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What we’ve learned so far….Atoms lose/gain electrons to form cations
and anions (8 valence electrons like a noble gas!)
Charged anions and cations form IONIC BONDS to make IONIC COMPOUNDS
FK+ -
What we’re learning today…
Atoms can also SHARE electrons to form COVALENT BONDS!
Why is this important? Covalent bonds are important in many organic
molecules (living creatures) Examples:
Cell membranes are made of lipids (We’d be puddles of liquid without them!)
Proteins (We’d starve without them!)
Covalent BondThe chemical bond formed when two
atoms SHARE electrons Why do atoms form bonds? To gain the same
number of valence electrons as a noble gas!
ClClFull octet (8)! Full octet (8)!
Bonding Pair
Types of Covalent Bonds
Single Bond: ONE pair of shared electrons
Double Bond (Pi Bond): TWO pairs of shared electrons
Triple Bond: THREE pairs of shared electrons
ClCl
OO
N N
How many bonds can an atom form?
# bonds formed = # of UNPAIRED electrons
# electrons atoms keep = # electrons in PAIRS
N N3 unpaired electrons
= 3 bonds!
O 2 pairs = four electrons the atom keeps!
Example 1How many bonds can Sulfur form?
How many electrons will Sulfur keep to itself?
STwo unpaired electrons!
Four electrons will be kept by sulfur!
STwo pairs of electrons!
Two bonds can be formed!
Example 2How many bonds can Carbon form?
How many electrons will Carbon keep to itself?
CFour unpaired electrons!
Zero electrons will be kept by carbon!
CNo pairs of electrons!
Four bonds can be formed!
Example 3How many bonds can Neon form?
How many electrons will Neon keep to itself?
NeZero unpaired electrons!
Eight electrons will be kept by neon!
NeFour pairs of electrons!
Zero bonds can be formed!
What we have learned so far…Covalent Bonds: when two atoms SHARE pairs of
electronsWhy do atoms bond together? To get the same
number of valence electrons as a noble gas!How many bonds can an atom form?
Equal to the number of UNPAIRED electrons!
How many electrons do atoms keep to themselves? Equal to the number of PAIRED electrons!
Lewis structures are pictures of molecules Right number of valence electrons
Follows the octet rule
What we are doing today
Drawing our own Lewis structures!
Why are we doing this?
1. Organic molecules have covalent bonds!
2. We function because of covalent bonds!
3. Lewis structures tell us the SHAPE of molecules
Shape can determine how a molecule will behave
How do we draw a Lewis structure?
1. Determine what elements and how many atoms you have in a molecule from the formula
2. Draw the electron dot structures for every atom
The atom(s) with the most unpaired electrons go in the middle
3. Connect the dots! (unpaired electrons)
4. Redraw the structure neatly so lines are straight
Step 1: Determine the elements and number of atoms from the formula
Symbols tell you what elements we haveLittle numbers tell you how many atoms of
each element
Example:
H2O2 Hydrogen
atoms!1 Oxygen
atom!
Step 2: Draw the electron dot structures for EVERY atom
Put the atom(s) with the most UNPAIRED electrons in the middle!
Example: H2O
(2 hydrogens, 1 oxygen)
H H O
Step 3: Connect the dots! (unpaired electrons)
Connect one unpaired electron from atom to an unpaired electron of another atom
These lines show COVALENT BONDS BETWEEN ATOMS!
Example: H2O
H HO
Step 4: Redraw the structureLines are straightLeave paired electrons in the structure
Example: H2O
H HO H HO
“Do Now” for 2/22/10
Draw the Lewis structures for the following molecules (look in your notes or p. 254 in your text)
C2H6 (common name ‘ethane’, an alkane) C2H4 (common name ‘ethylene’, an alkene) C2H2 (common name ‘acetylene’, an alkyne)
Bonus Point: What is the special name for the group of compounds that these 3 molecules fall under?
When you are done, QUICKLY finish drawing the compounds from Friday in your notes (or on your worksheet) so we can go over them:
H2S H2O2 NH2Cl HCN
“Do Now” response cont.
C2H2The name for the
group of organic compounds that all 3 of these molecules belong to is:
C CH H
C C HH
Hydrocarbons
Def: Simplest organic compounds composed of hydrogen and carbon EXCLUSIVELY.
Announcements We will be finishing Ch. 8 this week If by the end of class on Wed. you are struggling
with ionic or covalent bonds, it is highly recommended that you attend an after school review session
Response to lab ?’s are due TODAY New HW was posted in class and online last Fri and
is due on Thurs, 2/25 Exam on Ch. 7 and 8 is scheduled for THIS
Friday, 2/26, review guides on table – complete by Thursday!
