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exam 2 chapters 8-11

by: nicole l brown

exam 2 chapters 8-11 chem 131

Marketplace > Truman State University > Chemistry > chem 131 > exam 2 chapters 8 11
nicole l brown
Truman State
Chemical Principles II
R Baughman

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About this Document

this set of notes has the chapters 8 and 10. 11 well be added once Dr. Baughman decides what in chapter 11 the exam well include.
Chemical Principles II
R Baughman
exam 2 notes
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This 9 page Bundle was uploaded by nicole l brown on Saturday October 3, 2015. The Bundle belongs to chem 131 at Truman State University taught by R Baughman in Summer 2015. Since its upload, it has received 70 views. For similar materials see Chemical Principles II in Chemistry at Truman State University.

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Date Created: 10/03/15
Chem 131 Dr Baughman Chem 131 chapter 810812 Multielectron atoms The electrons to these atoms are more spread out compared to the hydrogen atoms because of the multiply electrons bonds these atoms become intertwined and hard to determine each electron The angular parts are unchanged and the radial parts are different Hydrogenlike orbitals the electron orbitals are the same as those of the hydrogen atom Attractive force of nucleus for an electron increases and the nuclear charge increases orbital energy becomes more negative with the increase of atomic number of the atom Orbital energies depend on the type of orbital Screenings of the electrons are caused by the electrons in the orbitals closer to the nucleus shield the nucleus for the farther electrons This effect reduce the effectiveness of the nucleus attract the more distant electron reduce nuclear charge The types of orbitals in the inner electron are in and the type of orbital that is screened electron can decide the magnitude of the reduction of the nuclear charge S orbitals high probability density at the nucleus Electrons in this orbital are effective at screening the nucleus from the outer electrons o Penetration electrons that allow them to get close to nucleus Penetration is better at screening a low penetration P amp D orbitals zero probability P densities at nucleus The electrons are less effective than the s orbital Radial probability distribution probability of finding electrons anywhere in the spherical shell of radius e and the infinitesimal thickness Found by multiplying radial P density RA2 by factor 4 pie rquot2 area of a sphere of radius r Quaintly 4 pie rquot2 RA2 this provides a different view into the behavior of electrons For 1s orbital for radial P density the maximum P for 1s electron is at the nucleus H atom electron is likely to be found 53 pm from nucleus radial P distribution reaches the maximum You have a 95 P of finding an electron is a larger sphere probalitly one with a radius about 141 pm 1s orbitals increased P of being close to nucleus compared to 2s and 2p 2s electron increased chance of being close to the nucleus and a greater penetration than 2p Increased degree of penetration ca block the view of an electron in an outer orbital that is looking for the nucleus 0 Atomic number nuclear charge that electron would experience if there weren t any intervening electrons is Z 0 Effective nuclear charge is the nuclear charge that an electron doe s experience is reduced by the intervening electrons to the value of Zeff 0 Less of the nuclear charge that the outer electron can see smaller value of Zeff then the smaller the attraction of that electrons s to the nucleus so it a higher energy of that orbital in which the electron is found 0 Under the Zeff the effective nuclear charge has no more splitting of energies within a subshell because all the orbitals in that subshell have the same radial characteristics 0 P orbitals of that principal shell have same energy five D orbitals have same energy 0 Energy levels overlap because of the combined effect of the decreased spacing btw energy levels at the higher quantum s the splitting of subshell energy levels because of the shielding and penetration 0 Electron configurations o Is how the electrons are distributed btw the various orbitals in principal shells and subshells 0 Rules for assigning electrons to orbitals 0 Electron occupies orbitals in a way that minimizes the energy of the atom order goes by this Is 2s 2sp and so on How you fill these orbitals depends on principally