Chemistry Lesson 18 Quantum Numbers and Electron Configurations

7/28/2019 Chemistry Lesson 18 Quantum Numbers and Electron Configurations

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Lesson 18 Quantum Numbers

and Electron Configurations

Objectives:

- 1. The student will define and explain the four

quantum numbers.

- 2. The student will explain and apply Hunds Rule

and the Pauli Exclusion Principle.

- 3. The student will write electron configurations for

elements, as well as determine what element is

represented by a specific electron configuration.


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I. Scientists whose theories led to

the understanding of the electron:a. Louis deBroglie: French graduate student in physics

who proposed The DeBroglie Hypothes, which statesthat particles have properties of waves as well as

properties of particles, the wave particle duality ofnature Formula ( = h/mv ) =wavelength, h=Plancks constant, m= mass in kg, v= velocity in m/s

b. Werner Heisenberg: German physicist who publishedthe Heisenberg Uncertainty Principle: it is impossible toknow the exact location and exact momentum of a

particle at the same time.c. Erwin Schrodinger: 1926, Austrian physicist who

treated electrons as waves to help determine probabilityof location within an atom. This led to the creation ofthe quantum mechanical model that we use to explain

the structure of the atom today.


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II. Labeling Electrons in atoms

a. Quantum numbers are used to differentiatebetween electrons

i. In quantum theory, each electron in an atom

is assigned a set of four quantum numbers.

ii. Three of these give the location of the

electron, and the fourth gives the orientationof the electron within the orbital

iii. Definitions of numbers

` 1. Principle Quantum NumbernThis numberdescribes the energy level that the electron occupies. It

can have a value of 1-7This defines the level of the

electron.


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2. Orbital Quantum Numberl(Azimuthal) this

number describes the shape of the orbital that the electron is

found in. It can have a value from 0-3. This defines the

sublevel of the electron. Also, the numbers can be replacedby letters according to the following:

a. 0 = s

b. 1 = pc. 2 = d

d. 3 = f

f orbital shapes:


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3. Magnetic Quantum Number - mlthis number

describes the orientation of the electrons in the orbitals. This

defines the orbital of the electron. There are 2l+1 orbitals ineach sublevel. This quantum number can have the following

values: (-l to +l)

a. Ifl

= 0, ml can equal 0b. Ifl = 1, ml

can equal1, 0, +1

c. Ifl = 2, ml can equal2, -1, 0, +1, +2

d. Ifl = 3, ml can equal3, -2, -1, 0, +1, +2, +3

4. Spin Quantum Numbermsthis number describes thedirection of spin of the electron in the orbitalelectrons in the

same level and sublevel must spin in opposite directions. This

can have a value or +1/2 or1/2 only.


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iv. According to the Pauli Exclusion Principle, no two

electrons can have the same four quantum numbers in the sameatom.

v. Think of these as City, Street, House Number, and

upstairs/downstairs apartment. No two people could have the

same complete address, but they could live in the same city, onthe same street, or even in the same house, but not the same

apartment.

b O bi l di d l fi i d l


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b. Orbital diagrams and electron configurations are models

for electron arrangements.

i. Orbital diagrams are used to show how electrons

are distributed among the different sublevels and also to

show the direction of spin.

ii. For orbital diagrams, you must fill in orbitals in the

same energy level with one electron each before pairing

up any electrons. This is known asHunds Rule

.

iii. Electron configurations are used to show similar

information, but are a much more abbreviated form.

iv. How many electrons can go in any level?(Maximum)

1. s = 2

2. p = 6

3. d = 104. f = 14


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v. What order do I fill the levels in? The Aufbau Principlestates that when predicting an atoms ground state electron

configuration, electrons will occupy the lowest energy

orbital available first.

1s

2s 2p

3s 3p 3d

4s 4p 4d 4f

5s 5p 5d 5f

6s 6p 6d

7s 7p

vii. This also could be written: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s,4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p


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III. Writing Lewis Structures or Lewis DotDiagrams for elements

a. This is a kind of short hand thatillustrates how many outer shell electronsan atom contains.

b. The purpose behind all of theconfigurations is because the number ofelectrons and their placement in theatom, strongly influences how the atom

will react, bond and the properties it willdemonstrate.


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c. Rules for writing dot diagrams:

i. Write configuration.

ii. How many e- are in the outer energy level?iii. Write the elements symbol.

iv. Draw dots around the symbol to represent

outer level electrons, each of the 4 sides

represents an orbital.

v. s electrons must be paired (1st two e-)

vi. Other three sides cannot be paired until each

has at least one e-. (Hunds Rule)


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d. Example:


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IV. Exceptions to electron configurationusing the Aufbau Diagram

a. A half full level is the next stable thing to afull level.

b. Some atoms will violate our predictions inorder to achieve stability. This can occur in

the transition metals when the predictedconfiguration ends in a d4 or d9.

c. It will steal a single electron from the full sshell that came before it to obtain 2 half full

shells or one half and one full shell.d. (s2 d4) becomes (s1 d5) and (s2 d9)

becomes (s1 d10)


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e. Actual exceptions:

* 5d1 fills before starting the 4f sequence

* 6d1 fills before starting the 5f sequence

Predicted configurations Actual configurations

Cr:[Ar]4s2, 3d4 Cr:[Ar] 4s1, 3d5

Cu:[Ar] 4s2, 3d9 Cu[Ar] 4s1, 3d10

Nb:[Kr]5s2,4d3 Nb:[Kr] 5s1, 4d4

Mo:[Kr] 5s2, 4d4 Mo[Kr] 5s1, 4d5Tc:[Kr] 5s2, 4d5 Tc[Kr] 5s1, 4d6

Ru[Kr] 5s2, 4d6 Ru[Kr] 5s1, 4d7

Rh[Kr] 5s2, 4d7 Rh[Kr] 5s1, 4d8

Pd[Kr] 5s2, 4d8 Pd[Kr] 5s0, 4d10

Ag[Kr] 5s2, 4d9 Ag[Kr] 5s1, 4d10

Pt[Xe] 6s2, 4f14, 5d8 Pt[Xe]6s1, 4f14, 5d9Au[Xe] 6s2, 4f14, 5d9 Au[Xe] 6s1, 4f14, 5d10


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Questions:

1. Make a chart, with the following columns:

Quantum number name, symbol, possible

values. Fill in the information for each of

the four quantum numbers.

2. What is the reason that an element cannot

have all four quantum numbers the same?

3. What is the rule which means spread

them out before you pair them up?


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Chemistry Lesson 18 Quantum Numbers and Electron Configurations

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