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Week 2 notes

by: Katheran Mccarroll

Week 2 notes Chem 162

Katheran Mccarroll
GPA 3.77
General Chemistry
Colleen Craig

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About this Document

This is my notes for 10/9/15 through 10/14/15 it includes the last half of lecture 4, and all of lectures 5,6, and 7 (including the video she posted online). I hope these notes are helpful good lu...
General Chemistry
Colleen Craig
Class Notes
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This 3 page Class Notes was uploaded by Katheran Mccarroll on Wednesday October 14, 2015. The Class Notes belongs to Chem 162 at University of Washington taught by Colleen Craig in Fall 2015. Since its upload, it has received 26 views. For similar materials see General Chemistry in Chemistry at University of Washington.


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Date Created: 10/14/15
Chemistry 162 Week 1 Notes Lectures 123 and first half oflecture 4 93015 through 10715 Types of chemical bonds Ionic bonds involve the quottransferquot of electrons from one atom to another and usually take place between a metal and a nonmetal o The word quottransferquot is in quotbecause in reality there is never a 100 transfer of the electrons between the atoms but conceptually it is easier for us to think of it this way 0 Melted ionic solids conduct electricity I The proves that there are charged species present 0 Ionic bonds can all technically be considered polar covalent bonds with varying degrees of ionic character dipole moment X Yexpiriment 39 Ionic character xloooO dipow moment XYcalculated Nonpolar covalent bonding occurs when two atoms with very similar electronegativity bond to each other most commonly two of the same atoms bond together like a carboncarbon bond or a C1C1 bond 0 Important note CarbonHydrogen bonds are essentially non polar even though they are two different atoms Polar covalent bonds occur when there is unequal sharing of electrons in the bonding pair 0 The most common example of this is the HydrogenFluorine bond Fluorine is so much more electronegative than hydrogen that it pulls Hydrogen s electron closer to it I This means there is a build up of negative energy around F and a buildup of positive energy around H 0 When one of these molecules is in an electric field they usually line up showing there charge 0 This separation of charge is called the dipole gt H F I This is what the dipole moment of HF looks like the sign is over the H atom to show its slightly positive charge and the arrow points towards the F atom to show the build up of electrons around it I The equation for the dipole moment is M QR 1 is the dipole moment Q is the charge magnitude R is the separation distance It is given in units of D which is coulumbsmeter I Some molecules have no net dipole moment because there are two equal dipoles moving in opposite directions An example of this is C02 Electronegativity is the ability of an atom to attract shared electrons to itself 0 Electronegativity generally increases as you move up a group and across a period to the right H 2 More electronegative Li Be B C mm m WW um my mm 39Lanlhamde series 57 53 59 5 639 62 3 5 be mean hng snap quotlgg ll wmn La Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Actinideseries Ac Th Pa U Np Pu Am Cm Bk Cf Es Md No o Ionization energy is higher when electronegativity is high 0 The greater the difference in electronegativity is the more ionic the bondis Lewis structures Atoms usually form as many bonds as they have holes in their Lewis dot structure 0 For example 0 usually makes two bonds F usually only makes one bond N usually makes 3 bonds and C usually makes 4 bonds Steps to drawing a Lewis dot structure 0 Sum up the total valence electrons for all of the atoms 0 Use pairs of electrons to make the bonding pairs and subtract the number of electrons used from the total sum you just calculated 0 Distribute the remaining valence electrons subtracting as you go so that you know you used the correct amount Special rules to remember 0 H and F can both only form one bond 0 O and H are usually bonded to each other 0 C atoms usually bond to each other to form a chain or loop this is what the study of organic chemistry is all about 0 C atoms REALLY do not like to have lone pairs they like to satisfy the octet rule by making bonds Resonance Structures 0 Resonance structures are when a molecule can form multiple correct Lewis dot structures 0 The real structure is an average of all the available resonance structures but that cannot be easily diagramed I An example of this is ozone 03 I o oo I Both of the bonds in ozone are somewhere in between a double and a single bond this is proven by bond lengths Structural Isomers 0 Formal charge is the apparent charge on every atom in a molecules Lewis Dot Structure LDS I Minimal formal Charge on ALL ATOMS tends to be te most correct LDS o Assigning Formal Charge I Lone pairs are assigned to the atom in question I Bonding pairs are split evenly between the two bonded atoms I The sums of formal charge for all atoms must equal the overall charge of the molecule or species Valence Shell Expansions 0 Period 3 elements can have more than 8 valence electrons by starting to fill the 3d orbital o Atoms that usually form valence shell expansions are P and S Sometimes Cl Br and I will as well 0 The atom with the expanded shell should be the central atom in the LDS Sub octet systems 0 These form stable bonds with less than 8 valence electrons o B and Be commonly do this 0 BF3 and Ber are the most common example


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