C4 Compounds and Stoichiometry
scientific notation 6.2 E 8 is 62 E _____ 6.2 E-9 is 62 E ____
62E7 62E-10 move to the right goes down
bomb calorimeter - what does it measure and what are the conditions that allow for this how would you solve for variables in terms of delta U of system and surroundings and heat?
a device for measuring the heat transferred to surroundings (potentially water) in the combustion of a substance under constant-volume conditions volumetric - means W=0 delta U system = - delta u surroundings q system = -q surroundings
expressing concentration as normality formula and why it is unique
# of equivalents / L SOLUTION *reaction dependent - in acid base we are looking at H+ concentration, oxidation reduction - e- concentration think of this as molarity of what is important / of interest
van hoff factor NaCl vs glucose
# of particles dissolved in sol'n glucose - not readily dissolved - i = 1 NaCl - i = 2
how to calculate percent composition
(mass in formula / molar mass) *100
if a gas within a system is being cooled and the PV graph shows a hyperbolic curve what kind of process may be indicated adiabatic isobaric isothermal isochoric
- cooled = change in temp = adiabatic
Isolated vs. closed vs. open systems
- isolated: no exchange matter or energy w/ environment - closed: exchange energy, but not matter - open: exchange both
Looking at the following equation proceeding forward what are the relative rates and what are the units? · 2A +B --> C
-[A]/2t = -[B]/t = +[C]/t -note negative/positive signs and mols in A -units: concentration over time --> M/s
Things to note about transition state
-can't be isolated -theoretical structure with more or equal energy required than Ea -can move towards products or reactants
what are the state functions and what what do they describe what are process functions
-describe a system in equilibrium, kinetics or thermodynamics, and are independant of path taken: pressure density temperature volume (H) enthalpy (U) internal temp Gibbs Free Energy (S) Entropy MNEMONIC: when i feel PRESSURE and DENSE i watch TV and want HUGS process functions - work and heat describe the path taken
spontaneous processes: things to know regarding free energy, ea, how they proceed
-gibbs free energy of rxn is negative - products more negative G than reactants -but tend to have high Ea - so may never occur - enzymes help biological processes which might need to occur to sustain life -they can go to completion but may also stay in equilibria state
zero order reaction what is rate equal to, how can we change the rate, what are the units of k, how does this relate to slope on a concentration vs. time graph
-product is independent of rectant concentration -only temperature can change rate by changing k rate = k (units: M/s) [concentration] vs. time k = -slope (linear)
second order reaction what is rate equal to, what are the units of k, how does this relate to slope on a 1/[concentration] vs. time graph
-rate is proportional to either the concentrations of 2 reactants or to the square of the concentrations of a single reactant Rate = k[A]1[B]1 or k [A]2 or k [B]2 k units: 1/ M*s 1/([A]) vs time - makes linear k= + slope
Mixed-order reactions define, what is rate equal to, what are the units of k note: not responsible for derivation but looking at a rate law understand how the order will change if given concentration changes, for example: rate = k1[A]2/k2+k3[A] with large [A]
-rate order that changes over time (course of rxn) -fraction or broken orders observed EX: since [A] is large k3[A] >>> k2 and rn will appear first order
When is solvation exothermic? favored at high or low temp?
-when the new interactions are stronger than the original ones -favored at low temperatures -usually gases dissolved in liquids are exothermic, le chat explains why lowering temp will result in gas dissolving in liquid
when is solvation endothermic? favored at high or low temp?
