Atomic Orbitals & Electron Configurations

Chemistry

1

Atomic Orbitals & Electron Configurations

1. Identify the relationships among a hydrogen atom’s energy levels, sublevels, and atomic orbitals.

2. Apply the Pauli exclusion principle, the aufbau principle, and

Hund’s rule to write electron configurations using orbital

diagrams and electron configuration notation.

3. Define valence electrons and draw electron-dot structures

representing an atom’s valence electrons.

For the Test:

-be able to answer the objectives

-know all vocabulary

-be able to answer all review/practice questions

-study your notes (anything can be asked from them) 2

Atomic Orbitals

Orbitals do not have an exactly defined size, but are just unoccupied spaces available for electrons should the atom’s energy increase (or decrease)..

As the energy level increases, the orbital becomes larger

-n = 1, 2, … 7

The shape of the orbital is represented by the sublevels in that orbital

-s, p, d, & f

3

Atomic Orbitals

Each atomic orbital is designated by energy level and

sublevel:

1s represents ‘s’ orbital in the 1st energy level

each suborbital holds 2 electrons, each with

opposite spin

~ s = 2

~ p = 6, 3 suborbitals (2 per suborbital)

~ d = 10, 5 suborbitals (2 per suborbital)

~ f = 14, 7 suborbitals (2 per suborbital)

4

5

6

Electron Configurations

electron configuration: arrangement of electrons in an atom.

-electrons fill in so they have the lowest possible

energy

-there are three rules we follow

1. Aufbau principle: each electron occupies the lowest energy orbital available

-follows the aufbau diagram

♦all orbitals in a sublevel are equal energy

♦in multi-electron level, energy sublevels within a

principle energy level have different energies 7

♦in order of increasing energy, the sequence of

energy sublevels within a principle energy level is s,

p, d, and f

♦orbitals related to energy sublevels within one

principle level can overlap orbitals related to energy

sublevels within another principle level; begins in

n = 3

2. Pauli-exclusion principle: a maximum of 2

electrons may occupy a single orbital, but only if

they have opposite spins

-arrows pointing up (↑)and down (↓) represent

electron spins

-paired electrons in same orbital: represented by ↑↓.8

3. Hund’s rule: single electrons with the same spin

must occupy each equal energy orbital before

additional electrons with opposite spins can occupy

the same orbital

-s sublevel: max 1 unpaired e-

p sublevel: max 3 unpaired e-

d sublevel: max 5 unpaired e-

f sublevel: max 7 unpaired e-

9

Orbital Notation/Electron Configurations

There are three ways to indicate electron configurations.

1. orbital notation: an unoccupied orbital is represented by a line, ____, with the orbital’s principle quantum number and sublevel letter written underneath the line.

___

1s -an unpaired electron is shown as ____ and

1s

paired electrons as shown as ____ .

1s 10

Orbital Notation Practice

1. Mg

2. Al

3. Si

4. P

5. S

6. Cl

7. Ar

8. Cr

9. As

10. Kr11

2. electron-configuration notation: eliminates the lines and arrows of orbital notation

-the number of electrons in a sublevel is shown by adding a superscript to the sublevel designation;

exs. H configuration is 1s1

(1 electron on the 1s orbital) He is 1s2

(2 electrons in the 1s orbital) -use the periodic table to write electron configuration notation: s, p, d, and f block. -does not show orbital distributions of electrons

12

Electron Configuration Notation Practice

1. Mg

2. Al

3. Si

4. P

5. S

6. Cl

7. Ar

8. Cr

9. As

10. Kr13

Noble Gas Notation

3. noble-gas notation: shorthand notation using the noble gas elements

-noble-gas configuration: outer main energy level fully occupied by eight electrons, except in He

-ex: Na 1s22s22p63s1 is the electron-configuration

notation (long form)

[Ne]3s1 is the noble gas notation (short form)

because all orbitals are completely filled

up to neon, who’s orbital notation is

1s22s22p6

-this notation is only used for elements above neon! (those in n = 3 and above)

14

Noble Gas Notation Practice

1. Mg

2. Al

3. Si

4. P

5. S

6. Cl

7. Ar

8. Cr

9. As

10. Kr15

Exceptions

Some transition metals do not follow this trend.

-includes groups 6 and 11 (Cr and Cu groups)

Cr [Ar]4s23d4 Cu [Ar]4s23d9

Instead, we write it so that all s and d orbitals after the noble gas are half filled-increases stablity.

Cr [Ar]4s13d5 Cu [Ar]4s13d10

Practice with Exceptions:

1. Mo

2. Au

16

Valence Electrons

Only certain electrons determine an element’s chemical properties

-valence electrons: electron’s in the atom’s highest most energy level

•sulfur has 16 electrons, but only 6 valence electrons

Valence electrons can be determined by writing electrons configurations:

-S [Ne]3s23p4 6

-Ga [Ar]4s23d104p1 3 (4th level highest)

[Ar] 3d104s24p1

17

Electron Dot Structures

Chemists use valence electrons to show how atom’s are involved in bonding.

-electron dot structure: consists of the element’s

symbol (which represents the nucleus and inner electrons) surrounded by dots (that represent the

valence electrons)

•unpair the electrons first along the four sides of the

symbol

S

18

Oxidation Numbers

When an atom bonds, it either gains or loses electrons, forming a charge on the atom

-oxidation number: charge on an atom after bonding

•metals lose to form a positive charge

Ca+2 loses two electrons

•nonmetals gain to form a negative charge

N-3 gains three electrons

Look at the electron dot structure to determine how many it will lose or gain.

19