through spectroscopy magnetic studies and experiment order based on experiment 0 1s 2s 2p 3s 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s 5f 6d 7p 0 No two electrons in an atom can have all four quantum numbers alikethe Pauli exclusion principle this concept comes from Wolfgang Pauli he explained how complex features of emission spectra that is associated with atom in magnetic fields by proposing that no two electrons in an atom can have all four quantum s alike First three quantum s are 11 3C and multitude factors determine the specific orbitals 2 electrons have to have 3 quantum numbers alike but they have to have different values of M4 spin quantum number So 2 electrons can occupy the same orbital but they have to have opposing spins So s subshell can get past the two electrons per orbital issue so s orbital can have 2 electrons but a P subshell has 3 orbitals with the total number it can have is 6 electrons 0 When orbitals if identical energy degenerate orbitals is available electrons initially occupy these orbitals singly also known as Hund s rule it says that atoms tend to have a lot of unpaired electrons Because the electrons carry the same electric charge and they get as far apart as possible they achieve this by finding empty orbitals of similar energy in preference to pairing up with and electron in a halffilled orbital 0 Different ways to show an electron configurations 0 Spdf notation condensed C 1squot2 2squot2 2pquot2 0 Spdf notations expanded C 1squot2 2squot2 2pxquot1 2pyquot1 o Orbital diagram C T V T T 2P Al 25 ls Spdf notation only shows the total number of electrons in each subshell Expanded shell shows Hund s rule Orbital diagram show each subshell in to single orbitals Arrow pointing up shows 12 Arrow pointing down shows l2 One arrow up and down shows opposing spins only 1s and 2s Aufbau process we assign electron configurations to elements in order of increasing atomic number To go form one atom to the next we as a proton neutrons to the nucleus Z1 H lowest energy state for electron is ls orbital 1squot1 22 He lsquot2 Z3 Li lsquot2 2squotl Z4 Be lsquot2 2squot2 Z5 B lsquot 2squot2 2pquot1 Z6 C lsquot2 2squot2 2pquot2 Z710 NNE the subshell are completed number of unpaired electrons reaches a maximum 3 with N and the decrease to 0 with Ne Z1 118 Na Ar 8 electrons and these go to 3s and 3p orbitals each of those element has ls 2s and 2p subshells filled Valence electrons those are the electrons that are added to the electron shell of the highest principal quantum number Z19 amp20 KampCa the next subshell to fill is a d orbital and that holds about 10 electrons There can be to ways to write this for scandium Sc Ar3dquot14squot2 or Sc Ar4squot23dquot1 There are two expects to these rules chromium and copper so 3d shell is halffilled with electrons Z3l36 GaKr 4p orbital is filled Z3754 RdXe the subshells fill 5s 4d 5p Z5586 CsRn the order is 6s 4f 5d 6p Z87 Fr 7s 5f 6d7p Electron configurations and the periodic table Elements in the same group of the table have similar electron configurations N the shell of the highest principal quantum number or outer most shell group 1 have one outershell electron in s orbital nsAl group 17 have 7 outershell elections nsquot2npquot5 group 18 have expect helium has 2 electrons 8 electrons nsquot2npquot6 0 how the periodic table is split up 0 S block it is the s orbital of the highest quantum number n this holds the groups of 1 and 2 plus He Example Sr Kr5squot2 0 P block it is the p orbital of highest quantum number n this holds the groups of 13 18 expect He Example Al Ne3squot23pquot 0 D block d orbitals the electron shell nl This holds the groups of 312 0 F block f orbital the electronic shell n2 This holds the elements lanthanides and actinides Chem 131 Dr Baughman Chapter 10 notes Lewis theory Lewis theory Electrons are a key component in chemical bonding especially the outermost valence electronic shell Electrons can be transferred from one atom to another The positive and negative ions are formed amp attract to each other because of the electrostatic forces called ionic bonds Electrons can love one or more pairs of electrons that are shared btw atoms Covalent bond is when the bond formed by the sharing of electrons btw atoms Electrons are shared in some sort of way and each atom acquires especially stable electron configuration Octet is when it has 8 outer shell electrons Lewis symbol is made up of the chemical symbol and the electrons of an atom This can have up to a maximum of 4 