CH 8: Electron Configuration & periodicity

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Vanessa N. Prasad-PermaulValencia Community College

CHM 1045

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Electron Configuration of Atoms

Electron Configuration of an atom: a particular

distribution of electrons among available subshells.

Li 3 electrons: 1s2 2s1

Orbital Diagram: a diagram that shows how the

orbitals of a subshell are occupied by electrons.

Li 3 electrons: 1s 2s

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Pauli Exclusion Principle: no two electrons in

an atom can have the same four quantum

Numbers. Rewritten: An orbital can hold at most

twoelectrons, and then only if the

electrons haveopposite spins.

SUBSHELL NUMBER OF ORBITALS

MAXIMUM NUMBER OF ELECTRONS

s (l = 0) 1 2p (l = 1) 3 6d (l = 2) 5 10f (l = 3) 7 14

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EXAMPLE 8.1Which one of the following orbital diagrams

or electron configurations are possible and which are impossible, according to the Pauli Exclusion Principle? Explain:

a. d. 1s32s1

1s 2s 2p

b. e. 1s22s12p7

1s 2s 2p

c. f. 1s22s22p63s23p63d84s2

1s 2s 2p

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EXERCISE 8.1Look at the following orbital diagrams and

electron configurations, which are possible and which are not according to the Pauli Exclusion Principle? Explain:

a. d. 1s22s22p4

1s 2s 2p

b. e. 1s22s42p2

1s 2s 2p

c. f. 1s22s22p63s23p103d10

1s 2s 2p

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The Building-Up PrincipleGround State: The electron configurationassociated with the lowest energy level of the

atom.Na 1s22s22p63s1

Excited State: The electron configuration associated

with an atom the energy levels other than the most

stable (ground state).

Na* 1s22s22p63p1 (emission of a yellow light at 589nm)

Energy s < p < d < f


Rules of Aufbau Principle: Lower n orbitals fill first.

Each orbital holds two electrons; each with different ms.

Half-fill degenerate (same energy level) orbitals before pairingelectrons. (p, d, & f)

NOT __ 3px 3py 3pz

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A mnemonic diagram of the Aufbau PrincipleIncreasing

Energy


Element Diagram ConfigurationLi (Z = 3) 1s2 2s1

1s 2s

Be (Z = 4) 1s2 2s2 1s 2s

B (Z = 5) __ __ 1s2 2s2 2p1

1s 2s 2px 2py 2pz

C (Z = 6) __ 1s2 2s2 2p2

1s 2s 2px 2py 2pz

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Element Diagram ConfigurationO (Z = 8) 1s2 2s2 2p4

1s 2s 2px 2py 2pz

Ne (Z = 10) 1s2 2s2 2p6

1s 2s 2px 2py 2pz

S (Z = 16) 1s 2s 2px 2py 2pz 3s 3px 3py 3pz

1s2 2s2 2p6 3s2 3p4 or [Ne] 3s2 3p4

abbreviations using the noble gases referred to as apseudo-noble gas core.Valence Electrons: an electron in an atom outside thenoble gas or pseudo-noble-gas core.

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Table of Electron Configuration using noble gas core


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Table of the Valence-shell configurations of the Elements


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The building-up order using the Periodic Table.


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EXAMPLE 8.2: Use the Aufbau Principle to obtain the complete electron configuration for the ground state of the Gallium atom (Z = 31). Abbreviate with the noble gas core and what is the valence shell configuration?

Gallium (Ga) Z = 31Full configuration: 1s22s22p63s23p64s23d104p1

Rearranged by shells: 1s22s22p63s23p63d104s24p1

Abbreviated configuration: [Ar] 3d104s24p1

Valence-shell configuration: 4s24p1


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EXERCISE 8.2: Use the Aufbau Principle to obtain the complete electron configuration for the ground state of the Manganese atom (Z = 25). Abbreviate with the noble gas core and what is the valence shell configuration?


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EXAMPLE 8.3: What are the configurations for the outer electrons of :a.Tellurium Z = 52

[Kr] 5s24d105p4

[Kr] 4d105s25p4

5s25p4

a. Nickel Z = 28[Ar]4s23d8

[Ar] 3d84s2

3d84s2


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EXERCISE 8.3: What are the configurations for the noble gas and the outer electrons of :

a.Arsenic

b.Bromine

c.Silver

d.Calcium


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EXERCISE 8.4: The lead atom has a ground state configuration of [Xe]4f145d106s26p2. find the period and group for this element. From it’s position in the periodic table, classify it as main-group element, a transition element or an inner transition element.

