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65 Cards in this Set
- Front
- Back
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when the elements were discovered |
metals are ancient elements (Gold, silver, ect.)
1735-1843 explosion/expansion period (p-block & transition metals 1923-1961 lanthanides/ actinides (world wars & cold war for nuclear purposes) elements 104-118 are man made (half-lives...what good does it do?) |
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Schrodinger's wave equations quantum numbers (n, l, ml, ms ) Pauli exclusion principle shell subshell orbital |
-existence & energy of electron in atom is described by its unique wave function, psi -no 2 electrons in an atom can have the same four quantum numbers (each stadium seat is uniquely identified (E, R12, S8); each seat can only hold one individual at a time) -electrons w/the same value of n -electrons w/the same values of n and l -electrons w/the same values of n, l and ml |
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Bohr's Mistake -----INCORRECT energy of orbitals in a single electron atom |
energy only depends on principal quantum number n En = -RH ( 1/ n^2 ) |
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REALITY energy of orbitals in a multi-electron atom |
energy depends on n AND l 1s n=1, l=0 2s n=2, l=0 2p n=2, l=1 3s n=3, l=0 3p n=3, l=1 4s n=4, l=0 3d n=3, l=2 4p n=4, l=0 5s n=5, l=0 4d n=4, l=2 |
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"fill up" electrons in lowest energy orbitals first (Aufbau principle) of same energy = degenerate list orbitals in increasing energy? higher the orbital energy, the further away the electrons are from the nucleus |
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d.... |
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Hund's Rule |
the most stable arrangement of electrons in subshells is the one with the greatest number of parallel spins |
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paramagnetic diamagnetic |
-form magnetic moment; unpaired electrons that all face the same way -all quantum numbers are being used in the subshell; all electrons paired |
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shorthand for orbital filling by electrons... |
follow atomic number on the periodic table |
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effective nuclear charge ( Zeff) |
-the "positive charge" felt by an electron Zeff = z - shielding constant Zeff is about Z, the number of inner/core electrons Z Core Zeff Radius (pm) Na 11 10 1 186 Mg 12 10 2 160 Al 13 10 3 143 Si 14 10 4 132 |
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General trend of effective nuclear charge |
increases from right to left increases from top to bottom the stronger the nucleus, the smaller the radius radius decreases as you go from left to right |
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Electron configuration ___________ [ up arrow ] --------------- Orbital diagram: H 1s^1 |
-how the electrons are distributed among the various atomic orbitals in an atom 1s^1 -the front number is the principal quantum number -letter is the angular momentum quantum number l -the exponent is the number of electrons in the orbital or subshell -all orbitals are allowed 2 electrons |
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outermost subshell is the part being filled with electrons |
s subshell 2 electrons p subshell 6 electrons d subshell 10 electrons f subshell 14 electrons |
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simplify electron configurations by using the Noble gas rule; if you pass a noble gas, use it in brackets give full electron configuration for Phosphorus and Potassium use the noble gas rule for Argon and calcium |
P --> 1s^2 2s^2 2p^6 3s^2 2p^3 K --> 1s^2 2s^2 2p^6 3s^2 2p^6 4s^1 Ar --> [Ar] Ca --> [Ar] 4s^2 |
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orbital energy stability due to electronic interactions of electrons, there is a hierarchy of stability.... odd electron configurations: Anomalous Series contain what elements? electron configuration of copper? |
Stability: stable to unstable empty orbitals --> full orbital --> 1/2 filled (Hund; spins in same direction) --> partially filled anomalous series: Cu, Au, Ag, Pd, Cr, Mo [mostly in 3d or 4d orbitals] [6B and 1Bs do this very often] -move 1 or 2 electrons to stabilize the rest Cu: [Ar] 4s^1 3d^10 b/c it's more stable!! |
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ground state electron configurations sulfur? palladium, which is diamagnetic? |
S --> 1s^2 2s^2 2p^6 3s^2 3p^4 paramagnetic Pd --> [Kr] 4d^10 |
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Classification of the elements -same electron state means that those elements act similar |
