Chapter 8 Review of Quantum Numbers Principal Quantum

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 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 3 rd 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 4 p 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 directionsspin up or spin down ms = +1/2 (spin up) ms = -1/2 (spin down)

**Write the electron configurations for potassium and titanium. potassium 19 e 1 s 22 p 63 s 23 p 64 s 1 **shorthand way = [Ar] 4 s 1 titanium 22 e 1 s 22 p 63 s 23 p 63 d 24 s 2 **shorthand way = [Ar] 3 d 24 s 2 Question: What are the four quantum numbers for each of the two e- in a 4 s 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 = ns 1 alkaline Earth metals = ns 2 transition metals = d block halogens = np 5 noble gases = np 6 inner transition metals = f block

**Predict the outer e- config for each element: 1) strontium 2) bromine 3) cadmium 1) 5 s 2 2) 4 s 24 p 5 3) 4 d 105 s 2 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 principal 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 a) 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 3) Rb, Ca, Si, S, F 4) 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 F 1 - = 1 s 22 p 6 e- config of neon Try an aluminum ion Aℓ 3+ = 1 s 22 p 6 e- config of neon

-transition metal cations lose e- in a different way -the 4 s will lose its e- before the 3 d even though the 4 s is in a higher energy level Ex- vanadium ion = V 2+ 1 s 22 p 63 s 23 p 63 d 3 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) Ga 3+ 1 s 22 p 63 s 23 p 63 d 10 diamagnetic 2) S 21 s 22 p 63 s 23 p 6 diamagnetic 3) Fe 3+ 1 s 22 p 63 s 23 p 63 d 5 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 S 2 - b) Ca or Ca 2+ c) Br or Bra) S 2 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 1) 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 1) 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