Anomalous Electron Configurations

19 of the predicted configurations from the periodic table are wrong Largely due to unusual stability of both half-

filled and fully filled subshellsCr (Z=24) expected configuration: 1s2 2s2 2p6 3s2 3p6 4s2

3d4

__4s 3d 3d 3d 3d 3d

actual configuration: 1s2 2s2 2p6 3s2 3p6 4s1 3d5

4s 3d 3d 3d 3d 3d

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Orbital Diagrams of Atoms; Hund’s Rule

Hund’s Rule: the lowest energy arrangement ofelectrons in a subshell is obtained by puttingelectrons into separate orbitals of the subshellwith the same spin BEFORE pairing the electrons.

* 1s 2s 2p

1s 2s 2p

1s 2s 2p20

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EXAMPLE 8.4: Write the orbital diagram for the ground state of the iron atom. Z = 26

Electron configuration: 1s22s22p63s23p63d64s2

Noble gas: [Ar] 3d64s2

Valence electron: 3d64s2

Orbital Diagram:

1s 2s 2p 3s 3p 4s 3d


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EXERCISE 8.5: Write the orbital diagram for the ground state of the phosphorus atom. Z = 15

Electron configuration:

Noble gas:

Valence electron:

Orbital Diagram:


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Paramagnetic Substance: a substance that is weakly attracted by a magnetic field this attraction is generally the result of unpaired electrons

Diamagnetic Substance: a substance tht is not attracted by a magnetic field or is very slightly repelled by such a field. This property generally means that the substance has only paired electrons

Magnetic Properties of Atoms

Periodic Properties

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The Periodic Law: When the elements are arranged by atomic number, their physical and chemical properties vary periodically.

•Atomic Radius

•Ionization Energy

•Electron Affinity

(important in discussions of chemical bonding)

Periodic Properties

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Representation of Atomic Radii of the Main-Group Elements

Periodic Properties

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Two Factors that primarily determine the size of the outermost orbital:

•Principle quantum number (n) of the orbital; the larger the n of the orbital, the larger the size of the orbital.

•The effective nuclear charge acting on an electron in the orbital; as the effective nuclear charge increases, the size of the orbital decreases by pulling the electrons inward.

•Effective nuclear charge: the positive charge that an electron experiences from the nucleus, equal to the nuclear charge but reduced by any shielding or screening from any intervening electron distribution.

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EXAMPLE 8.5: Refer to the periodic table use the trends noted for size of atomic radii to arrange the following in order of increasing atomic radius: Al, C, Si

C is above Si in Group IVA the radius of C is smaller than that of Si.Al and Si are in the same period, going to the right of the table the radius of Si is smaller than that of Al

C, Si, AlIn order of increasing radius

Periodic Properties

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EXERCISE 8.6: Using the periodic table, arrange the following in order of increasing atomic radius: Na, Be, Mg.

Periodic Properties

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Ionization Energy: the minimum energy needed to remove the highest-energy (the outermost) electron from the neutral atom in the gaseous state.

Li (1s22s1) Li+ (1s2) + e-

•Within a period, values tend to increase with atomic number the lowest values are found in Group 1A.•Elements with the lower ionization energy lose electrons easily•Noble gases have high ionization energy•Generally, as atomic numbers increase, ionization energy increases

Periodic Properties

Trends of First Ionization Energy, Ei

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Increase

Increase

Periodic Properties

Higher Ionization Energy, Ei1234…

Easy to remove an electron from a partially filled valence shell

Difficult to remove an electron from a filled valence shell

Large amount of stability associated with filled s & p subshells

Na: 1s2 2s2 2p6 3s1

Mg: 1s2 2s2 2p6 3s2

Cl: 1s2 2s2 2p6 3s2 3p5

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Ionization Energy, Ei

Some exceptions/irregularities to general trend Ei Be > Ei B we would expect opposite

Be 4 e 1s2 2s2

B 5 e 1s2 2s2 2p1

2s is closer to nucleus than 2p, Zeff for Be is stronger

2s is held more tightly and is harder to remove

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Periodic Properties

Ionization Energy, Ei Ei N > Ei O we would expect opposite

N 7e 1s2 2s2 2p3 __ __ __

O 8e 1s2 2s2 2p4 __ __ __

Only difference is that an electron is being removed from a half-filled orbital (N) and one from a filled orbital (O) Electrons repel each other and tend to stay as

far apart as possible, electrons that are forced together in a filled orbital are slightly higher in energy so it is easier to remove one O < N

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Periodic Properties

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EXAMPLE 8.6: Using the periodic table, arrange the following in order of increasing ionization energy: Ar, Se, S.

