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Chem 105 notes March 25-April 1st

by: Allie Evey

Chem 105 notes March 25-April 1st Chem 105

Marketplace > Washington State University > Chemistry > Chem 105 > Chem 105 notes March 25 April 1st
Allie Evey
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About this Document

These notes cover March 25-April1st
Chem 105
Class Notes
Chem, 105, Finnegan, notes, Chemistry




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This 10 page Class Notes was uploaded by Allie Evey on Saturday March 5, 2016. The Class Notes belongs to Chem 105 at Washington State University taught by Finnegan in Spring2015. Since its upload, it has received 121 views. For similar materials see Chem 105 in Chemistry at Washington State University.


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Date Created: 03/05/16
Lewis structures for covalently bonded compounds  Bonds are shown as lines  Lone pair electrons are pairs of dots  All atoms(except H) want 8 electrons(octet rule)  An atom usually forms one bond for each electron  Central Atom Premise: the central atom is the least electronegative(except hydrogen, it only has one electron) Example: CCl 4 Note the lines representing the bonds, and the dots representing lone electrons that aren’t bonded Important Rules  Molecules with multiple Carbon and Nitrogen will usually have multiple centers  Carbon is the ONLY element the forms long chains (ie. Sugars)  Avoid O—O and F—F bonds, NEVER EVER make chains(2 or more) of Oxygen and Fluorine!  Fluorine does not form double bonds, Ever Example: Draw C 3 O8 Example Draw C H 3 6 Cyclopropane is highly explosive, because of the triangular structure  DO NOT make triangles or squares with your bonds.  Pentagons and hexagons are okay.  Remember, rules can be broken, but when you do that bad things start happening. The more rules you break the worse you compound is. Draw C3H 6 It is an ISOMER, it has many different ways of drawing it, all of which are correct, but different With C 3 6 it is really easy to draw bad structures, these structures are highly combustible, if it looks awkward, it is probably because it is wrong. 3/28/16 Energyof bonds∈products (Energyof thebonds∈thereacta)t∑ (¿) Δ Hrxn∑ ¿ OR BondsFormed (Bonds Broke) ∑ (¿) ΔH rxn∑ ¿  Breaking bonds takes energy and is endothermic always O  2 is a double bonded compound Draw the Lewis structure for: H O 2 2 But you aren’t supposed to make oxygen-oxygen bonds, but we have no choice in this case, so its okay. −¿ BF4 The Rules for drawing Lewis Structures 1. Skeletal structures(Weakest electronegative in the center ,ie Boron, then the others around it) 2. Valence electron count (count the total number of valence electrons) 3. Octets (8 electrons for each element) 4. Multiple bonds (if you need a double or a triple to complete the oc❑et, do it) Try CS 2  How many valence electrons:16  Which one is in the center, the least electronegative, they are both 2.5, put the single one in the middle so Carbon  Fill in the electrons……but we don’t have enough electrons for carbon to have an octet. Make some double bonds! Final result: Now do COCl 2  Center atom?  Number of Valence Electrons?  Are there enough electrons to satisfy the central Atom?  Yes? Cool you’re done, No, what bonds would make it work? Remember, Oxygen really likes double bonds.  You must use lone pair electrons from the outer atoms to make the double bonds, because the point of double/triple bonds is to satisfy the central atom FORMAL CHARGE  Total the electrons around the atom(bonding electrons are split evenly)  Subtract this from an atoms normal charge(from the periodic table)  Formal Charge is used to decide between different possible structures  The best structures have a 0 formal charge Example: CS 2 Current Charge-Charge found on the periodic table=Formal Charge 6-6=0 4-4=0 6-6=0 3/30/16  The formal charge must add up to the charge of the ion. So in −¿¿ BF 4 the formal charge must be -1 because the charge on the ion is -1 −¿ Draw C H COO ¿ 3  Oxygen-oxygen bonds are bad, so a carbon should be in the center  24 valence electrons  the first carbon has three hydrogens,and one carbon, octet complete  the second carbon has 1 carbon and two oxygens, so a double bond is need to complete the octet  Having resonance structures makes a compound more stable Resonance (IT’S IMPORTANT)  They show the possible electron arrangements of a given ion  They really show the extreme. An electron that makes a double bond is actually switching back and forth. That’s hard to show in a picture so we have resonance structures  Must show resonance structures that have the optimal formal charge  There are structures that have non-optimal formal charge (necessary for later Chem classes)  IF MORE THAN ONE RESSONANCE STRUCTURE EXIXTS, ALL EQUIVALENT (BY FORMAL CHARGE) RESONANCE STRUCTURES MUST BE SHOWN LINKED BY DOUBLED HEADED ARROWS  In oxaoacids, “ionizable hydrogen” is not attached to the central atom. Resonance structures for HCN: How many resonance structures are there, 1. There is never 0 resonance structures. Every possible combination counts as a resonance structure Sulfite ion: Sulfur is special, it can actually have extra electrons. So instead of the formal charges being wrong for the sulfite ion, they can be correct 1 Bond order for S-O is 1 3 Resonance Hybrid  The average of the resonance structures, an attempt to show the molecules as it really exists How many Resonance structures does Nitric Acid have? 1. What’s the formula for Nitric Acid (HNO )3 2. What’s the central atom?(N) 3. Satisfy the outer atoms with octets first 4. Is the central atom satisfied? 5. Satisfy it 6. How many resonance structures are there?(2) Bond order 3/2, 3 3 electrons spread over 2 places  Only Carbon, Oxygen, Nitrogen, and Fluorine must obey the octet rule!  The other elements tend to obey the octet rule, but they don’t have too  Any Atom int the third period or lower can have more than eight valence electrons(up to 18) Example PCl 5  Has 40 electrons  No way to draw this without having more than 8 electrons on the central atom SF 6  48 electrons Odd numbers of electrons:  Compounds/ions with unpaired electrons are called radicals NO  11 electrons 


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