Chapter 7; Electronic Structure of Atoms

I. Electromagnetic Radiation

II. Flame Test/ Emission Spectra

III. Quantized Energy Levels

IV. Bohr Model/ Rydberg Equation

V. Principal Energy Levels, na) First Ionization Energy

b) 2nd , 3rd, 4th, etc Ionization Energy

Chapter 7; Electronic Structureof Atoms

VI. Sublevels (s, p, d, f)a) Photoelectron Spectroscopy

VII. Electron ConfigurationVIII. Valence Electrons/ CoreIX. Good/ Bad Point of Atom ModelX. Quantum Theory

a) Dual Nature of the Electronb) Heisenberg Uncertainty Principle

Chapter 7; ElectronicStructure of Atoms

XI. Quantum Numbers (n, l, ml, ms)

XII. Oribtal Diagramsa) Paramagnetism and Diamagnetism

Electronic Structure Model

Experimental Evidence1. Line Spectra

2. Ionization Energies

3. Photoelectron Spectrum

4. Intensity/detail of Line Spectra

What it means1. Electrons in quanitized ‘n’

2. # electrons in each ‘n’

3. # electrons in each ‘n’ and each sublevel

4. Indicates ‘n’ have sublevels associated with them

Electronic Structure

n # of Sublevel

# e- in n

(2n2)

Sublevel

Names

# e- in each sublevel

1 1 2 s s-2

2 2 8 s,p s-2, p-6

3 3 18 s,p,d s-2, p-6,

d-10

4 4 32 s,p,d,f s-2, p-6,

d-10, f-14

Order of orbitals (filling) in multi-electron atom

1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s7.7

What is the electron configuration of Mg?

Mg 12 electrons

1s < 2s < 2p < 3s < 3p < 4s

1s22s22p63s2 2 + 2 + 6 + 2 = 12 electrons

7.7

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

What is the electron configuration of Cl?

Cl 17 electrons 1s < 2s < 2p < 3s < 3p < 4s

1s22s22p63s23p5 2 + 2 + 6 + 2 + 5 = 17 electrons

Electron Configurations of Cations and Anions

Na [Ne]3s1 Na+ [Ne]

Ca [Ar]4s2 Ca2+ [Ar]

Al [Ne]3s23p1 Al3+ [Ne]

Atoms lose electrons so that cation 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]

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

Of Representative Elements

8.2

Na+: [Ne] Al3+: [Ne] F-: 1s22s22p6 or [Ne]

O2-: 1s22s22p6 or [Ne] N3-: 1s22s22p6 or [Ne]

Na+, Al3+, F-, O2-, and N3- are all isoelectronic with Ne

What neutral atom is isoelectronic with H- ?

H-: 1s2 same electron configuration as He

8.2

Electron Configurations of Transition Metals

• Completely filled or half-completely filled d-orbitals have a special stability

– Some “irregularities” are seen in the electron configurations of transition and inner-transition metals.

Electron Configurations of Cations of Transition Metals

8.2

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

Fe: [Ar]4s23d6

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

Fe3+: [Ar]4s03d5 or [Ar]3d5

Mn: [Ar]4s23d5

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

Order of filling; 3s<3p<4s<3dBut when removing electrons to form + ions for transition metalsOrder of removing electrons; 4s<3d<3p<3s

Electronic Structure

Good Points• Electrons in Quantized

Energy Levels• Maximum # electrons

in each n is 2n2

• Sublevels (s,p,d,f) and # electrons they hold

Bad Points• Electrons are placed in

orbits about nucleus• Only explains

emission spectra of H2

• Does not address all interactions

• Treats electron as particle

H

+1

Be

+4

There are less interactions to take into account in H than other elements

Interactions1. Attraction between + nucleus and negative electrons

Interactions1. Attraction between + nucleusand negative electrons2. Repulsion between electronsin same energy level.3. Shielding effect of filledprincipal energy levels.

Quantum Theory – Revised Electronic Structure Model

1. Dual Nature of the Electron

2. Heisenberg Uncertainty Principle

Dual Nature of Electron

Previous Concept;

A Substance is Either Matter or Energy

• Matter; Definite Mass and Position

Made of Particles

• Energy; Massless and Delocalized

Position not Specificed

Wave-like

Dual Nature of Electron

• Electron is both “particle-like” and “wave-like” at the same time.