representative elements are 1A through 7A noble gases are 8A transition metals are 3B through 1B zinc, cadmium and mercury are 2B (noble gases of transition metals) lanthanides 4f actinides 5f |
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a certain element has 15 electrons... ground-state configuration? how should it be classified? diamagnetic or paramagnetic? |
[Ne] 3s^2 3p^3 representative element paramagnetic |
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electron configuration of cations and anions of representative elements -valence electrons have highest principal quantum numbers and any unfilled subshell |
-atoms lose electrons so that cation has a noble gas outer electron configuration ex. Na, Al, Ca -atoms gain electrons so that anion has a noble gas outer electron configuration ex. H, F, O, N |
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isoelectronic species isoelectronic: |
-have the same number of electrons & hence the same ground state electron configuration -atoms are not identical & don't have the same properties -usually revert to noble gas form to be stable ex. Na+, Al 3+, F-, O 2-, N 3- |
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electron configurations of cations of transition metals Fe3+ and Mn2+ are isoelectronic |
-when a cation is formed from an atom of a transition metal, electrons are always removed first from the ns orbital, and then from the (n-1)d orbitals RULE OF THUMB: -add electrons using the atomic number/trend of periodic table; remove them from the highest principle quantum number FIRST! |
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how atomic radii are determined spheres = electron cloud probability to find electron metallic radius & covalent radius |
-measure from nucleus of one atom to another & cut it in half to get atomic radii |
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atomic radius trends exceptions: |
-increases from right to left (increase of effective nuclear charge in opp. direction) -increases from top to bottom (increasing principal quantum number) -Ga gets slightly smaller going down in p-block that has 3d contributions before it -Sb to Te show up as paramagnetic -Pb to Bi --4th electron destabilizes & gets bigger so atomic radii gets bigger to accommodate it -transition metals, lanthanides & actinides don't follow atomic radii rules (trench U shape) on left of U 5 electrons make them smaller (paramagnetic), one right of U 5 electrons make radii get bigger (pair up, diamagnetic) |
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list in order of increasing atomic radius: P, Si, N |
N < P < Si |
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comparison of atomic radii with Ionic radii |
ionic radii: -cations radii are smaller b/c it reverts back to noble gas configuration ( + charge ion, radii gets smaller) -anions radii are larger b/c decreasing effective nuclear charge (- charge ion, radii gets larger) F -->9protons & 9electrons; F- 9protons & 10e- |
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ionic compounds |
-cation is always smaller than atom from which it is formed -anion is always larger than atom from which it is formed Fe > Fe 2+ > Fe 3+ -left to right rule no longer works across the periodic table |
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which is larger? N3- or F- Mg 2+ or Ca 2+ Fe 2+ or Fe 3+ |
N3- b/c more e-'s to p+'s, decrease effective charge Ca 2+ b/c higher principal quantum number Fe 2+ b/c more p+'s (increase effective charge) |
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ionization energy I1+ X (g) ---> X+(g) + e- I1 first ionization energy I2 + X+(g) ---> X2+(g) + e- I2 second ionization E. I3 + X2+(g)--->X3+(g) + e- I3 third ionization E. I1 < I2 < I3 |
-the minimum energy (kJ/mol) required to remove an electron from a gaseous atom in its ground state *always remove electrons -smaller values, easier for e- to leave -larger values, harder for e- to leave *must have gas part of element *electrons adopt lowest quantum energy level -from first to second to third ionization energy, the kJ/mol required increases, as well as the effective nuclear charge -electrons become harder and harder to remove |
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general trends in first ionization energies |
-as you increase principal quantum number in a group (1A) it gets easier to remove electron -Be to B, boron is easier b/c changing block orbital than from same block orbital (take a p-block away instead of s-block orbital electron) -O to N, oxygen is easier b/c there is an opposing spin e- that wants to leave -ionization energy is all endothermic -noble gases have highest ionization energy b/c they are stable -have not done ionization energy for lanthanides & actinides -increases from left to right -increases from bottom to top |
