Today is Tuesday,March 10th, 2015

Pre-Class:You’ve probably heard of the special name we give to

electrons in the outermost principal quantum number.

Do you remember what it is?

Also get a small piece of a paper towel for you/your partner.

Stuff You Need:Periodic TablePaper Towel

In This Lesson:Valence Electrons

and Lewis Dot Structures

(Lesson 4 of 4)

Today’s Agenda

• Study Guide for Core Assessment• Valence Electrons• Lewis Dot Structures

• Where is this in my book?– P. 187 and following…

By the end of this lesson…

• You should be able to draw valence electrons as Lewis Dot Structures.

Valence Electrons

• Valence electrons are the ones available for bonding.

• Notice where they are:

Da ‘portant Stufz

• Valence electrons are electrons in the outermost shell (highest energy level).– The highest “coefficient.”

• IMPORTANT NOTE:– d and f sublevel electrons do not figure in bonding

because we only look at the highest principal quantum number electrons.

Finding Valence Electrons

• To find valence electrons, simply perform the usual electron configuration notation.

• Find the sublevel(s) with the highest principal quantum number. Count the electrons there, ignoring d or f sublevels, if any.

• Example: 1s22s22p4

• 2s22p4 are the valence electrons (6 total)

Determining Valence Electrons

• Let’s try some.• Grab your periodic tables and whiteboards.• Tell me the number of valence electrons in the

following elements [next slide].

Valence ElectronsElement Electron Configuration # of Valence

Electrons/Capacity

Oxygen (O) 1s22s22p4 6/8

Hydrogen (H) 1s1 1/2

Xenon (Xe) 1s22s22p63s23p64s23d104p65s24d105p6 8/8

Rubidium (Rb) 1s22s22p63s23p64s23d104p65s1 1/8

Helium (He) 1s2 2/2

Boron (B) 1s22s22p1 3/8

Carbon (C) 1s22s22p2 4/8

Fluorine (F) 1s22s22p5 7/8

Aluminum (Al) 1s22s22p63s23p1 3/8

More Practice

• Electrons Review Worksheet– Do the first page (landscape orientation), but only

these columns:• Atomic Number• Electron Configuration• Number of valence electrons

The Octet Rule

• One other thing…• Remember when we said that atoms want a full

valence electron shell like those super awesome noble gases?

• Well, atoms “want” to be like noble gases because a full valence shell makes them more stable than having a partial valence shell.– And who doesn’t like stability?

• They do this by adding or dropping electrons.

The Octet Rule

• What is the electron capacity of a full s sublevel plus a full p sublevel?– 8 (s=2, p=6)

• This idea, of having 8 electrons in the valence shell to be full, is called The Octet Rule.– Note: Hydrogen and helium are exceptions. What

is capacity for their valence shells?• 2, so they only want 2 electrons to be stable.

Group IA (alkali metals) have 1 valence electron (1+)

Group IIA (alkaline earth metals) have 2 valence electrons (2+)

Group IIIA elements have 3 valence electrons (3+)

Group IVA elements have 4 valence electrons (4+)**

Group VA elements have 5 valence electrons (3-)

Group VIA elements have 6 valence electrons (2-)

Group VIIA (halogens) have 7 valence electrons (1-)

Group VIIIA (Noble gases) have 8 valence electrons, except helium, which has only 2 (no charge)

Transition metals (“d” block) have 1 or 2 valence electrons (1+ or 2+)

Lanthanides and actinides (“f” block) have 1 or 2 valence electrons (1+ or 2+)

About transition metals

• Transition metals do weird things.– Yes, they do have 1 or 2 valence electrons, but

they form lots of different ionic charges.• The first thing to be aware of is that while full

energy levels are the most stable, half-filled sublevels are still mostly stable.

• To understand this better, let’s take a look at Cu, Fe, and Mn.

• Copper’s valence orbital notation:

• So you can see that copper has a full s sublevel but only an almost full d sublevel.– We would expect it to drop it the two 4s electrons,

making a charge of 2+.• However, because a half-filled sublevel is

preferable to this setup, Copper flips one electron up from the 4s sublevel to fill 3d.

Copper (Cu)

____ ____ ____ ____ ____ ____

4s 3d

Copper (Cu)

• Then, copper can just drop the s electron.

• As a result of these two possibilities, copper can have two possible ionic charges: 2+ or 1+.

____ ____ ____ ____ ____ ____

4s 3d

Iron (Fe)

• Iron does something kinda similar:

• For iron, the first thing it can do is drop both s electrons.– 2+ ion.