• Previous model only considered “particle-like” nature of the electron

Heisenberg Uncertainty Principle

• Act of measuring the position and energy of electron changes the position of electron

– Better one variable is known (energy); the less well the other variable is known (position)

Orbitals Replace Orbits

• Orbits- Both electron position and energy known with certainty

• Orbitals – Regions of space where an electrons of a given energy will most likely be found

Quantum TheoryOrbitals Replace Orbits

Orbits Orbitals

Schrodinger Wave Equation ()

Describes size/shape/orientation of orbitals

7.5

• Wave Equation is based on…

1. Dual Nature of Electron (Electron both particle and wave-like at the same time.)

2. Heisenberg Uncertainty Principle(Orbitals describe a region in space an electron will most likely be.)

Wave Equation ()

• Wave Equation describe the size, shape, and orientation of the orbital the electron (of a given energy) is in. There are four variables in the function

-n; Energy and size of orbital– l; Shape of orbital

– ml; Orientation of orbital

– ms; Electron Spin

(n, l, ml, ms)

1. Each electron has a unique set of 4 Quantum Numbers

2. Each orbital described by the Quantum Numbers can hold a maximum of 2 electrons.

Schrodinger Wave Equation;1st Quantum Number

fn(n, l, ml, ms)

principal quantum number n

n = 1, 2, 3, 4, ….

n=1 n=2 n=3

7.6

distance of e- from the nucleus

= fn(n, l, ml, ms)

angular momentum quantum number l

for a given value of n, l = 0, 1, 2, 3, … n-1

n = 1, l = 0n = 2, l = 0 or 1

n = 3, l = 0, 1, or 2

Shape of the “volume” of space that the e- occupies

l = 0 s orbitall = 1 p orbitall = 2 d orbitall = 3 f orbital

Schrodinger Wave Equation2nd Quantum Number

7.6

Principal Energy Level, n

Sublevel,

l

Quantum # Electron Configuration

1 0 (1,0, , ) 1s

2 0 (2,0, , ) 2s

1 (2, 1, , ) 2p

3 0 (3,0, , ) 3s

1 (3, 1, , ) 3p

2 (3,2, , ) 3d

4 0 (4,0, , ) 4s

1 (4, 1, , ) 4p

2 (4, 2, , ) 4d

3 (4, 3, , ) 4f

l = 0 (s orbitals)

l = 1 (p orbitals)

7.6

l = 2 (d orbitals)

f-orbitals

Orbital Shapes

Orbital Type Shape Name

s Spherical

p Dumbbell

d Complex

f More complex

= fn(n, l, ml, ms)

magnetic quantum number ml

for a given value of lml = -l, …., 0, …. +l

orientation of the orbital in space

if l = 1 (p orbital), ml = -1, 0, or 1if l = 2 (d orbital), ml = -2, -1, 0, 1, or 2

Schrodinger Wave Equation3rd Quantum Number

7.6

Number of Degenerate Orbitals Needed for Each Type of Orbital (Sublevel)

Type of Orbital Maximum # of electrons in

Orbital

# of Degenerate Orbitals

s 2 1

p 6 3

d 10 5

f 14 7

ml = -1 ml = 0 ml = 1

ml = -2 ml = -1 ml = 0 ml = 1 ml = 2

= fn(n, l, ml, ms)

spin quantum number ms

ms = +½ or -½

Schrodinger Wave Equation4th Quantum Number

ms = -½ms = +½

7.6

Valid Possibilities for Quantum Numbers

Chemistry; The Science in Context; by Thomas R Gilbert, Rein V Kriss, and Geoffrey Davies, Norton Publisher, 2004, p125

How many 2p orbitals are there in an atom?

2p

n=2

l = 1

If l = 1, then ml = -1, 0, or +1

3 orbitals

How many electrons can be placed in the 3d subshell?

3d

n=3

l = 2

If l = 2, then ml = -2, -1, 0, +1, or +2

5 orbitals which can hold a total of 10 e-

7.6

Three Manners to Convey How Electrons are Arranged

1. Electron Configuration ; List Orbitals and Number of Electrons in Each

(1s22s22p63s2…)

2. Quantum Numbers (2,0,0,+1/2)

3. Orbital Diagrams; List Orbitals and show location of electrons and their spin

1s 2s 2p

Pauli exclusion principle - no two electrons in an atomcan have the same four quantum numbers.

The most stable arrangement of electrons in subshells is the one with the greatest number of parallel spins (Hund’s rule)or maximum # of unpaired electrons.

Orbital Diagrams

7.7

Orbital Diagrams

1s 2s 2p

Carbon; 6 electrons

Electron Configuration; 1s22s22p2

Orbital Diagram

Orbital Diagrams

1s 2s 2p

Oxygen; 8 electrons

Electron Configuration; 1s22s22p4

Orbital Diagram

Paramagnetic

unpaired electrons

2p

Diamagnetic

all electrons paired

2p7.8