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smaller first ionization energy? O or S higher second ionization energy? Li or Be |
sulfur b/c further away e- is easier to remove (principal quantum number) Lithium b/c should be easier to remove 2s e- from Be+ than to remove a 1s e- from Li+ |
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electron affinity X(g) + e- ---> X-(g) F(g) + e- ---> F-(g) deltaH = -328 kJ/mol EA= +328 O(g) + e- ---> O-(g) deltaH = -141kJ/mol EA= +141 |
-the negative of the energy change that occurs when an electron is accepted by an atom in the gaseous state to form an anion *exothermic but EA is positive -zero value means upset to gain e- -from left to right EA increases -N doesn't want one b/c it's paramagnetic (all same spins) -F to Cl exception b/c Cl wants e- more than F due to atomic size -noble gases don't want to be messed with -alkaline earth metals don't want e- b/c it doesn't want to fill another shell (negative or sm. values) |
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elemental relationships |
diagonal relationships: -Li acts like Mg b/c charge densities act very similar 1A elements: -reactive metals, explode in water, more reactive as you go down 2A elements: -less reactive metals, Rd is radioactive outlier 3A elements: -redox elements, reduced or oxidized, Ga can melt in your hand 7A elements: -reactivity increases as it goes up 8A elements: -fairly unreactive; completely filled ns and np subshells; highest ionization energy of all elements; no tendency to accept extra electrons three common halogens: Cl2 Br2 and I2 |
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valence electrons |
the outer shell electrons of an atom - valance electrons are the electrons that participate in chemical bonding Group: # of valence e- 1A 1 2A 2 3A 3 4A 4 5A 5 6A 6 7A 7 |
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ionic bond Li ---> Li+ + e- (ionization energy) endothermic e- + F ----> F- (electron affinity) exothermic creates stable compound through ion combination |
the electrostatic force that holds ions together in an ionic compound -very strong -both are happy -give up or gain e-, then pair up -reach noble gas notation before combining |
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use lewis dot symbols to show the formation of aluminum oxide Al2O3 |
both reach [Ne] electron configuration && then bind Al3+ & O2- |
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lattice energy (U) E = [ k (Q+)(Q-) ] / r E is the potential energy Q+ is the charge on the cation Q- is the charge on the anion r is the distance between the ions |
-the energy required to completely separate one mole of a solid ionic compound into gaseous ions -lattice energy increases as Q increases and/or r decreases Compound Lattice Energy (KJ/mol) MgF2 2957 Q: +2, -1 MgO 3938 Q: +2, -2 LiF same 1036 LiCl charge 853 r F- < r Cl- 1. difference in charge takes preference 2. atomic size (bigger radius means energy gets smaller [lower lattice energy]*use when charges are the same) |
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covalent bond |
-a chemical bond in which two or more electrons are shared by two atoms (no charges) -why share electrons? both electrons are happier; share energy (takes less energy/effort) -lone pairs are associated to only 1 atom -single covalent bond (shared pair, covalent pair, bonding pair) |
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single covalent bonds: ex. water double bond (two atoms share two pairs of electrons) ex. CO2 triple bond (two atoms share three pairs of electrons) ex. N2 |
lewis structures of each example? |
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lengths of covalent bonds bond lengths: triple bond < double bond < single bond bond stability: single bond < double bond < triple bond |
-bond length indicates bond strength -smaller bond length btw 2 atoms, the stronger it is O-H 96 pm bond length (makes it very strong b/c of polar covalency) |
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ionic vs. covalent |
-solids are mostly ionic -liquids and gases are mostly covalent -ionic melting points, molar heat of fusion, boiling point, molar heat of vaporization, density, and solubility in water are much higher than covalent compounds -both have poor electrical conductivity as solids, but ionic compounds are good conductors as a liquid |
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writing lewis structures |