dots to a side Lewis structure this is the combined equation for the chemical bonds and the Lewis symbols To show that a bond is ionic it shows to show that a bond is covalent it shows a dot The use square brackets to show that it is an ion Binary ionic compounds is made out of monatomic cations and monatomic anions Ternary ionic compounds is made out of monatomic and polyatomic ions Bonding of the polyatomic ions is covalent Each cation is surrounded by anions and each of the anion has a cations Ionic crystal this is when large numbers of ions are arranged in a neat order network Covalent bonding The lower the ionization energy the more metallic an element for example Na is more metallic than H Covalent bonds are when two to atoms share and electron to fill that outer most valence shell Octet rule this is when a shell as a limit of 8 electrons to be more stable H is the expect of this rule with only 2 electrons Single covalent bond is when two atoms share a single pair of electrons Bond pair is the term used for a pair of electrons in covalent bond Lone pair this is the term used for a pair of electrons that are not in a bond Coordinate covalent bond it is a covalent bond with a single atom that shares both of the electrons in a shared pair Fyi once a bond is formed it is impossible to tell which electrons came from which atom Double covalent bond it is when two atoms share 2 electron pairs Triple covalent bond is the sharing of three pairs of electrons btw two atoms Multiple covalent bonds this is double and triple bonds Polar covalent bonds and electrostatic potential maps Ionic bonds are when the electrons go to another atom an atom steals an electron Polar covalent bond in a covalent bond the electrons are shared un equal btw the two atoms The electrons are pulled more to the nonmetallic element This type of sharing leads to a somewhat negative charge on the nonmetallic element and a positive charge to the element that is more metallic Electrostatic potential map it is a way to look at a picture and see the charge distribution within a molecule Electrostatic potential the work done in moving a unit of charge at a constant speed from one region to another charge wants a more electronrich region a region with excess charge after all of the charges of nuclei and electrons are accounted for then the electrostatic potential will be To have an electrostatic potential be then a point charge is placed in an electron poor region this is a region with excess charge and these well want to repelled Electronegativity EN is when an atom can compete for electrons with other atoms that it is bonded to Related to ionization energy I and electron affinity EA AB AABA DeltaE1 IaEAb AB gtAA1BA Delta E2 IbEAa Delta E1 delta E2 means the bonding of electrons are shared equally Nonpolar IaEAbIbEAa ENa delta IaEAa the atom is related to the electronegativity of the atom by this formula High ionization energy and electron affinity large and has a large electronegativity relative to an atom with a low ionization energy and small electron affinity According to Linus Pauling the lower EN then the more metallic the element is and the Higher EN it is more nonmetallic So looking at a periodic table the electronegativity decrease from top to bottom and it increase from left to right Electronegativity difference delta EN so for EN for two atoms are small then the bond btw they are covalent But is EN is large the then the bond is ionic For an intermediate causes of delta EN the bind is polar covalent Big EN differences are btw more metallic and more nonmetallic elements produce ionic binds Small EN differences are for two nonmetal atoms and bond btw they could be covalent Writing Lewis structures All electrons in the valence shell has to be shown Most electrons are paired 0 Most atoms have and octet in the outer shell H only has 2 outer shell 0 Sometimes a multiple covalent bonds can be double or triple bonds are needed 0 Skeletal structure Is when all of the atoms in the structure arranged in order which they are bonded to another atom 0 Central atom an element that is the center of the Lewis structure 0 Terminal atom Are elements that are on the outside of the central atom o H atoms are always a terminal atom because it can only have 2 electrons in its valence shell 0 The lowest electronegativity is most likely the central atom 0 C are always central atom 0 Molecules and polyatomic ions are most likely have compact symmetrical structures How to draw a Lewis structure 0 Find