In This Lesson:
Valence Electrons
and Lewis Dot
Structures
(Lesson 4 of 4)
Today is Thursday,
March 3rd, 2016
Stuff You Need:
Periodic Table
Paper Towel
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.
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 Electrons
Element
Electron Configuration
# of Valence
Electrons/Capacity
Oxygen (O)
1s22s22p4
6/8
Hydrogen (H)
1s1
1/2
Xenon (Xe)
1s22s22p63s23p64s23d104p6
5s24d105p6
8/8
Rubidium (Rb)
1s22s22p63s23p64s23d104p6
5s1
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 (Cu)
• Copper’s valence orbital notation:
____
4s
____ ____ ____ ____ ____
3d
• 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)
• Then, copper can just drop the s electron.
____
4s
____ ____ ____ ____ ____
3d
• As a result of these two possibilities, copper
can have two possible ionic charges: 2+ or 1+.
Iron (Fe)
• Iron does something kinda similar:
____
4s
____ ____ ____ ____ ____
3d
• 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.
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 and
others).
• 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:
____
4s
____ ____ ____ ____ ____
3d
• Manganese can drop both s electrons.
– 2+ ion.
• Or, it can drop all 7 electrons.
– 7+ ion.
• Or about five other possibilities!
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 5+
• 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.
Lewis Dot Notation
• Lewis Dot Notations for Period 2 elements.
Lithium
Beryllium
Boron
Carbon
Li
Be
B
C
Nitrogen
Oxygen
Fluorine
Neon
N
O
F
Ne
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.