# of electrons present

energy

1s

Electronic configuration

Orbital Energy Level Diagram


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The Electronic Configuration of HeliumHe: Atomic # of 2, 2 electrons in a neutral He atom

H 1s1

He 1s

2

He 1s1s

the maximum number of electrons in an orbital is TWO

if there are 2 electrons in the same orbital they must have anopposite spin.

This is called Paulis Exclusion Principle


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1s

Lithium (Li)Li: Z=3 Li has 3 electrons.

2nd

shell

1s

The 2nd shell (n= 2) has 2subshells which are sandp.

The s subshell fills first!(Aufbau Principle)

2s

2p

Li 1s

2

2s

1

2s

Li 1s

Electronic configuration

Orbital Energy Level Diagram


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Berylium (Be)Be: Z=4 Be has 4 electrons.

Be 1s22s

2

2sBe 1s

Electronic configurationOrbital Energy Level Diagram

1s2

ndshell

2s 2p

B 1s

2

2s

2

2p

1

2p2s

B 1s

Boron (B) has 5 electrons, the s subshell is full so the 5th

electron occupies the first orbital in the p subshell


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Subshells so far- designated by s, and p- refers to the shape(s) of

the volume in which the electron

can be located.- also designates an energy level

within the shell.- relative energy: s


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Carbon (C)C: Z=6 C has 6 electrons.

1s2

ndshell

2s 2p C 1s

2

2s

2

2px

1

py

1

2p2s

C 1s

C 1s22s

22p

2

The 6th

electron occupies anempty p orbital. This illustratesHundsRule electrons do notpair in orbitals until each orbital

is occupied with a single electron.

The electron configuration is

But always written as

C l h filli f


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2p2s

N 1s1s

22s

22p

3

2p2s

O 1s 1s22s

22p

4

2p2s

Ne 1s 1s22s

22p

6

Can we relate the filling ofthe subshells with theionization energy data?

I i ti t d


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Ionization energy trends

Down a group : ionization energy decreases- ENC constant but atoms larger so easier to ionize

Across a period : ionization energy increases- increasing ENC therefore smaller size (e- closer to nucleus)so harder to ionize


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Explaining the dips support for sandporbital model

Be to B dip- because s shields p and lowers ENC

N to O dip- because repulsions between electronpair in first full orbital (experimentalevidence supporting Aufbau and Hund)


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Electron Configurations and the Periodic Table

So far, we have seen how the subshell model provides

and explanation for the patterns in ionization energywe see in the periodic table.

You have also seen how to write electron configurations

Example CALCIUM 1s2

2s2

2p6

3s2

3p6

4s2

Principle energy level subshell # of e-

Calcium can also be written shorthand as:

[Ar]4s2


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Practice

Use the sheets provided to fill out orbital diagrams and

determine the electron configuration for the followingelements

1. Fluorine

2. 56Fe

3. Magnesium - 224. 131I

5. Potassium 42

6. 75Ge

7. Zirconium 90

8.41

Ca2+

P actice


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Practice

Use the sheets provided to fill out orbital diagrams anddetermine the electron configuration for the following

elements

1. Fluorine 1s22s

2p

5

2.56

Fe 1s22s

2p

63s

23p

64s

23d

6

3. Magnesium

22 1s22s2p63s2

4.131

I 1s22s

2p

63s

23p

63d

104s

24p

64d

105s

25p

5

5. Potassium 42 1s22s

2p

63s

23p

64s

1

6.75

Ge 1s22s

2p

63s

23p

64s

23d

104p

2

7. Zirconium 90 1s22s

2p

63s

23p

64s

23d

104p

65s

24d

2

8.41

Ca2+

1s2

2s2

p6

3s2

3p6

Th i ti f th P i di t bl l t di tl t


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The organization of the Periodic table correlates directly toelectron structure

C d d l t fi ti f l th l t


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Condensed electron configurations for example the electronconfiguration of bromine can be written [Ar] 4s

23d

104p

5

Read questions carefully many IB questions require you

to write the FULL electron configuration

El t fi ti f i


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Electron configuration of ions:

The exception: TRANSITION METAL IONS

In general, electrons will be removed from orbitals (ionization) in thereverseorder that the orbitals were filled. In other words, electronsvacate higher energy orbitals first.

When these ions form, electrons are removed from the valence shell sorbitals before they are removed from valence dorbitals when transitionmetals are ionized.

For example: Cobalt has the configuration [Ar] 4s23d

7OR [Ar] 3d

74s

2

The Co2+

and Co3+

ions have the following electron configurations.

Co2+

: [Ar] 3d7 Co

3+: [Ar] 3d

6

Condensed electron configurations for example the electron


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Condensed electron configurations for example the electronconfiguration of bromine can be written [Ar] 4s

23d

104p

5

1. Si ___________________________

2. S2- ___________________________3. Rb+ ___________________________

4. Se ___________________________

5. Ar ___________________________6. Nb ___________________________7. Zn2+ ___________________________8. Cd ___________________________

9. Sb ___________________________

You are responsible for configurations up to Z 54 (Xe) The table


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You are responsible for configurations up to Z = 54 (Xe). The tableworks well for this with the exception of Cr and Cu


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Chromiums configuration is:

[Ar]4s13d

5

Coppers configuration is:

[Ar]4s1

3d10

These configurations are energetically more stable

than the expected arrangements. KNOW THEM!

Successive ionization


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1st

737.7

2nd

1450.7

3rd

7732.735458 31653 25661 21711 18020 13630 10542.5

169988189367.7

Successive ionizationenergy data supports theelectron configuration model


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Review: the principles involved

Hunds Rule: the most stable arrangement of electrons in

orbitals of equal energy is where there is the maximum number

of unpaired electrons all with the same spin.

Aufbau Principle: electrons will fill the lowest energy orbitalsfirst

Paulis Exclusion Principle: A maximum of two electrons can

occupy a single orbital. These electrons will have oppositespins.

Atomic Structure Hl

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