Unit 6 chapters 11 12 pages 295 366 atomic electron configurations and periodicity
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Unit 6: Chapters 11-12. Pages 295-366 ATOMIC ELECTRON CONFIGURATIONS AND PERIODICITY PowerPoint PPT Presentation

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Unit 6: Chapters 11-12. Pages 295-366 ATOMIC ELECTRON CONFIGURATIONS AND PERIODICITY. Bohr Model. First model of the electron behavior Vital to understanding the atom Does not work for atoms with more than 1 electron. Collision of Ideas. Matter. Dalton. Thompson. Rutherford. Bohr. ?.

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Bohr Model

  • First model of the electron behavior

  • Vital to understanding the atom

  • Does not work for atoms with

    more than 1 electron

Collision of Ideas







De Broglie






The Photoelectric Effect

Duality of Light

  • Wave behavior

  • Particle behavior



de Broglie’s Novel Notion

Light was “known” (thought) to be a wave, but

Einstein showed that it also acts particle-like

Electrons were particles with known mass & charge

What if ……

electrons behaved as waves also

Evidence for de Broglie’s Notion

Diffraction pattern obtained with firing a beam of electrons through a crystal.

This can only be explained if the electron behaves as a wave!

Nobel Prize for de Broglie in 1929

Electron Characteristics

  • Extremely small mass

  • Located outside the nucleus

  • Moving at very high speeds

  • Have specific energy levels

  • Standing wave behavior

Baseball vs Electron

A baseball behaves as a particle and follows a predictable path.


An electron behaves as a wave, and its path cannot be predicted.

All we can do is to calculate the probability of the electron following a specific path.

What if a baseball behaved like an electron?

  • Characteristic wavelength

  • baseball 10-34 m

  • electron 0.1 nm

All we can predict is…..

Werner Heisenberg(1901-1976)

The Uncertainty Principle

  • Proposed that the dual nature of the electron places limitation on how precisely we can know both the exact location and speed of the electron

  • Instead, we can only describe electron behavior in terms of probability.



Erwin Schrodinger(1887-1961)

Wave Equation & Wave Mechanics

  • In 1926, Austrian physicist, proposed an equation that incorporates both the wave and particle behavior of the electron

  • When applied to hydrogen’s 1 electron atom, solutions provide the most probable location of finding the electron in the first energy level

  • Can be applied to more complex atoms too!

Solutions to Schrodinger’s Wave Equation

Gives the most probable location of electron in 3-D space around nucleus (probability map)

- most probable

location called an


- orbitals can hold a

maximum of 2 e-

“Most Successful Theory of 20th Century”








De Broglie










Quantum Mechanics ModelDescribes the arrangement and space occupied by electrons in atoms

Electron’s energy is quantized



Mathematics of waves to define orbitals

(wave mechanics)

Bohr Model v. Quantum Mechanics

BohrQ. Mech.




Dartboard Analogy

Suppose the size of the probability distribution is defined

as where there is a % chance of all hits being confined.

Quantum Mechanics Model

The electron's movement cannot be known precisely.

We can only map the probability of finding the electron at various locations outside the nucleus.

The probability map is called an orbital.

The orbital is calculated to confine 90% of electron’s range.

Arrangement of Electrons in Atoms

Electrons in atoms are arranged as

SHELLS (n) = distance from nucleus

1, 2, 3, …

SUBSHELLS (l) = shape of region of probability

s, p, d, f

ORBITALS (ml) = orientation in space

Arrangement of Electrons in Atoms

  • There is a relationship between the quantum number (n) and its the number of subshells.

Principal quantum number (n) = number of subshells

Representing s Orbitals

Comparison of 1s and 2s Orbitals

The 2s orbital is similar to the 1s orbital, but larger in size.

”Larger” means that the highest probability for finding the electron lies farther out from the nucleus.

Each can hold a maximum of


Probability Maps of the Three 2p Orbitals

The 2p orbital is in the n = energy level.

There are 2p orbitals oriented in three directions.

Each orbital can hold a maximum of electrons.

The maximum number of electrons in the 2p sublevel is .

Adding all 2p orbitals would result in a sphere.

Probability Maps of the Five 3d Orbitals

The five 3d orbitals are generally oriented in different directions.

Adding all five orbitals, would result in a sphere.

The five orbitals, taken together, make up the d subshell of the n = 3 shell.

Each orbital can hold a maximum of two electrons.

This sublevel has a maximum of electrons.

Probability Maps of 7 f Orbitals

Arrangement of Electrons in AtomsElectron Spin Quantum Number- ms

Each orbital can be assigned no more than 2 electrons! And each electron spins in opposite directions.

Electron Spin Quantum Number

Diamagnetic: NOT attracted to a magnetic field

Paramagnetic: substance is attracted to a magnetic field. Substance has unpaired electrons.




n ---> shell1, 2, 3, 4, ...

l ---> sublevels, p, d, f

ml ---> orbital -l ... 0 ... +l

ms ---> electron spin+1/2 and -1/2

Pauli Exclusion Principle- No two electrons in the same atom can have the same set of 4 quantum numbers.

Determine the quantum numbers for the outer two valence electrons in the lithium atom.

Aufbau Principle-Electrons fill open lower energy levels sequentially lower energy to higher energy

spdf notation

for H, atomic number = 1


no. of




value of l

value of n

Writing Electron Configurations

Two ways of writing configs. One is called thespdf notation.