-when the new interactions are weaker than the original (heat supplied to form weaker less stable solutions) -favored at high temperatures **more common to be endothermic
standard measurements for pressure 1 atm = _________ mm hg = ___________ torr = ___________ kPa
1 atm = 760 mm hg = 760 torr = 101.325 kPa
cm (cubed) to mL
1 cm cubed = 1 mL
which of the following has the most exothermic standard heat or combustion? endothermic? A. ethane B. propane C. n-butane D. n-pentane
1. D --> most exothermic - more products and heat realeased - aka longest hydrocarbon 2. A --> Endo - shortest hydrocarbon
Steps to determine empirical formula given percentages composition
1. Find # mols each substance assume a 100g sample and make the % --> g and divide by the molar mass of that given element (EX: 40% C --> 40g/(12g/mol) --> # mols 2. divide by smallest mols 3. try to make empirical formula 4. to find molecular formula divide molecular/empirical molar mass and then multiply coefficients
For the following generalized formulas of sparingly soluble salts what would the generalized Ksp and what would represent the molar solubility in each situation assuming no common ion effect (it may help to draw out balanced equation) 1. MX --> EX: CuBr 2. MX2 --> EX: CaF2 3. MX3
1. Ksp = [M+][X-] = x^2 2. Ksp = [M2+][2X-]^2 = 4x^3 3. Ksp = [M3+][3X-]^3 = 27x^4 notice Ksp increases --> solubility increasing in all cases x is the molar solubility assuming no common ion effect
rules when determining the limiting reagent
1. always compare in mols 2. combine absolute mols with stoichemetric ratio for given reaction
How do the following effect the reaction coefficient "Q" of a reaction 1. increasing product 2. increasing reactant 3. increasing pressure 4. increasing volume 5. increasing temp
1. increasing product = Qc increase 2. increasing reactant = Qc decrease 3. increasing pressure = change Qp such that it will shift to side containing fewer mols of gas 4. increasing volume = decrease p = change Qp such that it will shift to side containing more mols of gas 5. increasing temp --> no effect on Q (stays the same), this will change Keq
Fill in the blanks with ionic / polar covalent / NP covalent compounds _____ make good electrolytes and dissolve most readily ____ are the weakest and do not form current carrying ions
1. ionic compounds --> not in solid state (after ionic strong polar covalent will dissolve most readily -->HCl in water) 2. NP covalent compounds --> gases and organic compounds such as o2, glucose
understanding an isobaric process: 1. what is constant, if anything 2. explain if and how it alters the first law of thermodynamics 3. how is this curve represented on PV graph
1. pressure is constant 2. does not alter --> deltaU = Q-W 3. flat horizontal line on PV graph with slope=0
explain enthalpy, free energy, entropy and how they change during an explosion
1. releasing heat energy - exothermic ΔH<0 2. entropy is increasing, disperse energy ΔS > 0 (positive) 3. ΔG < 0 spontaneous and exergonic since ΔG = ΔH - TΔS
understanding an isothermal process: 1. what is constant, if anything 2. explain if and how it alters the first law of thermodynamics 3. how is this curve represented on PV graph
1. temperature is constant - therefore the internal energy remains constant --> deltaU = 0, Q=W 2. This means that on a P vs V graph the area under the hyperbolic curve = W = Q (heat added = work done by the system)
understanding an isochoric or isovolumetric process: 1. what is constant, if anything 2. explain if and how it alters the first law of thermodynamics 3. how is this curve represented on PV graph
1. volume is constant - therefore no work is done (gas is not expanding or compressed) 2. deltaU=Q -change in internal energy = heat added to system 3. straight vertical line, 0 area under curve, no work done
1/9 2/9 3/9
1/9 = 0.111 2/9 =0.222 3/9 = 0.333
1 mol of ideal gas at STP is how many L? What is STP?
22.4 L standard pressure = 1 atm standard temp = 0C = 273K *don't confuse with standard conditions
What is a colligative property? list them
A physical property derived solely from the number of particles present, not the chemical nature of those particles *usually associated with dilute solutions: freezing point depression boiling point elevation vapor pressure depression osmotic pressure
which will increase molar solubility? A. Complex ion formation B. Common ion effect C. Both
A. Complex ion formation
isochoric (isovolumetric) process is best described as: A. internal energy is = to heat added to system B. internal energy is = heat removed from a system C. internal energy is = work done on a system D. internal energy is = work done by a system
A. internal energy is = to heat added to system no work done so deltaU=Q
a system in which the heat added is = to the work done by the system A. isothermal B. adiabatic C. isobaric D. isochoric
A. isothermal temp is constant, internal energy constant so Q=W
solubility rules regarding hydroxides *know generally
All hydroxides are insoluble, except when formed with Ca2+, Ba2+, Sr2+
explain how VP depression and BP elevation are related
As VP of a solution decreases --> higher temp required to reach atmospheric pressure --> raise BP
As more solute B is dissolved in solvent A what happens to the vapor pressure or each? the mol fraction of each? explain
B dissolves into A: VP and mol fraction of B increases VP and mol fraction of A decreases - the presence of solute blocks solvents ability to evaporate, but has no effect on condensation rate (molecular basis of raoults law)
VP of water at boiling point?