the total number of valence electrons that most appear in structure 0 Id the central atom and the terminal atom 0 Guess a skeletal structure use single covalent bonds 0 For each of those bonds 2 from the total of valence electrons 0 Complete the octets for terminal atoms then complete the ones for the central atoms 0 Then is there are incomplete octets us double triple bonds 0 Formal charges PC are the charges in the certain atoms from the Lewis structures because ether atoms have not contributed equal s of electrons to the covalent bonding 0 Count lonepair electrons these are the electrons that are found in an atom they didn t originate from 0 Divide bondpair electrons equally btw bonded atoms FC number valence e in free atom number lonepair e 12 number bondpair e 0 Some Lewis structures based off of formal charges 0 Sum of formal charges must 0 for a neutral molecule and have to magnitude of charge for polyatomic ion 0 When formal charges are need they should be as small as possible 0 formal charges appear on the most electronegative atoms on least electronegative atoms 0 The structures that have formal charges of the same sign on adjacent atoms are unlikely o Resonance o It is when you can write a Lewis structure different ways 0 The only thing you change btw structures is the bonds of the atoms or the electrons around the terminal and central atoms 0 Exceptions to the octet rule 0 NO has 11 valance electrons is paramagnetic 0 Free radicals highly reactive fragments with one or more unpaired electrons are highly reactive species 0 Incomplete octet is not having an enough of electrons we fix this problem by moving lonepairs electrons from terminal atoms to make multiple bonds This is true for beryllium boron aluminum compounds 0 Expanded valence shells they are 10 or even 12 valence electrons around that central atom These are usually nonmetal atoms of 3 periods they are bond highly electronegative atoms The only down side to this is you have to show were the extra electrons go 3s3p subshells from the central atom get 8 electrons in that shell and the extra electrons move to 3d subshell 0 Shapes of molecules 0 bond angles are the angles btw adjacent lines representing bonds 0 diatomic molecule one bond no bond angle they are linear 0 triatomic molecule 2 bonds and one bond angle 180 degree the atoms are in a straight line so linear other bonds the shape is angular or bent or v shaped 0 Valence shell electronpair repulsion theory VSEPR the pairs of electrons repel each other it doesn t matter if they are in the chemical bonds or unshared electrons and the pairs become orientations about an atom to minimize repulsions o A wedge shows that a bond points towards us and a dashed wedge shows the bond points away The symbols depend on the following 0 Clean lines show bond that lie in the plane of paper 0 The solid color wedges shoe bonds that pair towards the person reading 0 Dashed wedges show the bonds point away from reader behind the plane of paper 0 Electron group geometry the geometric arrangement of the atomic nuclei it is the actual determinant of molecular shapemolecular geometry 0 AX2E2 2 atoms or the groups are bonded to central atom A central atom 2 lone pairs of electrons E o Electrongroup geometries o 2 electron groups linear 0 3 electron groups trigonal planar 0 4 electrons groups tetrahedral 0 5 electrons group trigonal bipyramidal 0 6 electron groups octahedral 0 Molecular geometry electiongroup geometry but it is only when all electron groups are bond pairs 0 Stronger repulsion closer together 2 groups of electrons are forced stronger at 90 degree than 120 degree or 180 degree o Lonepair electrons that are spread out more than do bondpair electrons order of repulsive forces is from strongest to weakest repeal of 1 lone pair of e for another lone pair is greater 0 How to predict shapes of the molecules 0 Draw a Lewis structure 0 Determine number of E groups around the central atom and ID them as a bondpair electron groups pair lone pairs of electrons 0 Find out the electrongroup geometry around central atom linear trigonal planar tetrahedral trigonal bipyramidal octahedral 0 Find molecular geometry for the atoms that are bond directly to central atom 0 Polar molecule it shows using an arrow which atom attaches the electrons more 0 dipole moment 5 W M has value of 33410quot 30 coulombmeter Cm I has value of l debye This is based on the behavior of polar molecules in the electric eld


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