Broad Periodic Table Classifications

  • Representative Elements(main group): filling s and p orbitals (Na, Al, Ne, O)

  • Transition Elements: filling dorbitals (Fe, Co, Ni)

  • Lanthanide and Actinide Series(inner transition elements): filling 4fand 5forbitals (Eu, Am, Es)

Writing Orbital Notations

Two ways of writing configs. Other is called theorbital box notation.

One electron has n = 1, l = 0, ml = 0, ms = + 1/2

Other electron has n = 1, l = 0, ml = 0, ms = - 1/2

Energy ordering of orbitals for multi-electron atoms

Different subshells within the same principal shell have different energies.

The more complex the subshell, the higher its energy. This explains why the 3d subshell is higher in energy than the 4s subshell.

Rules for Filling Orbitals

Bottom-up (Aufbau’s principle)

Fill orbitals singly before doubling up (Hund’s Rule)

Paired electrons have opposite spin (Pauli exclusion principle)



Atomic Number

Full Configuration

Valence Configuration

Shorthand Configuration

Orbital diagram and electron configuration for a ground state lithium atom

Orbital diagram and electron configuration for a ground state carbon atom

Hund’s Rule- electrons in the same sublevel will spread out into their own orbital before doubling up.

Silicon's valence electrons

Selenium's valence electrons

Core electrons and valence electrons in germanium

Outer electron configuration for the elements

The periodic table gives the electron configuration for As

Valence Electrons by Group

Ion charges by group

Periodic Law

All the elements in a group have the same electron configuration in their outermost shells

Example: Group 2

Be2, 2

Mg 2, 8, 2

Ca 2, 2, 8, 2


Specify if each pair has chemical properties that are similar (1) or not similar (2):

A. Cl and Br

B. P and S

C. O and S

Higher effective nuclear charge

Electrons held more tightly

Larger orbitals.

Electrons held less


General Periodic Trends

1. Atomic and ionic size2. Electron affinity

3. Ionization energy 4. Metallic Character

Effective Nuclear Charge, Z*

  • Z* is the nuclear charge experienced by the outermost electrons. Screen 8.6.

  • Explains why E(2s) < E(2p)

  • Z* increases across a period owing to incomplete shielding by inner electrons.

  • Estimate Z* by --> [ Z - (no. inner electrons) ]

  • Z = number of electrons

  • Charge felt by 2s e- in Li Z* = 3 - 2 = 1

  • Be Z* = 4 - 2 = 2

  • B Z* = 5 - 2 = 3and so on!

Effective Nuclear Charge

Figure 8.6

Electron cloud for 1s electrons

Effective Nuclear Charge, Z*

  • AtomZ* Experienced by Electrons in Valence Orbitals

  • Li+1.28

  • Be-------

  • B+2.58

  • C+3.22

  • N+3.85

  • O+4.49

  • F+5.13

Increase in Z* across a period




Atomic Size

  • Size goes UPon going down a group. See Figure 8.9.

  • Because electrons are added further from the nucleus, there is less attraction.

  • Size goes DOWNon going across a period.

Atomic Radii

Figure 8.9

Trends in Atomic SizeSee Figures 8.9 & 8.10

Ion Sizes

Does the size go

up or down when losing an electron to form a cation?




, 78 pm

2e and 3 p

Ion Sizes

Forming a cation.

  • CATIONS are SMALLER than the atoms from which they come.

Li,152 pm

3e and 3p

Ion Sizes

Does the size go up or down when gaining an electron to form an anion?



F, 71 pm


, 133 pm

9e and 9p

10 e and 9 p

Ion Sizes

Forming an anion.

  • ANIONS are LARGER than the atoms from which they come.

Trends in Ion Sizes

Figure 8.13

Ionization EnergySee Screen 8.12

IE = energy required to remove an electron from an atom in the gas phase.

Mg (g) + 738 kJ ---> Mg+ (g) + e-

Ionization EnergySee Screen 8.12

Mg (g) + 735 kJ ---> Mg+ (g) + e-

Mg+ (g) + 1451 kJ ---> Mg2+ (g) + e-

Mg2+ (g) + 7733 kJ ---> Mg3+ (g) + e-

Energy cost is very high to dip into a shell of lower n. This is why ox. no. = Group no.

Trends in Ionization Energy

Trends in Ionization Energy

  • IE increases across a period because Z* increases.

  • Metals lose electrons more easily than nonmetals.

  • Metals are good reducing agents.

  • Nonmetals lose electrons with difficulty.

Trends in Ionization Energy

  • IE decreases down a group

  • Because size increases.

  • Reducing ability generally increases down the periodic table.

  • See reactions of Li, Na, K


  • A measure of the ability of an atom that is bonded to another atom to attract electrons to itself.

Electron Affinity

A few elements GAIN electrons to form anions.

Electron affinity is the energy involved when an atom gains an electron to form an anion.

A(g) + e- ---> A-(g)

E.A. = ∆E





+ electron


O atom

Electron Affinity of Oxygen

∆E is EXOthermic because O has an affinity for an e-.

EA = - 141 kJ

Trends in Electron Affinity

  • See Figure 8.12 and Appendix F

  • Affinity for electron increases across a period (EA becomes more negative).

  • Affinity decreases down a group (EA becomes less negative).

Atom EA

F-328 kJ

Cl-349 kJ

Br-325 kJ

I-295 kJ

Trends in Electron Affinity

Metallic character trends in the periodic table

Metallic Character

The text links metallic character to the tendency to lose electrons in chemical reactions, and nonmetallic character to the tendency to gain electrons in chemical reactions. The metallic character trends therefore follow the ionization energy trends

The metallic character trends explain the location of metals, metalloids, and nonmetals

Which is the more metallic element, Sn or Te?

Which is the more metallic element, Si or Sn?

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