BP = 100C VP = 1 atm
When a nonvolatile solute is dissolved in a volatile solvent, which characteristic is greater for the solution than for the solvent?
BP is greater for solution than pure solvent formula given by: deltaTb = i*Kb*molality
a leak of helium gas through a small hole occurs at a rate of 3.22E-5 mol/s. How will leak rates of Ne and O gases compare? A. Ne leaks faster, O leaks slower B. Both leak faster C. Ne leaks slower, O leaks faster D. Both leak slower
Both leak slower Graham's Law r1/r2 = inverse square root of molar mass (M2/M1) O and Ne both have > MM than Ne
in which of the following situations is it impossible to predict how the pressure will change for a gas A. gas cooled at constant V B. gas heated at constant V C. gas heated, volume increased simultaneously D. gas cooled, volume increased simultaneously
C not D because decreasing temp and increasing volume will both have same effect --> decrease P
which of the following phase changes is associated with largest decrease in entropy? fusion solidification deposition sublimation
C. deposition decrease disorder - gas to solid
adiabatic process is best described as: A. internal energy is = to heat added to system B. internal energy is = heat removed from a system C. internal energy is = work done on a system D. internal energy is = work done by a system
C. internal energy is = work done on a system since deltaU= -W
when water molecules become ordered around an ion as it dissolves - ordering makes a negative contribution to change in entropy of the solution. An ion with more charge density will have a greater hydration effect or ordering of water molecules. Based on this info which of the following compounds will have the most negative contribution to entropy of the solution? KCl LiF CaS NaCl
CaS --> since Ca2+ and S2- the rest all have +/- 1 charge densities
which of the following statements is true of a process that is spontaneous in the forward direction? ΔG > 0, Keq < Q ΔG > 0, Keq > Q ΔG < 0, Keq < Q ΔG < 0, Keq > Q
D Explained: has to be c or d --> negative ΔG = spontaneous ΔG = RTln(Q/Keq) --> ln of a fraction is negative so Keq > Q will make this true
be able to decide if entropy is increasing or decreasing
EX: - increasing if moving to gas phase - increase if heat is being transferred between system and surroundings - increase if moving towards more moles of gas - increase if moving towards more particles
Celcius to Fahrenheit
F=9/5C+32
What is heat capacity?
Heat capacity is mass*specific heat - measure of the heat required to raise the temperature of any given amount of substance 1 C
what law states that the solubility of a gas is directly proportional to it's partial pressure how is the equilibrium between the gas and liquid state expressed
Henry's Law equilibrium between condensation and evap: [A]1/P1 = [A]2/P2 = henry's constant "k" or [A] = K P
how does common ion effect change Ksp?
IT DOESN'T IT DECREASES MOLAR SOLUBILITY
Chelation -what is its process, structures associated, and what is it used for
In complex ions --> the process of binding central metal cations (lewis acid/ e- acceptor) to the same ligand at multiple points used to sequester toxic metals
Acetic acid dissociates in solution according to the following equation; CHCOOH <--> CH3COO- + H+ If sodium acetate, is added to a solution of acetic acid in excess water, which of the following effects would be observed in the solution? Decreased pH Increased pH Decreased pKeq (pKa) Increased pKeq (pKa)
Increased pH addition of acetate - yield acetate ions - which is a product - according to Le Chat the rxn will then shift toward the left favoring reactants (less H+) - increase pH - more basic
According the Kinetic Molecular Theory of Gases if all gases are at the same temperature what can be said about their KE how can we compare their speed
KE is the same thus we can compare their speed solely from their mass - larger = slower
K (equilibrium constant) don't confuse with reaction rate constant "k"!!!!!!!