1. draw skeletal structure of compound showing what atoms are bonded to each other; put least electronegative element in the center 2.count total number of valence e-; add 1 for each negative charge; subtract 1 for each positive charge 3.complete an octet for all atoms except hydrogen 4. if structure contains too many electrons, from double and triple bonds on central atom as needed |
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lewis structures: NF3 HNO3 CO3 2- |
(: check yo answers |
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there are two possible structures of formaldehyde (CH2O) formal charge? -the sum of the formal charges of the atoms in a molecule or ion must equal the charge on the molecule or ion |
-the difference between the number of valence electrons in an isolated atom and the number of electrons assigned to that atom in a Lewis structure formal charge equation: formal charge on an atom in a Lewis structure = total number of valence electrons in the free atom ---minus--- total number of nonbonding electrons ---minus--- 1/2 (total number of bonding electrons) |
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find the correct lewis structure of the formaldehyde compound |
*the species with no formal charge is the stable molecule |
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formal charge & lewis structure rules |
1. for neutral molecules, a lewis structure in which there are no formal charges is preferable to one in which formal charges are present 2. lewis structures with large formal charges are less plausible than those with small formal charges 3. among lewis structures having similar distributions of formal charges, the most plausible structure is the one in which negative formal charges are placed on the more electronegative atoms |
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formal charge of carbonate ion? draw lewis structure of CH2O |
-2 |
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resonance structure |
-one of two or more Lewis structures for a single molecule that cannot be represented accurately by only one Lewis structure ex. ozone; N2O |
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incomplete octet rule exception odd-electron molecules ex. NO the expanded octet rule exception |
-less than 8 electrons needed to reduce formal charges ex. BeH2 , BF3 -central atom with principal quantum number n>2) ex. SF6 |
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draw lewis structure AlI3 PF5 |
SO4 2- XeF4 |
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polar covalent bond/ polar bond |
-a covalent bond with greater density around one of the two atoms (unfair sharing) ex. HF -more red= more negative electron density -more blue= more positive |
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electronegatively |
-the ability of an atom to attract toward itself the electrons in a chemical bond Electron Affinity --measurable, Cl is highest X(g) + e- -----> X- (g) Electronegativity --relative, F is highest -electronegativity increases going bottom to top -electronegativity increases going from left to right -increased nuclear charge, increases electronegativity |
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Difference Bond Type 0 to .5 covalent greater than or equal to 2 ionic between .5 and 2 polar covalent |
covalent ex: H2, Cl2 (same elements bonded together) ionic: unequal sharing ex: LiF, KBr polar covalent: equal sharing ordered in increasing diff. in electronegativity covalent ----> polar covalent ----> ionic share e- ---> partial transfer of e- -->transfer e- |
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classify the following bonds as ionic, polar covalent, or covalent HCl KF the CC bond in H3CCH3 |
-polar covalent .9 -ionic 3.2 -covalent .3 |
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Valence shell electron pair repulsion (VSEPR) model |
predict the geometry of the molecule from the electrostatic repulsions btw the electron (bonding and nonbonding) pairs -different central atoms will have different molecular geometry -if there are no lone pairs, the molecular geometry is the same as the electron geometry |
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Class: AB2 # of atoms bonded to central atom: 2 # lone pairs on central atom: 0 arrangement of electron pairs: :-A-: 180 degrees; linear; 2-D motif molecular geometry: linear B-A-B 180 degrees ex. BeCl2 |
class: AB3 # of atoms bonded to central atom: 3 # lone pairs on central atom: 0 : arrangement of electron pairs: | trigonal planar; 120 degrees :-A-: B molecular geometry: | trigonal planar; 2-D motif B-A-B ex. BF3 -3 terminal atoms on central atom |
