Chapter 8

Electronic Configuration and Periodicity

AnnouncementsEXAM 3Wednesday, Sept 306-7:30 PMChapter 6-840 multiple choice/2 problems

COMPREHENSIVE MAKE UP EXAMWednesday, Oct 76-7:30 PMChapter 1-1060 multiple choice

We will use recitation as lecture on Wednesday, so we don’t have to rush. 4:30-6PM this Thursday??

I. The Periodic Law and the Periodic Table

• 1864 Newland “Law of Octaves

• 1869 Dimitri Mendeleev and Lother Meyer

When the elements are arranged in order of increasing atomic mass, certain sets of chemical and physical properties recur periodically.

• 1913 Henry Mosely relates X-ray frequency to atomic number

When the Elements Were Discovered

Predicted Elements were Found

1. Principal Quantum Number (n): Defines the size and energy level of the orbital. n = {1,2,3,4,.....}. Also called a shell (K = 1, L = 2, M = 3, N = 4, .....).

2. Angular Momemtum Quantum Number (l): Defines the “shape” of the orbital which is a volume in space where the electron is likely to be found. Also called a subshell. l = {0,1,2,3...up to n-1} where (0=s, 1=p, 2=d, 3=f)

3. Magnetic Quantum Number (ml): Defines the spatial orientation of an orbital of the same energy. ml = {-l, 0, +l}

4. Magnetic Spin Quantum Number (ms): Defines the orientation of “electron spin”. ms = {+1/2 or -1/2}.

Quantum numbers (n,l,ml,ms) specify “allowed states” or “orbitals” which are regions of space where electrons are likely to be found around the nucleus.

Each electron can described by a unique set of 4 quantum numbers n, l, ml, ms. You can think of the quantum numbers as specifying the “house” or wavefunction or orbital where the electron is in space.

Name

principal

Symbol Permitted Values Property

n positive integers (1,2,3...) orbital energy (size)

angular momentum

lintegers from 0 to n-1

orbital shape (0, 1, 2, and 3 correspond to s, p, d, and f orbitals, respectively.)

magnetic ml integers from -l to 0 to +l orbital orientation in space

spin ms +1/2 or -1/2 direction of e- spin

angular momentum

magnetic

spin

magneticangular momentumprincipal

A shell describes the n quantum number (n = 1 = K; n = 2 = L and so on). A subshell refers to the specific n,l quantum number as in a 2s subshell or 2p subshell or 3s subshell or 3d subshell.

Each orbital can hold two electrons of opposite spin!

n value 1 2Shell K L Subshell 1s 2s 2p

3 Li 2 1 4 Be 2 2 5 B 2 2 1 6 C 2 2 2 7 N 2 2 3 8 O 2 2 4 9 F 2 2 5 10 Ne 2 2 6

s-orbitalp-orbital d-orbital

f-orbital

The shape of an orbital is given by the “l” quantum number (l = {0,1,2 up to n-1}. The number of orbitals and its orientation in space is given by the angular momentum quantum number ml {-l,..0..+l}.

The shapes are radial boundry plots that incorporate 95% probability of finding an electron.

The “Shapes” Are Boundry Surface Plots

Boundary Surface Plot

(“orbital shape”)

Probability atpoint r in space

Probability of pointsadded up in a circularstrips r + rdr

1s 2s 3s

Radial Probability Distribution Plots

Probability Density Plotsnodes (0 probability)

The Shapes of Orbitals

1s 2s 3s

1s-Orbital Shape is spherical and gets larger with the n-quantum number.

Notice the nodes or regions of 0 electron density in the plots.

Boundary Surface Plot having 95% of all probability

2-D Slice of probability atpoints in space

1s

1s

2s

2s

3s

1s 2s 3s

s-orbitalp-orbital d-orbital

f-orbital

Your job is to know the shapes of the orbitals. We will use them to form “chemical bonds” in molecules by overlapping them in space.

The shapes are radial boundry plots that incorporate 95% probability of finding an electron.

The p Orbitals ( l = 1)

2-D Slice of probability atpoints in space

Boundary Surface Plot having 95% of all probability

Radial ProbabilityDistribution Plotfor p-orbital.

The 3 p orbital Boundary Surface Plots (Shapes)

Three p-orbitals superimposed

2 px 2 py 2 pz

The 5-d Boundry Surface Plots (Shapes)

Five d-orbitals superimposed

dyz dxydxz

dz2

dx2 - dy2

Ener

gy

Electrons willfill lowest energy orbitals first!

