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65 Cards in this Set
- Front
- Back
Mass number and Atomic number of an element
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quantum #s: electron address
n l m(l) m(s) |
n: 1...infinity. aka energy level
l: 0...(n-1) subshell (e.g. s,p,d,f) m(l): -l....+l orientation m(s): +1/2 or -1/2 up or down |
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Quantum Numbers
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Electron Configuration exceptions
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those elements located at d4 and d9 want to make orbitals half-full or full.
e.g. Cr [Ar]4s1 3d5 Cr, Cu, Mo, Ag, Au |
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Paramagnetic
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when there is unpaired electrons
odd # e- always paramagnetic attracted to magnetic field |
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Diamagnetic
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when electrons are paired at each orbital
very slight repulsion to magnetic field even # e-....either para or diamagnetic. |
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How many electrons per shell?
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2n^2 e-/shell
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alpha decay
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alpha = helium
only Z>83 |
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beta- decay
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n to p...because n/p >> 1
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beta+ (positron) decay
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p to n...N/Z << 1
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Electron capture
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p to n...N/Z << 1
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gamma decay
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mass defect
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mass is less than the sum. Iron 56 has highest Nuc binding energy.
E/nucleon is proportional to ∆mass/nucleon |
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Unstable Nuclei
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Odd # of protons and/or neutrons
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Stable Nuclei
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N/Z ~ 1 and Z < 20
magic #s for N or Z: 2, 8, 20, 50, 82, 126 |
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Effective Nuclear Charge (Zeff)
Zeff = group # |
Generally speaking, effective nuc. charge is charge felt by the valence electrons after you have taken into account the number of shielding electrons that surround the nucleus.
Example: Lithium, three protons and electron config. 1s2 2s1. The e- in 2s orbital is shielded from full attraction of the protons by the e- in the 1s orbital. Thus, Z* felt by the 2s electron should be one rather than three. |
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Atomic Radius
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down and to the left.
Na > Na+ Cl- > Cl |
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Ionization Energy
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up and right.
energy required to remove 1e-. 2nd e- is much harder to remove. (Endothermic) d5 and d10 Elements are more stable. |
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Electron Affinity
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increase from left to right.
Energy change associated with gaining an e- (Exothermic) |
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Electronegativity
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up and right.
smaller radius = higher e- neg. |
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Bonding
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Ionic
(strongest, meta/nonmetal, high m.p., high b.p.) Metallic (metal/metal, malleable, ductile) Covalent (nonmetal/nonmetal, lower m.p, b.p.) |
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Covalent
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sharing pairs of e-
molecular: h2O network solid: diamond, graphite |
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Formal Charge
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Normal # V.E. - Actual # V.E.
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Intermolecular Forces
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Hydrogen Bonding (strongest of Inter's)
Dipoles London Dispersion/Van der Waals |
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Branching of alkanes
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Sandwiches
decreases b.p. increases m.p. |
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Longer Alkane (effects on b.p.)
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longer carbon chain increases SA, greater london dispersion force
therefore increases b.p. |
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Phase Diagram
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Phase Diagram for CO2
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1atm is lower than the triple point. This is why it sublimes (solid to gas)
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Phase Diagram for H2O
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This explains why ice floats (ice is less dense than liquid water).
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Calorimetry
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q = mc∆T
q = heat change m = mass c = specific heat ∆T = Tf - Ti |
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Phase Transitions
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s-l-g (H>0, S>0)
fusion, sublimation, vaporization g-l-s (H<0, S<0) condensation, deposition, crystallization |
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Ideal Gas
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Each gas molecule has no volume.
No intermolecular forces. All collisions are elastic. KE ∝ T 1atm = 760 torr = 760 mmHg STP = 0∘C, 1atm 1 mole gas = 22.4L |
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Ideal Gas Law
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PV = nRT
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Dalton's Law of Partial Pressure
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P1 = X1 Ptot
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Graham's Law of Effustion
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Rate is proportional to velocity, lighter gas has faster veloctiy
M = molar mass |
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Combined Gas Law
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P ∝ 1/V
P ∝ T V ∝ T |
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Solubility Rules
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1) All Group I metal, NH4+, and NO3- salts are soluble
2) Most Ag+, Pb2+, and Hg₂2+ salts are insoluble |
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Freezing Pt. Depression
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m is number of ions, Ba(OH)2 has 3 ions when dissociated
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Vapor Pressure Depression (Raoult's Law)
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Boiling pt. Elevation
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Boiling pt. Elevation
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Osmotic Pressure
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water moves across membrane from (low solute) high water conc. to low conc.
pie = MRT or P = (n/V)RT |
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Catalyst
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Speeds up reaction by lowering Ea.
Does not get used up. starts as reactant and ends up as product. (an intermediate is the reverse) |
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Rxn Coordinate Diagram
Identify Transition State, Intermediates, delta H |
delta H < 0 because final is less than initial
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Rate Law
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only reactants in rate law.
coefficients do not determine order. x, y exponents indicate order |
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Common Ion Effect
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increasing concentration of ion will decrease solubility
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Le Chatelier's Principle
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Add something to a rxn, the rxn will adjust to regain equilibrium.
In Exothermic (T is product), increasing T will cause a shift to LEFT increase in P = decrease in V increase V then shift to the side with more moles of gas |
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Strong Acids
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perchloric, chloric, hydrochloric....
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Strong Bases
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Binary Acid Trend (anything with H)
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more e.n. the atom, stronger pull on H and makes it easier to give up H.
bigger the atom, longer and weaker the bond, therefore, easier to give up. |
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OxyAcid Trend
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everything else being equal, more Oxygen, more acidic.
more e.n., more acidic |
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[H+] of WEAK acids
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this equation is only for weak acids
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3 LAWS of Thermodynamics
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1. Conservation of Energy
2. ∆Suniv. > 0 3. S = 0 at 0K |
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Systems: open, closed, isolated
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open: exchange of heat, work, matter
closed: exchange of heat and work isolated: no exchange |
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Neutral Cations
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Conjugates of strong bases e.g. group 1A metals
all other cations are acidic |
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Neutral Anions/Bases
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Conjugates of strong acids e.g. Cl-, Br-, I-, NO3-, ClO4-, ClO3-
1 strong conjugate: HSO4- almost all other anions are bases |
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Henderson-Hasselbalch equation (or buffer equation)
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Bond Enthalpies
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∆H = ∑Dbroken - ∑Dformed
D = bond bond breaking is endo bond making is exo (think alphabetical b to endo, m to exo) |
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Enthalpies of Formation (Hess's Law)
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products or reactants in elemental state have ZERO formation values. e.g. Cl2, O2 etc.
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∆G = ∆H - T∆S
(Get Higher Test Scores) |
remember trends:
postive ∆S helps to go spont. negative ∆H helps to go spont. T amplifies S |
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Electrochemical Cells (for all cells)
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AN-ode OX-idation
RED-uction CAT-hode e- always flow from anode to cathode (because e- are lost at anode, gained at cathode, cathode gains mass) salt bridge: anions flow to anode, cations to cathode |
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Galvanic/Voltaic Cell
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cathode +
anode - |
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Electrolytic Cell
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cathode -
anode + consumes energy: e- goes to neg. cathode (not spontaneous) |
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Oxidation States
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1) Elements in their elemental form are in the zero oxidation state.
2) Group 1 metals are +1 and Group 2 metals are +2. 3) Hydrogen is +1 except when bonded to metals (when it’s –1). 4) The most electronegative elements get their typical ox state. 5) The last element not assigned balances the charge of the compound/ion. |
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Spontaneous Rxn's
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∆G < 0
Ecell > 0 Q < K |