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class: AB4 (8 electrons around central atom) # of atoms bonded to central atom: 4 # lone pairs on central atom: 0 arrangement of electron pairs: tetrahedral 3-D 109.5 degrees : B | | :-A-: B-A-B | | : B molecular geometry: tetrahedral ex. methane (organic molecules) |
class: AB5 # of atoms bonded to central atom: 5 # lone pairs on central atom: 0 top& : bottom arrangement of electron pairs: : | 90 degree tigonal bipyramidal \ A-: middle atom 120 degree / | : : molecular geometry: trigonal bipyramidal ex. PCl5 (D orbitals) |
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class: AB6 -middle, equatorial atoms w/in the same plane -top&bottom axial positions (outside plane of central plane) # of atoms bonded to central atom: 6 # lone pairs on central atom: 0 : : : arrangement of electron pairs: \ | / octahedral; 90 degrees A / | \ : : : molecular geometry: octahedral ex. SF6 |
-for each lone pair, change degree angle by 2 to 3 degrees -more lone pairs, more they compress other atoms -know unaltered bond pairs, then change it based off that atom ex. CH4; all bonded --> 109.5 degrees NH3; one lone, rest bonded --> 107.3 degrees H2O; two lone, rest bonded --> 104.5 degrees lone-pair vs. lone-pair repulsion > lone-pair vs. bonding-pair repulsion > bonding-pair vs. bonding-pairs repulsion |
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class: AB2E # of atoms bonded to central atom: 2 # lone pairs on central atom: 1
arrangement of electron pairs: : trigonal planar B-A-B molecular geometry: bent 117-118 degrees ex. SO2 |
class: AB3E # of atoms bonded to central atom: 3 # lone pairs on central atom: 1 arrangement of electron pairs: : tetrahedral 107.5 degrees B-A-B solid triangle=plane towards you | dotted lines=plane away from you B molecular geometry: trigonal pyramidal ex. NH3 |
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Class: AB2E2 # of atoms bonded to central atom: 2 # lone pairs on central atom: 2 : : A arrangement of electron pairs: / \ tetrahedral B B 104-105 degrees molecular geometry: bent ex. H2O |
class: AB4E # of atoms bonded to central atoms: 4 # lone pairs on central atom: 1 B B arrangement of electron pairs: | / trigonal bipyramidal : A | \ 179-180 degree & 118-120 B B -take away equatorial position b/c it gives more space than axial position molecular geometry: distorted tetrahedral (seesaw) ex. SF4 |
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class: AB3E2 *take equatorial e- away first # of atoms bonded to central atom: 3 # lone pairs on central atom: 2 B arrangement of electron pairs: \ : trigonal bipyramidal B- A 177-178 degrees / : & 116-117 degrees B -leave as many atoms in plane as possible molecular geometry: T-shaped ex.ClF3 |
class: AB2E3 # of atoms bonded to central atom: 2 # lone pairs on central atom: 3 B arrangement of electron pairs: | : trigonal bipyramidal : A 180 degrees | : B molecular geometry: linear ex. I3- |
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class: AB5E *take away axial position 1st # of atoms bonded to central atoms: 5 # lone pairs on central atom: 1 B B B arrangement of electron pairs: \ | / octahedral A / : \ 87-88 degrees B B molecular geometry: square pyramidal ex. BrF5 |
class: AB4E2 *must take opposite of 1st e- away of 2nd # of atoms bonded to central atom: 4 # lone pairs on central atom: 2 B B arrangement of electron pairs: \ : / octahedral A / : \ 85-86 degrees B B molecular geometry: square planar ex. XeF4 |
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predicting molecular geometry *only care about central atoms |
1. draw lewis structure for molecule 2.count number of lone pairs on the central atom and number of atoms bonded to the central atom 3. use VSEPR to predict the geometry of the molecule |
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use VSEPR model to predict the geometry: AsH3 OF2 AlCl4- I3- C2H4 |
trigonal pyramidal geometry && tetrahedral tetrahedral && bent geometry tetrahedral && tetrahedral geometry trigonal bipyramidal && linear geometry trigonal planar |
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dipole moments & polar molecules |
u = Q x r Q = the charge r = the distance btw charges 1 D = 3.36 x 10^-30 Cm |
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how polar molecules react to magnetic fields |
field off: positive to negative ends occur, but aren't too organized field on: positive to negative ends occur && are extremely organized dipole moments in NH3 && NF3 ?? |
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predict whether each of the following molecules have a dipole moment: BrCl BF3 (trigonal planar) CH2Cl2 (tetrahedral) |
polar w/ Cl at the negative end; yes Fl more negative than B, but cancel out; no CL is more negative than carbon which is more negative than H; yes |
