5-3Electron Configurations and

Periodic Properties

PreClass: What does the numerical value of the period (ie. Period 4) represent in addition to just physically locating a particular row on the Periodic Table?

• The period = n value of the outermost electrons

• Ie. If n=4, than all elements on the 4th period have their outermost electrons in the 4 energy level !!

Beaker Breaker

1. Mendeleev’s periodic table had the elements ordered in increasing ________

2. What is the element in the same Period that follows potassium, K ?

3. What is the name of the group on the far left of the Periodic table?

4. What element ends with 3p1 ?

Atomic Radius

• One-half the distance between the nuclei of identical atoms that are bonded together

Group Trend for Atomic Radii

• Down a group the radius of an atom gets larger

• The principal quantum number increases and the cloud grows by one shell

Period Trend for Atomic Radii

• Across a period the radius of an atom gets smaller

• You’re adding electrons to approximately the SAME region (same “n”) while the electrons are being pulled in more tightly as nuclear charge increases (more protons)

• Exception: Noble Gas Family – atoms don’t interact and pull together like in other atoms because atoms already have full outer level

Questions

• Of the elements Li, O, C, and F, identify the one with the largest atomic radius and the one with the smallest atomic radius

• Of the elements Br, At, F, I and Cl, identify the one with the largest atomic radius and the one with the smallest atomic radius

Answers

• Li = largest F=smallest

• F = smallest At = largest

Why does the atomic radii not decrease very much as you move

across the d-sublevel?

• Increasing # inner electrons shield outer level electrons from nucleus

Why would the atomic radius of hafnium (#72) be LESS than that

of zirconium (#40)?

• Hf has such a greater nuclear charge

Beaker Breaker

1. Which has the largest and which has the smallest atomic radius of the following:

C, Ge, Sn, Si

2. Which has the largest and which has the smallest atomic radius of the following:

K, Cu, As, Br

Ionization energy, IE

• Energy required to remove an electron from a neutral atom (or first ionization energy, IE1)

• Ion: charged particle

• Ionization: process of an electron being lost or gained from an atom which results in the formation of an ion (Na+

and Cl-)

Group Trend for Ionization Energies

• Down a group the ionization energy of an atom generally gets smaller

• Electrons are at a greater distance from the nucleus

• Outer electrons are shielded from the nucleus by inner electrons

Period Trend for Ionization Energies

• Across a period the ionization energy increases

• Nuclear charge gets greater while atomic size decreases

Examples

• Which of the following has the largest ionization energy?

Na, Mg, P, Cl

Which of the following has the smallest ionization energy?

Be, Mg, Ca ,Sr

Homework

• Pg 156-157

# 28 (all), 29 (all), 30 (all), 31 (a &b only)

Beaker Breaker

• Which alkaline earth metal has the lowest ionization energy?

• Which halogen has the largest atomic radius?

What will determine whether an electron is easily lost or not?

(4 things)

4 Factors Affecting Ionization Energy

1. Nuclear charge (greater the charge, the greater the IE)

2. Shielding Effect (greater the shielding, the lower the IE)

3. Radius (the greater the radius, the less the IE)

4. Sublevel configuration (an electron from a half-full or full sublevel requires more IE)

2nd Ionization energy, IE2 is the energy required to take a 2nd electron away from an atom.

Why is IE2 always greater than IE1?Why is the IE2 of Na so much greater that the greater of Mg ??(Hint: look at the electron config.)

Multiple Ionization Energies

• Energy required to remove the 2nd, 3rd, etc. electron from an atom

• IE3 > IE2 > IE1 because remaining electrons will be held more tightly as the electron repulsion decreases and the cloud is pulled in more tightly

Why does the IE go UP as you go down a d-sublevel group but

go down in an s-sublevel?

• “f” sublevel has minimal shielding effect while the nuclear charge continues to grow (“Lanthanide Contraction”)

Electron Affinity p. 147• Energy change that occurs when an electron is

acquired by a neutral atom

• Most atoms RELEASE energy when they acquire an electron: A + e- A- + energy– energy has negative sign

• Some atoms gain energy when they acquire an electron: A + e- + energy A-

– energy has positive sign; atom is unstable & loses electron spontaneously

Group Trend for Electron Affinities

• Down a group the electron affinity of an atom tends to get smaller

• Although there is an increasing nuclear charge, there are more levels so the size is greater

• Adding an energy level usually dominates!