•Se is below S I Group VIA ionization energy of Se should be lower than S•S and Ar are in the same period with Z increasing from S to Ar the ionization energy of S should be lower than that of Ar.

• Se > S> Ar

Periodic Properties

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EXERCISE 8.7: The first ionization energy of the chlorine atom is 1251 kJ/mol. State which of the following values would be the more likely ionization energy for the iodine atom. Explain.

a. 1000kJ/mol or b. 1400kJ/mol

Periodic Properties

Ionic Radii or size

Atoms expand when converted to anions III A ns2 np1 __ __ __ IV A ns2 np2 __ __ __ V A ns2 np3 __ __ __ VI A ns2 np4 __ __ __ VII A ns2 np5 __ __ __

Adding one electron to each of these will notadd another shell it will just fill an alreadyoccupied p subshell

Therefore the expansion is due to the decrease in Zeff and the increase in the electron-electron repulsions

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Ionic Radii or size

Atoms contract when an electron is removed to form a cation.

Dec. # of shells

Inc. Zeff : Less electrons, less shielding, outer electrons more attracted to nucleus, therefore smaller more compact

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Higher Ionization Energy, Ei1234…

Ionization is not limited to one electronM + Energy M+ + e Ei1M+ + Energy M2+ + e Ei2M2+ + Energy M3+ + e Ei3

Larger amounts of energy are needed for each successive ionization, harder to remove an electron from a positively charged cation

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Electron Affinity, Eea Energy change that occurs when an electron

is added to an isolated atom in the gaseous state.

The more negative the Eea , the greater the tendency of the atom to accept an electron

Group 7A (halogens) have the most negative Eea, high Zeff and room in valence shell

Group 2A and 8A have near zero or slightly positive Eea

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Periodic Properties

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EXERCISE 8.8: Using the general comments that were discussed in this section, decide which has the larger negative electron affinity: C or F.

Periodic Properties

Alkali Metals Group 1A (ns1)

Metallic Soft Good Conductors Low melting point Lose 1 electron in redox reactions;

powerful reducing agent Very reactive Not found in elemental state in nature

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Periodicity in the Main-Group Elements

Alkaline Earth Metals Group 2A (ns2)

Harder, but still relatively soft Silvery High melting point than group 1A Less reactive than group 1A Loses 2e- in redox reaction; powerful

reducing agent Not found in elemental form in nature

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Group 3A (ns2np1) All but Boron which is a metalloid

Silvery Good conductor Relatively soft Less reactive than 1A & 2A metals

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Halogens Group 7A (ns2np5)

Non-metals Diatomic molecules Tend to gain e- during redox reaction.

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Noble Gases Group 8A (ns2np6)

Colorless, odorless, unreactive gases Stable because of the filled subshell

Makes it difficult to add electrons or remove electrons

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Example 1: Electron Config. And NG Abb.

1. Sodium

2. Titanium

3. Argon

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Example 2: Ionic Radii

Which of the following in each pair has a larger atomic radius?

1.Carbon or Fluorine2.Chlorine or Iodine3.Sodium or Magnesium4.O or O2-

5.Ca or Ca2+

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Example 3: Quantum Numbers and Electron Configuration

What are the 4 quantum numbers for the following? Remember you are only interested in the last electron!!

1. C

2. Na+

3. S

4. N3- 48

Example 4: Electron config. and NG Abb.

1. Cl-

2. F-

3. Ca2+

4. Na+

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CH 8: Electron Configuration & periodicity

Documents

electron configurations

valence shell configuration

valence electrons

gallium atom z

electrons inan atom

pzne z

1s2 2s1 orbital diagram

1s2 2s2 2p4