Kc = [products]^n / [reactants]^n Kp (for gases) = Pproducts^n / Preactants^n stoichiometric coefficients are n here!! in rate law remember that they do not matter!!!!!!!
things to know about equilibrium constant 1. Keq = 2. if: Q>Keq Q<Keq 3. what substances are negligible 4. Dependent upon 5. reverse reaction is 6. if Keq is a strong negative or positive exponent which direction is favored
Keq = [products]n/[reactants]n n= Stoich. Coeff. Q>Keq - more products, rxn move reverse, positive gibbs free E Q<Keq - more reactants, rxn move forward, negative gibbs free E o Pure solids and liquids Keq = 1, and can be ignored o Dependent upon temperature o Reverse is 1/Keq strong negative - way less than 1 - favors reactants
What is Kf (formation / stability constant)? what is this formation referring to, what does it depend on, explain role of le chatelier
Kf - constant represents formation of a complex ion in formed solution (same equation as Keq = [products]^n / [reactants]^n -This step is dependent on the dissolution of a given compound in sol'n to form ion first before ion forms a complex -Kf > Ksp of compound providing ion -Formation of complex ion will increase initial compound dissolution --> shift reaction to the right to form more ions to be used in the complex (le chatelier)
if asked to find the equilibrium constant of an equation containig solely gases what are you looking for
Kp!! not Kc - this is concentration
ksp vs kf
Ksp is the solubility constant - using constituents of dissolved solute at equilibrium Kf is the formation or stability constant - formation of complex ion within a solution -larger than Ksp compound must dissolve in solution (rate limiting step to produce a given ion) --> the ion will then be sequestered to produce a complex ion
which conditions are least likely to result in ideal gas behavior high P low T high P high T low P low T low P high T
LEAST: low temp and high P
water equilibrium temp between liquid and solid phase - melt point? BP
MP 0C BP 100C (VP=1atm)
***super important things to know**** the natural log of less than 1 for example ln(0.5) will be negative or positive or zero? e is approximately
NEGATIVE 2.7
NaCl dissolution in water process note entropy and enthalpy for both and overall
NaCl solute; h2o solvent Enthalpy: -Na and Cl ionic bonds broken and H-bonds of H2O broken - endothermic -forming of new dipoles is exothermic, but still less than the process of breaking -overall process endothermic (only slightly), and favored at high temps Entropy: -entropy increases for nacl - freed from lattice structure, more microstates -entropy of water decreases - more restricted -entropy increase Nacl > decrease water - overall + entropy overall process is spontaneous
Gay-Lussac's Law
P1/T1=P2/T2 (directly proportional)
Combined Gas Law
P1V1/T1=P2V2/T2 USE SI UNITS
Boyle's Law
P1V1=P2V2 volume and pressure are inversely proportional in isothermal conditions
how does doubling temp affect the volume of an ideal gas? doubling pressure? both at same time
PV=nRT A. 2T = 2V B. 2P=0.5V C. no change
Raoults Law --> Pa=XaPa* explain the variables
Pa=XaPa* Pa = VP solvent A when solutes present Xa = mol fraction of solvent in solution Pa*= VP solvent A in pure state
unlike changes in concentration and pressure/volume - when temperature is altered in a system what is changing
Qc and Qp effected in concentration and pressure, respectively temperature doesn't change reaction coefficient - but alters Keq itself (remember Keq temp. dependant)
the vapor pressure of a solution is directly proportional to the mole fraction of solvent present in ideal solutions according to
Raoult's Law
irreversible reactions key characteristics
Reactions that precede in one direction only (no equilibrium), and reaction goes to completion Maximum amount of product is formed and determined by limiting reagent
a reaction is found to stop just before all reactants are converted to product. Which is true: Rxn is irreversible, forward rate > reverse rate Rxn is reversible, forward rate > reverse rate Rxn is irreversible, forward rate < reverse rate Rxn is reversible, forward rate > reverse rate Rxn is reversible, forward rate = reverse rate Rxn is reversible, forward rate < reverse rate
Rxn is reversible, forward rate = reverse rate EXPLAINED: scenario described indicates - equilibria is reached, which is far to the right - more products this cannot be irreversible reaction because it does not go to completion - can't be option A, B, C Since the rxn is at equilibrium forward rate must = reverse rate
DO NOT CONFUSE STP AND standard conditions - what are the difference
STP - gas law calculations - standard T = 0C (273K) and standard P = 1 atm standard conditions of a system/substance - most stable form - 25C (298K), 1 atm, 1M concentrations
Charles' Law
V1/T1=V2/T2 volume and temp are proportional when P is constant
energy required to reach the transition state
activation energy
what is avogadros principle (conditions and definition) and what is the equation associated with this in relation to the ideal gas law
all gases at a constant pressure and temperature have a volume proportional to the number of mols of gas At STP 1 mol is 22.4 L n1/V1=n2/V2
what are 4 anions and cations contained in salts--> just know they are always water soluble according to solubility rules 1 and 2
ammonium (Nh4+) alkali metals - group 1 acetate (Ch3coo-) nitrate (NO3-)
which phases of solvent and solute can form a solution?