Period Trend for Electron Affinity

• Across the p-sublevel, the energy change increases (becomes more negative)

–Electron config is close to being full and the size is smaller

Why is the electron affinity of nitrogen so low when compared

to carbon or oxygen?

• Adding an electron to carbon half fills the 2p sublevel

• Adding an electron to nitrogen forces the config to go from stable (half-filled) to less stable (no spec arrangement)

Multiple Electron Affinities

• It is always more difficult to add a 2nd electron to an already negatively charged ion..therefore, all 2nd electron affinities are positive

Beaker Breaker

1. Which alkaline earth metal has the largest electron affinity?

2. Which of the following has the smallest electron affinity Na, Mg, P or Cl

Ionic Radii of Cations

• Cation: positive ion

• Formed by an atom losing electron(s)

• Always smaller because electron cloud is smaller (less repulsion) & sometimes even one less energy level!

• Anion: negative ion

• Formed by an atom gaining electron(s)

• Always larger because electron cloud is greater (more repulsion among electrons)

Ionic Radii of Anions

Group Trend for Ionic Radii

• Down a group the ionic radius of an atom generally gets larger

• Electrons are at a greater distance from the nucleus (higher E level) and have more shielding

Period Trend for Ionic Radii

• Metals (left side):form cations

– Cationic Radius: Decreasing ionic radius as nuclear charge increases without adding an energy level

• Non-metals: form anions

– Anionic Radius:Decreasing ionic radius as nuclear charge increases without adding an energy level

Practice

• Pg 157 # 41

• Pg 157 # 46

PC: What are VALENCE electrons?

Valence Electrons

• Electrons available to be lost, gained, or shared in the formation of chemical compounds

• Often the outermost electrons because they are held most loosely

What would be the # of valence electrons in……..

• Calcium

• Lithium

• Chlorine

• carbon

What would be the # of valence electrons in……..

• Calcium – 2 : 4s2

• Lithium – 1: 2s1

• Chlorine – 7: 3s23p5

• Carbon – 4: 2s22p2

Beaker Breaker

1. How many valence electrons does Te have?

2. Which has a smaller ionic radius Na+1 or

Ca+1 ?

Electronegativity

• Measure of the ability of an atom in a chemical compound to attract electrons

• Fluorine is the MOST electronegative element – assigned an arbitrary value of 4.0

• All other values are relative to F

• 3 highest values: F – O - N

Group Trend for Electronegativity

• Tend to decrease down a group (or stay the same) as the atoms gets larger

Period Trend for Electronegativity

• Tend to increase across the period as the atoms gets smaller, the nuclear charge becomes greater, and the atom is getting closer to a noble gas configuration

Determine the likely charge for the following elements: Ca, O, Al

1. Write the noble gas configuration of the element

2. Determine if electrons will be LOST or GAINED to make the element stable

3. ID the noble gas whose electron configuration by losing/gaining these electrons

4. Write the formula for the ion

5. ID it as a cation OR anion

Determine the likely charge for the following elements: Ca, O, Al

• Ca: [Ar]4s2

• Ca will LOSE 2 e- • Ca now has the Ar config• Ca+2 • cation• O: [He]2s22p4

• O will GAIN 2 e-

• O now has a Ne config• O-2 , an anion

Al

• Al: [Ne]3s23p1

• Al will LOSE 3 e-

• Al now has a Ne config

• Al+3 , a cation

How do d-block elements form ions?

• Electrons in the highest occupied sublevel are always removed first

Why does zinc become a +2 ion?

Which electrons are lost when titanium becomes a +2, a +3 and a

+4 ion?

Beaker BreakerWhich one is larger?

• Na or K

• Na or Mg

• Na or Na+

• O or F-

• O or O-2

Which one is larger?• Na or K

• K

• Na or Mg

• Na

• Na or Na+

• Na

• O or F-

• F-

• O or O-2

• O-2