We can build up the electron configuration of all the ground state elements (and their ions) by filling in the orbitals with electrons.

Hydrogen Atom Multi-electron atoms

Orbitals are “degenerate” or the same energy in Hydrogen!

Quiz 5 or 6

Blue light scattered by air molecules in the sky has a frequency of about 7.5 X 1014 Hertz.

1) Calculate the wavelength, in nm, associated with this radiation knowing that the speed of light is 3.00 x 108 m/s.

2) Calculate the energy of a single photon of this light

The lowest energy (ground state) electronic configuration of all elements are constructed by filling lowest energy orbitals sequentially in what is called the “Aufbau Process”.

1. Lower energy (n-quantum number) orbitals fill first.

2. Hund’s Rule-degenerate orbitals fill one at a time before electrons are paired.

3. Pauli Exclusion Principle: No two electrons can have same 4-quantum numbers)

Electrons fill the lowest energy orbitals first, 2 at a time!

The Pauli Exclusion principle states: “No two electrons can have the same 4-quantum numbers”. Using this fact, and the Aufbau process, we can build the electronic configuration of the elements.

Li 1s22s1

Example:

each electron can be described by a unique set of quantum set of four quantum numbers.

Atomic Number/Element

Orbital BoxDiagram

Full-electronicconfiguration

Condensed-electronicconfiguration

[He]2s1

Chemists use spdf notation and orbital box diagrams to denote or show the “ground state electronic configuration” of elements.

spdf Notation

orbital box diagram

H

He

1s1

1s2

Element

electron shellprincipal quantum #

orbital typeangular quantum #

# of electronsin orbital

An arrow denotes an electron with “spin up” or “spin-down”.

Remember, no two electrons can have the same 4 quantum numbers!

The order of filling of the orbitals can be remembered using a mnemonic device. Memorize it!

Building electronic configuration using Aufbau and Hund

H

He

Li

Be

1s1

1s2

1s22s2

1s22s1

Atomic Number/Element

Orbital BoxDiagram

Full-electronicconfiguration

Condensed-electronicconfiguration

1s1

1s2

[He]2s1

[He]2s2

written with noble gas configuration

1s22s22p3

1s22s22p4

1s22s22p5

1s22s22p6

1s22s22p1

1s22s22p2

B

C

Atomic Number/Element

Orbital BoxDiagram

Full-electronicconfiguration

Condensed-electronicconfiguration

[He]2s22p1

[He]2s22p2

[He]2s22p3

[He]2s22p4

[He]2s22p5

[He]2s22p6

Anamolies occur when filling the d-orbitals

Odd-filling behavior here!4th and 9th position.

Paramagneticunpaired electrons

2p

Diamagneticall electrons paired

2p

• Diamagnetic atoms or ions:– All e- are paired.– Weakly repelled in a magnetic field.

• Paramagnetic atoms or ions:– Unpaired e- exist in an orbital– Attracted to an external magnetic field.

Unpaired electrons in orbitals gives rise to paramagnetism and is attracted to a magnetic field. Diamagnetic species contain all paired electrons and is “repelled” by the magnetic field.

Paramagnetic Diamagnetic

Unpaired electrons in orbitals gives rise to paramagnetism and is attracted to a magnetic field. Diamagnetic species contain all paired electrons and is “repelled” by the magnetic field.

Magnetic field off Magnetic field on Magnetic field on

Paramagentic Diamagentic

Write a set of quantum numbers for the third electron and a set for the eighth electron of the F atom.

What is the electron configuration of Mg?Mg 12 electrons

1s22s22p63s22 + 2 + 6 + 2 = 12 electrons

Abbreviated as [Ne]3s2 [Ne] 1s22s22p6

What are the possible quantum numbers for the last (outermost) electron in Cl?

Cl 17 electrons 1s < 2s < 2p < 3s < 3p < 4s 1s22s22p63s23p5 2 + 2 + 6 + 2 + 5 = 17 electronsLast electron added to 3p orbitaln = 3 l = 1 ml = -1, 0, or +1 ms = ½ or -½

What neutral atom is isoelectronic with H- ?H-: 1s2 same electron configuration as He

Use the orbital diagram to find the third and eighth electrons.

