chem 103 week 12
chem 103 week 12 CHEM 103 - 002
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CHEM 103 - 002
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This 4 page Class Notes was uploaded by Karlee Nelsen on Friday November 20, 2015. The Class Notes belongs to CHEM 103 - 002 at University of Wisconsin - Eau Claire taught by Sanchita Hati in Fall 2015. Since its upload, it has received 36 views. For similar materials see General Chemistry I in Chemistry at University of Wisconsin - Eau Claire.
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Date Created: 11/20/15
For K(19) the nobel gas configuration is [Ar] s s and any other electrons that would come after the noble gas are valence electrons which are easily removed Effective nuclear charge, EA, and IE increase while atomic side decreases Effective nuclear Periodic table charge, IE, and EA decrease. Atomic size decrease Electron Affinity [EA] – energy change for the process of adding an electron to a neutral atom in the gaseous state to form a negative ion - Electron added = energy release EA is higher according to magnitude of the change of Energy Cl (-349) > I (-295.2) - Noble gasses have NO EA values - Often exothermic Chemical Bonds - Bonding occurs with valence electrons - Electrons are either valence or core electrons - Atoms want to have 8 electrons Max hold 2 X electrons - Lewis symbol (lewis dot structure) - Ions need [brackets] and a charge 1 Li(3) [He] 2s 2 Li C(6) [He} 2s 2p C Octet Rule – G.N. Lewis – when atoms bond, they gain/loss/share electrons (due to the presence of d – orbitals, elements in period 3 and onwards exhibit expansion of the octet rule (they can bond more times O C O Types of Bonds Ionic Bonds – electrons transferred Covalent – electron sharing Metallic – mobile valence electrons Isoelectronic – chemical species (atom, ion, molecules) having the same number of electrons. - + - Ex. Ne(10) 10e and Na 10e F Lewis Structures Cl CC4 C Cl 4 valence electrons 7 valence electrons per atom Cl C Cl Total = 4 Total = 7 x 4 = 28 Cl Total valence electrons for molecule are 32 Each line represents two electrons and each dot is one electron Phosphate ion 3- PO4 3- O O P O O Ch 7 book notes Pauli exlusion principle – no two electrons can occupy the same orbital. They cannot have the same quantum numbers. n, l, m, m l s Subshell filling order. Lower (n+l) fills first, if two n+l equal the same number then the lower n level fills first ZeffEffective Nuclear Charge – net charge experienced by a particular electron in a multielectron atom resulting from a balance of the attractive force of the nucleus abnd the repulsive forces of other electrons Zeffs<p<d<f Hund’s Rule – the most stable arrangement of electrons in a subshell is that with the maximum number of unpaired electrons, all with the same spin direction Ions are formed by gaining/losing electrons in the atom’s effort to reach noble gas configuration and thus be stable - In transition metals the s electrons are lost first Paramagnetic – elements and compounds that have unpaired electrons are attracted to a magnet Diamagnetic – all orbitals are filled and paired and a slight repulsion occurs when exposed to a magnet IE – ionizing energy – is energy required to remove an electron from a gaseous atom Increases as electrons are removed Electron attachment enthalpy - ∆ H– EAthalpy change occurring when a gaseous atom adds an electron to form an ion Closely related to EA EA – electron affinity – equal magnitude, opposite sign of the internal energy change associated with gaseous atom adding and electron Cation atom size is smaller than ion size because more charge is acting on the remaining electrons so it can pull them closer. Anions are larger for the same reason as well as electron-electron repulsion
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