Post on 19-Dec-2015
transcript
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.