• Or, it could drop both s electrons and one d.– 3+ ion.

____ ____ ____ ____ ____ ____

4s 3d

Fun Fact: Iron (Fe)

• The fact that iron has four unpaired electrons in its d sublevel is the reason iron is/can be magnetized at room temperature (along with cobalt and nickel).

• Having four electrons spinning all in the same direction makes for easy magnetic field induction, and it’s called having an orbital magnetic moment.

• Note that there are other factors at play here, one of which is the “sea of electrons” concept you’ll learn next unit.

Manganese (Mn)

• Now for Manganese:

• Manganese can drop both s electrons.– 2+ ion.

• Or, it can drop all 7 electrons.– 7+ ion.

• Or about five other possibilities!

____ ____ ____ ____ ____ ____

4s 3d

Multivalent Elements

• These, and other metals, are multivalent – they have several different configurations of their valence electrons.

• Therefore, they form multiple charges.• Here’s a present for you – a periodic table with

a listing of multivalent metals:– Periodic Table – Polyatomic Ions and Multivalent

Elements Only

About Group IVA

• Group IVA also does some weird things.• Carbon, for example, would like to either gain

or lose 4 electrons. But how many does it have total?– 6

• So gaining/losing 4 is kinda hard for such a small atom.– Even Si is too small.

• So C and Si share electrons instead of losing or gaining.

About Group IVA

• However, Ge, Sn, and Pb are all big enough to ionize.– Their outer electrons are very far away, and what’s 4/82

to lead?• So, Ge tends to lose all four valence electrons.– 4+ ion.

• Sn and Pb either lose all four…– 4+ ion.

• …or just lose the p sublevel electrons.– 2+ ion.

• So Sn and Pb have two different possible oxidation states (ionic charges).

About Group IVA

• In addition to Group IVA, other nearby large elements do the same sort of thing:– Antimony (Sb) – 3+ or 5+

• 3+ = p dropped; 5+ = s and p dropped.

– Bismuth (Bi) – 3+ or 5+• 3+ = p dropped; 5+ = s and p dropped.

– Thallium (Tl) – 1+ or 3+• 1+ = p dropped; 3+ = s and p dropped.

– Polonium (Po) – 2+ or 4+• 2+ = one from p dropped, one from s dropped; 5+ = all of

p dropped and one from s.

More Practice

• Electrons Review Worksheet– Finish the first page.

Lithium

LiBeryllium

BeBoron

BCarbon

C

Nitrogen

NOxygen

OFluorine

FNeon

Ne

Lewis Dot Notation

• Lewis Dot Notations for Period 2 elements.

Creating Lewis Dot Structures

• Step 1: Determine the number of valence electrons.

• Step 2: Draw them around the element abbreviation one-by-one.

• Step 3: Check your answer. Make sure you only have electron pairs if you already have four electron “singles.”

Lewis Dot Practice

Cl Se Al

K Si Ca

Practice

• Electrons Review Worksheet– Try the reverse side.– Tough ones: Zn, Ag, Fe (TRY THEM!)

Summary

• Valence Electrons– Electrons in the outermost energy level (highest n

number) – just s and p sublevels.– Electrons can be shown via Lewis Dot Diagrams.

• Octet Rule– Atoms react to get eight electrons in their

outermost shell.• Except H and He.

– That makes ‘em stable.

Closure Part 1

• Draw the dot structure for the element Bromine:

• Draw the dot structure for the element Thallium:

• Draw the dot structure for the element Selenium:

• Draw the dot structure for the element Magnesium:

Closure Part 2

• Draw the dot structure for the element Potassium:

• Draw the dot structure for the element Helium:

• Draw the dot structure for the element Aluminum:

• Draw the dot structure for the element Hydrogen:

Closure Part 3

• If you have an atom on the right side of the table, let’s say Chlorine, how many electrons does it need to get to 8?– 1

• From where might it get that electron?– A cation.

• And in which group would that cation be?– Alkali metals (Group I), because they each have one electron

they’d like to give away.• And what would be a possible “donor” element?– Na, Li, K, Rb, Cs, et cetera. They make salts like in our flame

tests!

Closure Part 4

• Electron Configuration:– 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p1

• Shorthand Electron Configuration:– [Kr] 5s2 4d10 5p1

• Valence Electron Configuration:– 5s2 5p1

• Orbital Notation:– [arrows]

• Dot Notation:– Uh…dots.