G Chem Flashcards
Energy of a photon (equation)
E = hf h = Planks constant = 6.63x10^-34 Js E = hv (v in Hz) E = cλ c = 3.00 x 10^8 m/s
EM Spectrum (order, E and f levels)
Gamma X-ray UV VIBYOR IR Microwave Radio
E high; f high, λ low E low; f low, λ high
Charge of e-
-1.6022 x 10^-19 C
Mass of proton
1.673 x 10^-27 kg
Mass e-
9.11 x 10^-31 kg
Relationship between c, λ, and v
c = λv
Principle quantum number
n
Angular quantum number
l (n-1) l = 0 s orbital l = 1 p orbital l = 2 d orbital l = 3 f orbital
Magnetic quantum number
ml
-l to +l
l = 0 ml =0 s orbital
l = 1 ml = -1, 0, +1 p orbital (py, px, pz)
Spin quantum number
ms
+1/2 spin up
-1/2 spin down
Common reducing agents (4)
LiAlH4
NaBH4
Metals
Fe2+
Common oxidizing agents (7)
Halogens Permanganate-(MnO4) salts Peroxide compounds (ie H2O2) Ozone (O3) Osmium tetroxide (OsO4) Nitric acid (HNO3) Nitrous oxide (N2O)
Lewis acid
e- pair acceptor
Lewis base
e- pair donor
Bronsted Lowry acid
H+ donor
Bronsted Lowry base
H+ acceptor
Electron and molecular geometry
Number lone pairs: 1
Number bonding pairs: 2
Trigonal planar (120), bent (sp2)
Electron and molecular geometry
Number lone pairs: 1
Number bonding pairs: 3
Tetrahedral (109.5), trigonal pyramidal (sp3)
Electron and molecular geometry
Number lone pairs: 2
Number bonding pairs: 2
Tetrahedral, bent (sp3)
Steric number (# hybridization orbitals in a molecule)
sigma bonds + # lone pairs
Electron and molecular geometry
Number lone pairs: 0
Number bonding pairs: 2
Linear (180)
Structural isomers
Same molecular formula, different compounds
Stereoisomers
Differ in 3-D arrangement (max = 2^n) n= # chiral centers
Enantiomers
Stereoisomers not superimposable mirror images
Diastereomers
Stereoisomers that are non-superimposable non-mirror images (differ at SOME chiral centers, not all)
Z configuration
Two top priority groups on same side of double bond
E configuration
Two top priority groups on opposite side of double bond
Sublimation
Solid to gas
Condensation
Gas to liquid
Freezing
Liquid to solid
Deposition
Gas to solid
Vaporization
Liquid to gas
Melting
Solid to liquid
Molality
(m) = 1 mole / 1000 g
Normality
(N) = one equivalent / 1 Liter
log base a of a
1
log base a (M^k)
KlogaM
log base a (MN)
loga(M) + loga(N)
log base a(M/N)
logaM - logaN
10^(log10(M))
M
First law of thermodynamics (equation)
ΔE = Q - W Q= heat W = work
Q greater than 0, heat is….?
Absorbed
Q<0, heat is…?
Released
Work done on the system, W=…?
W less than 0
Work done by the system, W=…?
W is greater than 0
Calorimetry
Q=…?
Q=…?
Q=mcΔT
Q=mL
Change in Gibbs Free Energy (equation)
ΔG = ΔH - TΔS
Rate equation
Rate = k[A]^m[B]^n
A and B are reactants
Rate = Δ[concentration]/Δt
Relationship between Keq and ΔG (equation)
ΔG(naught) = -RTlnKeq
Hz
(1/s)
Wave period
T = 1/f (s)
Define: ionization energy
How much energy is required to remove an e- from a neutral atom in a gaseous phase (kJ or eV)
Relationship between E photon and BE (binding energy) and KE of photon
E = BE + KE
Therefore, hv = BE + KE
Higher binding energies correspond with e- in ______ energy subshells
Lower, e- are closer to nucleus, require more energy to remove
For gases: as temps increase, solubility…?