Write the electronic configuration of F using orbital box notation. What is the set of quantum numbers for the third electron and a set for the eighth electron of the F atom.

9F1s 2s 2p

The third electron is in the 2s orbital. Its quantum numbers are n = 2 l = 0, ml = 0 and ms = 1/2

The eighth electron is in a 2p orbital. Its quantum numbers are

n = 2 l = 1, ml = -1,+1 or 0 and ms = 1/2

Using the periodic table on the inside cover of the text and give the full and condensed electrons configurations, partial orbital diagrams showing valence electrons, and number of inner electrons for the following elements:

(a) potassium (K: Z = 19)

(b) molybdenum (Mo: Z = 42)

(c) lead (Pb: Z = 82)

(b) for Mo (Z = 42) 36 inner electrons and 6 valence electrons1s22s22p63s23p64s23d104p65s14d5

[Kr] 5s14d5

(c) for Pb (Z = 82) 78 inner electrons and 4 valence electrons.

[Xe] 6s24f145d106p2

condensedpartial orbital diagram

full configuration

condensed

partial orbital diagram

full configuration 1s22s22p63s23p64s23d104p65s24d10

5p66s24f145d106p2

6s2 6p2

5s1 4d5 5p

(a) for K (Z = 19)1s22s22p63s23p64s1

[Ar] 4s1

4s1

condensedorbital diagram

full configurationThere are 18 inner electrons.

3d 4p

SOLUTION:

paramagnetic(a) Mn2+(Z = 25) Mn([Ar]4s23d5) Mn2+ ([Ar] 3d5) + 2e-

(b) Cr3+(Z = 24) Cr([Ar]4s23d6) Cr3+ ([Ar] 3d5) + 3e-

paramagnetic

(c) Hg2+(Z = 80) Hg([Xe]6s24f145d10) Hg2+ ([Xe] 4f145d10) + 2e-

not paramagnetic (is diamagnetic)

Use condensed electron configurations to write the reaction for the formation of each transition metal ion, and predict whether the ion is paramagnetic.

(a) Mn2+(Z = 25) (b) Cr3+(Z = 24)

Write the electron configuration and remove electrons starting with ns to match the charge on the ion. If the remaining configuration has unpaired electrons, it is paramagnetic.

(c) Hg2+(Z = 80)

Metals lose electrons so that cation has a noble-gas outer electron configuration.

Na [Ne]3s1 Na+ [Ne]

Ca [Ar]4s2 Ca2+ [Ar]

Al [Ne]3s23p1 Al3+ [Ne]

Non-metals gain electrons so that anion has a noble-gas outer electron configuration.

H 1s1 H- 1s2 or [He]

F 1s22s22p5 F- 1s22s22p6 or [Ne]

O 1s22s22p4 O2- 1s22s22p6 or [Ne]

N 1s22s22p3 N3- 1s22s22p6 or [Ne]

Metals loose electrons (oxidized) to become cations. Non-metals gain electrons to become anions. The electronic configuration of each reflects this change in the number of electrons.

Metals and non-metal ions tend to form electronic states closest to their nearest noble gas configuration.

Metals and non-metals form ions with electronic configurations closest to their nearest noble gas configuration.

Fe: [Ar]4s23d6 Fe2+: [Ar]4s03d6 or [Ar]3d6

Mn: [Ar]4s23d5 Mn2+: [Ar]4s03d5 or [Ar]3d5

Fe3+: [Ar]4s03d5 or [Ar]3d5Fe: [Ar]4s23d6

When a cation is formed from an atom of a transition metal, electrons are removed first from the ns orbital, then from the (n-1)d orbital.

Na+, Al3+, F-, O2-, and N3- are all said to be “isoelectronic with Ne” as they have the same electronic configuration....all subshells are filled.

Isoelectronic species are two different elements that have the same electronic configuration (but not the same nuclear configuration).

Na: [1s22s22p63s1] =====> Na+: [1s22s22p6] = [Ne]oxidation

oxidationAl: [1s22s22p63s23p1] =====> Al3+: [1s22s22p6] = [Ne]

N: [1s22s22p3] =====> N3-: [1s22s22p6] = [Ne]reduced

O: [1s22s22p4] =====> O2-: [1s22s22p6] = [Ne]reduced

F: [1s22s22p5] =====> F-: [1s22s22p6] = [Ne]reduced