Chapter 8 Review of Quantum Numbers Principal Quantum Number (n) -tells you the energy level -n can be equal to 1, 2, 3, 4, 5, 6, 7… -distance e- is from.

Chapter 8Review of Quantum Numbers

Principal Quantum Number (n)

-tells you the energy level

-n can be equal to 1, 2, 3, 4, 5, 6, 7…

-distance e- is from the nucleus inc. as n inc.

Angular Momentum Quantum Number (l)

-determines the shape of the orbital

-shapes are s, p, d, or f

-when given a value of n, l can be any integer including zero up to n - 1

Value of l Shape of orbital

l = 0 s

l = 1 p

l = 2 d

l = 3 f

Magnetic Quantum Number (ml)

-specifies the orientation of the orbital

-equal to integer values, including zero ranging from +l to -l

Examples:

1) What are the quantum numbers of the orbitals in the 3rd energy level?

n = 3

l = 2, 1, 0

ml = +2, +1, 0, -1, -2 (represents d orbitals)

+1, 0, -1 (represents p orbitals)

0 (represents s orbital)

2) What are the quantum numbers of the 4p orbital?

n = 4

l = 1 (because it is in the p orbital)

ml = +1, 0, -1 (p orbitals have three orientations)

Spin Quantum Number (ms)

-electron spin is represented by the direction of the arrow (which represents the electrons)

-all e- have the same amount of spin

-electrons can only spin in one of two directions- spin up or spin down

ms = +1/2 (spin up)

ms = -1/2 (spin down)

**Write the electron configurations for potassium and titanium.

potassium 19e-1s22s22p63s23p64s1

**shorthand way = [Ar] 4s1 titanium 22e-1s22s22p63s23p63d24s2

**shorthand way = [Ar] 3d24s2

Question: What are the four quantum numbers for each of the two e- in a 4s orbital?

n = 4 l = 0 ml = 0 ms= +1/2

n = 4 l = 0 ml = 0 ms= -1/2

Orbital Diagrams

-shows arrangement of electrons in orbitals

-symbolizes electrons as arrows and orbitals as boxes

Rules for orbital diagrams:

1) Aufbau Principle

-electrons enter orbitals of lower energy first

s p d f

-atomic orbitals are represented as boxes

s = 1 box (1 orbital) p = 3 boxes (3 orbitals)

d = 5 boxes (5 orbitals) f = 7 boxes (7 orbitals)

2) Pauli-Exclusion Principle

-an atomic orbital can hold at most 2 electrons

-electrons are represented as arrows

-spins are opposite

-first electron is +1/2 ↑

-second electron is -1/2 ↓

-number of e- must equal number of arrows

3) Hund’s Rule

-one electron enters each orbital of equal energy until orbitals contain one electron, then they can hold two e-

-it is more stable to have partially filled orbitals than empty orbitals

**Draw orbital diagrams for beryllium and sulfur.

**Draw orbital diagrams for potassium and titanium.

Electron Configuration and the Periodic Table

periodic property- property that is predictable based on an element’s position within the periodic table

Modern periodic table is set up according to Dmitri Mendeleev’s:

periodic law- when elements are arranged in order of increasing mass, they arrange into groups with other elements having similar properties

-Henry Moseley later said it would be better to arrange according to increasing atomic number because not all masses are greater as you move across

Ex- tellurium and iodine

valence electrons- electrons in the outermost energy levels (highest energy level)

-important for chemical bonding because they are held most loosely and are easier to share or lose

-elements in the same group have similar # of valence e- and similar chemical properties

-in transition metals the d e- are included in the valence electrons even though they are not in the outermost energy level

core electrons- all other e- besides the valence e-

*Identify the valence and core e- for potassium, titanium and germanium

K = 1 valence e- and 18 core e-

Ti = 4 valence e- and 18 core e-

Ge = 4 valence e- and 28 core e-

-electron configurations can determine the group of the element on the periodic table

alkali metals = ns1

alkaline Earth metals = ns2

transition metals = d block

halogens = np5

noble gases = np6

inner transition metals = f block

**Predict the outer e- config for each element:

1) strontium 2) bromine 3) cadmium

1) 5s2 2) 4s24p5 3) 4d105s2

Summary

-periodic table is divided into four blocks (s, p, d, and f)

-the group # of a main-group element is equal to the number of valence e-

-the row # of a main-group element is equal to the highest principle quantum # of that element