Decreases (gas molecules move too much)
Gases: increasing pressure ______ solubility
Increases
List forces in order of relative strength from strongest to weakest (polar covalent, non polar covalent, metallic, ionic bonds)
1 (strongest) metallic
2 ionic
3 polar covalent
4 (weakest) nonpolar covalent
When adding an enzyme, does the overall energy released (ΔH) change?
NO (kinetics of rxn are not directly related to thermodynamics of rxn)
What is an acid catalyst?
H+ ion
What is a base catalyst?
OH- ion
Define: heterogeneous catalyst
In a different phase than the reactants
Define: surface catalyst
Reactant molecules are adsorbed onto a solid surface before reacting with the catalyst to form product
- rate increases with the surface area of the catalyst in contact
Keq > 1
Reaction goes in forward direction (favors products)
Le Chatelier’s principle, increase pressure favors which side?
Side w/ fewer molecules
Define: neutralization reaction
Double replacement reaction between an acid and a base
Acid + base = H2O + salt
Define: combustion reaction
Redox reaction between a compound and molecular oxygen (O2) to for, oxygen-containing products (when one of the reactants is a hydrocarbon, the products include CO2 and H2O)
Define: disproportionation reaction
When a single reactant is both oxidized and reduced
ie: 3ClO- -> ClO3- + 2Cl-
Reduction half reaction
X+ + e- = X
Oxidation half reaction
X -> X+ + e-
What does Keq describe? What is excluded?
Molar concentrations (mol/L) at equilibrium for a specific temp Pure solids and pure liquids are excluded
Kp (equation) (2)
Kp = (Pproducts)^coefficient / (Preactants)^coeff.
Kp = Kc(RT)^Δn
Δn = mole of product gas - mole of reactant gas
Describes the ratio of product and reactant concentrations at equilibrium in terms of partial pressures
What is the oxidation state of oxygen when bound to another oxygen (ie in H2O2)?
OS = 1
What is the oxidation state of oxygen when bound to F?
+1 (F is more electronegative than O)
Pressure (equation)
P = F/A F = ma
When temp increases, what happens to pressure?
P increases
When increasing # particles, P does what?
Increases
When volume decreases, P does what?
Increases
What are the three conditions for a gas to be ideal?
- No IMF
- No volume for individual particles
- Collisions between particles are perfectly elastic (KE is conserved)
Boyle’s law
PiVi = PfVf (at constant T)
Charles’s law
(Vi/Ti) = (Vf/Tf) (at constant P)
How much volume (L) does one mole of gas take up?
22.4L
Avogadro’s law
(Vi/ni) = (Vf/nf) (at constant T & P)
Due to Van de Walls forces between gas particles, what is the relationship of Preal to Pideal?
Preal < Pideal
Due to Van de Walls forces between gas particles, what is the relationship of Vreal to Videal?
Vreal > Videal
What are the 3 R values for ideal gas law?
R = 8.3145 (m^3*Pa)/(mol*K) R = 8.3145 (L*kPa)/(mol*K) R = 0.082 (L*atm)/(mol*K)
How do you calculate partial pressures (given total mass of total gas and percent compositions)
For each gas:
- multiply percent composition by total mass (g)
- divide by molar mass of gas to get mol of each gas (n)
- Plug into PV=nRT to get P, which is total P
- Multiply mol of individual gas as percent (ie 3 mol gas = 3%) by total pressure to get partial pressure of each gas
Define: vapor pressure
Pressure at which vapor and liquid state are at equilibrium (pressure is high enough for H2O molecules in liquid to reach high enough KE to change to gas, but low enough to force vapor H2O molecules to convert back to liquid)
Define Specific heat and give the equation
Amount of heat energy needed to raise 1g of water by 1C
Q = mcΔT Q = amount of heat released or absorbed (J) c = specific heat (amount of heat energy stored before T increases) m = mass
Heat of fusion/vaporization (equation)
Q = mL Q = amount of heat released or absorbed (J) L = latent heat (how much heat energy needed to change phase) m = mass
At what T does the ideal gas law start to break down?
Low T
If an inert gas is added, what happens to the Ptotal and partial pressures of other gases (Le Chatelier’s principle)?
Ptotal increases, but partial pressures of gas stay the same
Q > Keq, favors …?
Reactants
Q < Keq favors…?
