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 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)

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

Figure: 08-07-02UN Title: Small periodic table for Ar, Se. Caption: The valence electrons can be determined by the row and orbital block location of the element.

Write the electron configurations for potassium and titanium **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: 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)

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

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 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: 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 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 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: Choose the larger atom for these pairs: nitrogen or fluorine N carbon or germanium Ge c) nitrogen or aluminum Aℓ aluminum or germanium unable to tell based on trends

Choose the larger atom: tin or iodine b) germanium or polonium iron or selenium d) chromium or tungsten Sn b) cannot tell c) Fe d) W Place in order of decreasing radius: sulfur, calcium, fluorine, rubidium, silicon Place in order of increasing radius: nitrogen, lithium, carbon, oxygen, beryllium Rb, Ca, Si, S, F 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. Ga3+ 1s22s22p63s23p63d10 diamagnetic 2) S2- 1s22s22p63s23p6 diamagnetic 3) Fe3+ 1s22s22p63s23p63d5 paramagnetic

Trends Continued: 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

Choose the larger ion or atom S or S2- b) Ca or Ca2+ c) Br or Br- a) S2- b) Ca c) Br- Arrange in order of increasing ionic size: Ca, Sr, Be, Mg, Ba (all have 2+ charge) Be, Mg, Ca, Sr, Ba

Trends Continued: 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: tin/tellurium 2) phosphorus/antimony germanium/indium 4) sulfur/bromine 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. aluminum/sulfur 2) arsenic/antimony nitrogen/silicon 4) oxygen/chlorine 5) tin/iodine 6) carbon/phosphorus S 2) As 3) N 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 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