Periodic Trends

1) Atomic Size

-looking at atomic radius:

-half the distance between the nuclei of two atoms bonded together

Trend:

1) atomic radius tends to increase as you move down a group

-as you move down a group, the n value (energy level) increases resulting in larger atoms

2) atomic radius tends to decrease as you move across a period

-because there are more valence e- as you move across, there is a stronger attraction between the outermost e- and the nucleus and it makes it more tightly bound and therefore smaller

Examples:

1) Choose the larger atom for these pairs:

a) nitrogen or fluorine

N

b) carbon or germanium

Ge

c) nitrogen or aluminum

Aℓ

d) aluminum or germanium

unable to tell based on trends

2) Choose the larger atom:

a) tin or iodine b) germanium or polonium

c) iron or selenium d) chromium or tungsten

d) Sn b) cannot tell c) Fe d) W

3) Place in order of decreasing radius:

sulfur, calcium, fluorine, rubidium, silicon

4) Place in order of increasing radius:

nitrogen, lithium, carbon, oxygen, beryllium

5) Rb, Ca, Si, S, F

6) O, N, C, Be, Li

Electron Configs and Magnetic Properties of Ions

-remember ions are atoms or groups of atoms that have either lost or gained e- and have a charge

-when forming ions, atoms try to achieve e- config of closest noble gas

Ex- write e- config of a fluoride ion

F1- = 1s22s22p6 e- config of neon

Try an aluminum ion

Aℓ3+ = 1s22s22p6 e- config of neon

-transition metal cations lose e- in a different way

-the 4s will lose its e- before the 3d even though the 4s is in a higher energy level

Ex- vanadium ion = V2+

1s22s22p63s23p63d3

paramagnetic- when atoms or ions have unpaired e- in their e- configs (have an s, p, d or f orbital only partially filled)

diamagnetic- when all e- are paired in an atom or ion’s e- config (all orbitals contain max amount of e-)

Write e- configs and orbital diagrams for the following ions and determine if they are diamagnetic or paramagnetic.

1) Ga3+

1s22s22p63s23p63d10 diamagnetic

2) S2-

1s22s22p63s23p6 diamagnetic

3) Fe3+

1s22s22p63s23p63d5 paramagnetic

Trends Continued:

2) Ionic Size

-cations are much smaller than their corresponding atoms

-anions are much larger than their corresponding atoms

**as you move across a period the ionic size will decrease (comparing cations to cations or anions to anions)

**as you move down a group ionic size will increase

1) Choose the larger ion or atom

a) S or S2- b) Ca or Ca2+ c) Br or Br-

a) S2- b) Ca c) Br-

2) Arrange in order of increasing ionic size:

Ca, Sr, Be, Mg, Ba (all have 2+ charge)

Be, Mg, Ca, Sr, Ba

Trends Continued:

3) Metallic Character

-as you move down a group, metallic character increases

-as you move across a period, metallic character decreases

-makes sense with distribution of metals on the periodic table

Examples:

Choose the more metallic element:

1) tin/tellurium 2) phosphorus/antimony

3) germanium/indium 4) sulfur/bromine

4) Sn 2) Sb 3) In 4) cannot tell

Arrange in order of increasing metallic character:

silicon, chlorine, sodium, rubidium

Cℓ, Si, Na, Rb

Trends Continued:

4) Ionization Energy (IE)

-the energy required to remove an e- from the atom or ion in the gaseous state

first IE- energy needed to remove the first e-

-IE tends to decrease as you move down a family b/c e- in the outermost energy level become farther away from the + charged nucleus and are held less tightly

-IE tends to increase as you move across a period b/c valence e- experience greater attraction with the nucleus

Examples:

Choose the element with the higher first IE.

1) aluminum/sulfur 2) arsenic/antimony

3) nitrogen/silicon 4) oxygen/chlorine

5) tin/iodine 6) carbon/phosphorus

4) S 2) As 3) N

4) cannot tell 5) I 6) cannot tell

Put in order of decreasing first IE:

sulfur, calcium, fluorine, rubidium, silicon

F, S, Si, Ca ,Rb

second ionization energy- energy needed to remove the second electron

**Read page 347 on second and successive IE’s

Electron affinity (EA)

- energy change associated with the gaining of an e- by an atom in the gaseous state

**Read pages 352-356