Favors products
CO(g) + 2H2(g) -> CH3OH (g)
If Q < Kc, reaction favors _____ , which favors side with (fewer, more) molecules, causing Ptotal to (increase/decrease)?
Favors products, favoring side with fewer molecules, Ptotal decreases
Define: hydration
Water molecules break ionic bonds to dissolve
Calculating solubility (g/L) requires what type of table?
ICE
Solubility rules mnemonics (2)
All soluble: CASH-N-GIA
Chlorates, acetates, sulfates, halides (not F), nitrates, Group 1 metals, ammonium
All insoluble: HAPpy
Hg2 2+, Ag+, Pb2+)
Why is water less dense when frozen than liquid?
Due to the way H bonds are oriented as it freezes, the water molecules are pushed further apart than they are in liquid water
Why does it take a lot of heat to raise the T of water?
Because some of the heat must be used to break H bonds between the molecules (water has a high specific heat capacity)
Define: evaporative cooling
As water molecules evaporate, the surface they evaporate from gets cooler because the molecules with the highest kinetic energy are lost to evaporation
Define: cohesion
Attraction of molecules for other molecules of the same kind (responsible for surface tension)
Define: adhesion
Attraction of molecules of one kind for molecules of a different kind (responsible for capillary action)
Define: 0th Law of Thermodynamics
No heat is transferred between two objects in thermal equilibrium (they are the same T)
Define: work
The energy required to move something against a force
What is the relationship between W, P and V for expansion or compression of gases (work done by gases for expansion/compression)?
W = -PΔV (J)
**P is the external P
When a system does work on its surroundings, the systems internal energy _____?
Decreases (negative work)
Relate W, force, and displacement (equation)
W = F x d
What is the work done on a gas (or by a gas) in a rigid container (no change in volume is possible)? How do changes in energy occur?
No possible change in volume means it is not possible for gas to do work because ΔV =0, so
W =0
Changes in energy must occur through other methods such as heat
Define: endothermic
Heat absorbed or released from a system under constant P
Breaking bonds between atoms requires (adding/removing) energy?
Adding
Forming new bonds (releases/takes) energy
Releases
Define: exothermic
Reactions where the products have lower PE than reactants
Define: 2nd Law of Thermodynamics
The entropy of the universe is always increasing for a spontaneous process
ΔG = ΔH + TΔS is true when what are held constant?
Temperature and pressure
When ΔG is negative at constant T and P, the reaction is _____ and considered (exergonic/endergonic)?
Negative, exergonic
When ΔH < 0 and ΔS > 0, ΔG is…?
ΔG < 0, spontaneous at all T
When ΔH < 0 and ΔS < 0, ΔG is…?
Spontaneous at low T (when TΔS is small)
When ΔH > 0 and ΔS > 0, ΔG is…?
Spontaneous at high T (when TΔS is large)
When ΔH > 0 and ΔS < 0, ΔG is…?
Non-spontaneous at all T (ΔG > 0)
Define: 1st Law of Thermodynamics
Energy can neither be created nor destroyed
Define: open system
A system that freely exchanges both energy and matter with its surroundings
Define: closed system
A system that exchanges only energy with its surroundings, not matter
If ΔS > 0, the reaction is (spontaneous/non-spontaneous)?
Spontaneous
If ΔS < 0, the reaction is (spontaneous/non-spontaneous)?
Non-spontaneous
Water going from a solid to liquid is an (endothermic/exothermic) process?
Endothermic, heat absorbed, ΔH positive (chemical bonds broken)
Water going from a liquid to gas is an (endothermic/exothermic) process?
Endothermic, heat absorbed, ΔH positive (chemical bonds broke )
Water going from a gas to liquid is an (endothermic/exothermic) process?
Exothermic, heat released, ΔH negative (chemical bonds formed)
Water going from a liquid to solid is an (endothermic/exothermic) process?
Exothermic, heat released, ΔH negative (chemical bonds formed)
Equation for ΔH (relating to bonds formed/broken)
ΔH = ΔH (bonds broken in reactants) - ΔH (bonds formed in products)
Enthalpy equation ( H, E, P, V)
H = E + PV
Define: critical point
Point on phase diagram where phase can coexist (very high T and P)
Define: triple point
Point on phase diagram where all three phases coexist (intersection point), equilibrium between phases