You do NOT need to complete the graphic organizer for steps to naming covalent compounds (Ch. 8.2 HW). You may for EC if you like.
You NEED your book Tues-Thurs. No book = no chair for you! (until you have a chance
to correctly answer a ? at least)
Subscripts, Superscripts and Coefficients
Al3(SO4 )2coefficient
subscripts
superscripts
ALUMINUM SULFATE
SO43- Al2+
5
Title: Ch. 8 Overview Cornell Notes
Review: What is the difference between a compound and a molecule?
Compound: term used to describe elements that have ionic bonds
chemical combination of 2 or more different elements
can be broken down into simpler substances has properties different than those of the
elements which make it up
Molecule: term used to describe atoms of either the same or different elements that share e- (are covalently bonded)
Strength of covalent bondsBond Length
Def: Distance from the center of 1 nucleus to the center of another between bonded atoms.
As # of shared e- pairs _________, bond length_________.
As bond length _________, the strength of the bond _________.
Bonds and energy If you form a bond energy is released. If you break a bond energy must be added, or
used. The energy required to break a bond is called
bond-dissociation energy.
increasesdecreases
increasesdecreases
Molecular Structures Quickly sketch the graphic below of the different
types of models for showing molecules in your notes.
Lewis structures for Polyatomic ions
Though polyatomic ions as a unit form ionic bonds, the elements that make them up are covalently bonded to each other.
Main difference between a molecule and a polyatomic ion: difference in the # of e- that are available to bond (hence the + or – charges)
Steps:1. Determine the # of e-
Look at the # of valence e- each atom should have and add them together to get the total # of e-
S O4 2-
6 (x1) 6(x4)+ = 30 total 2. Check the charge on the polyatomic atom
on your cheat sheet Add it to your # of e- if the charge is - Subtract it from your # of e- if the charge is +
30 + 2 = 32 total valence e-
3. Determine which will be the central atom (Remember – it will be the atom that wants to make the most bonds!)
• Both S and O have ____ valence e-, so both want to make ____ bonds.• So what do we do??
• The element with the lowest electronegativity will be the central atom.• Remember: electronegativity is an atom’s ability to attract e-• Look at p. 194 in Ch. 6.3 for the trend…which you SHOULD
have memorized!
• General rules:• H will always be a terminal atom• 1st element listed will usually be the
central atom
26
In this case, it is S
4. Draw single bonds around the central atom Connect it to all surrounding atoms
5. Add the remaining e- around the terminal atoms
Fulfill the octet rule (unless H or He!)
S O4 2-
S OO
OO
8 e- used in the single bonds
How many do we still need to place?
24 e-
Is the octet rule satisfied?(Does each atom have 8 e-?)
Yep!
“Do Now” for 2/231. Which of the following molecules has the
strongest bonds? 1st draw the Lewis structures for the molecules 2nd look back over your notes to see how to determine
bond strength. (p. 246 in text if you were absent)
a.) CN- b.) CH3NH2Hint: the C and the N are
central atoms
C N C NH
HH
H
H
As the # of shared pairs of e- increases, the bond length decreases, which means the bond is stronger.
Naming Covalent MoleculesCh. 8.2
Binary Molecular Compounds (2 nonmetal atoms)1. Name the first element in the
compound2. Name the 2nd element using the
root and the suffix “ide” (Just like in monatomic ionic compounds!)
3. Use prefixes to denote the # of each of the elements.
Prefixes (Memorize these! p. 248)
# of Atoms Prefix
1 Mono (only used for Oxygen)
2 di
3 tri
4 tetra
5 penta
6 hexa
7 hepta
8 octa
9 nona
10 deca
Let’s Practice! (Answer in your notes)
1. CO2
2. SO2
3. CCl4
3. P2O5
Carbon
Sulfur
Carbon
phosphorous
oxide
oxide
oxide
chloride
di
di
tetra
pentdi
Naming Binary Acids(H + 1 other element) Write these down!!
1. Add the prefix “hydro” to the name of the 2nd element
2. Remainder of 1st word is the root of the 2nd element with the suffix “ic”
3. The 2nd word is always acid.
HCl
Hydrochloric acid
Naming Oxyacids (H + an oxyanion) Oxyanion = polyatomic ion that contains 1 or
more oxygen atoms.
Write these down!!1. Identify the oxyanion present.2. 1st word of name begins w/ the root of
the oxyanion (and the prefix “per” or “hypo” if it is part of the name)
3. The suffix of the 1st word is: “ic” if the oxyanion ends w/ “ate” “ous” if the oxyanion ends w/ “ite”
4. The 2nd word is always “acid”.
HNO3
Nitrateic acidNitr