any combo of phases as long as at leads homogenous mixture
what is the specific heat of liquid water and what does specific heat imply
c h20 (l) =1 cal/gK or 4.184 J/gK the amount of heat energy needed to raise 1 g of substance 1 degree C
device most appropriate to measure heat capacity of a liquid
calorimeter
phase changes key points -what conditions -what kind of system
can occur under standard or non-standard conditions isolated system - no matter or energy exchanged w/ environment changes are reversible (reverse = forward rate, as with equilibrium, and the amount of substance in 1 phase = to amount in the phase change)
solubility rules for carbonates (CO32-) and phosphates (PO43-) and sulfides (S2-) and sulfites (SO32-) *know generally
carbonates (CO32-) and phosphates (PO43-) and sulfides (S2-) and sulfites (SO32-) insoluble unless formed with alkali metals and ammonium
what is a complex ion aka coordinate compound think of some examples what bonds hold them together
cation "lewis acid" is bonded to at least 1 e- pair donor "ligand, lewis base" EX: H+ bonded to H2O EX: Fe of Hb bonded to O2/Co2 EX: transition metals within coenzymes bind ligands coordinate covalent
what does solubility depend on
changes in enthalpy and entropy temperature the type of solvent complex ion formation (Kf) - rare, increases solubility common ion effect gas phase --> pressure
freezing point depression boiling point elevation vapor pressure depression osmotic pressure are all?
colligative properties - physical properties derived solely from the number of particles present, not the chemical nature of those particles usually associated with dilute solutions
physical properties derived solely from the number of particles present, not the chemical nature of those particles defines a
colligative property
in hemoglobin the Fe cation forms a ________ with o2, co2 via a _______ bond Fe cation is the electron _______ aka lewis _______? o2 or co2 are the ? which is the "ligand"?
complex ion aka coordinate compound coordinate covalent bond Fe--> electron acceptor, lewis acid gases --> electron donor, ligand, lewis base
Amount of solute dissolved in a solvent is referred to as
concentration
how are condensation and evaporation rates affected by solute particles
condensation rate unaffected evaporation rate blocked overall vapor pressure of solvent decreases (molecular basis of raoults law)
gas to liquid? endothermic or exothermic? at what temp or pressure conditions does this occur
condensation; exothermic lower temp or high pressure
Under what conditions are enthalpy and heat equal? if heat is absorbed into a system the process is? if it is released? what is standard enthalpy of formation?
constant pressure absorbed - +deltaH- endothermic released- -deltaH - exothermic std H formation - enthalpy required to produce 1 mol of compound from elements in standard states
under what conditions does entropy increase upon dissolution (forming a solution)
constant temp and pressure
protons are rarely isolated in a solution and typically form what kind of bond with an electron pair carrier (donor)
coordinate covalent bond
rearranging ideal gas law to solve for density
d=m/V but with gas law d=PM/RT
decreasing the partial of pressure of a gas above a liquid will effect solubility how?
decrease solubility
how does freezing point change due to additional solute to a pure solvent what is the formula to calculate
decreases - solute particles interrupt solid lattice structure - more E must be removed in order to freeze ΔTf = i*Kf*m ΔTf - decrease in FP (not the FP itself) i - van hoff factor Kf - proportionality constant characteristic of solvent
how does a catalyst effect a reaction? the free energies of products or reactants? the activation energy?
decreases Ea - energy required to reach transition state no effect on free energy of products or reactants will increase reaction rate- reaction constant increase will move to equilibrium FASTER- low energy state where products = reactants but does not change the position
when will molarity and molality be approximately equal
dilute aqueous solutions at 25C
is bond formation or dissociation generally endothermic?
dissociation = endothermic
a ligand in a coordinate compound is the
electron pair donor (lewis base) to which a cation is coordinate covalently bound
Solvation -what is this process -aka -explain relation to endo or exothermic and which occurs more often
electrostatic interaction; process of surrounding solute particles with solvent particles to form a solution, involves breaking intermolecular forces and forming new ones aka dissolution or hydration when referring to water solvent endothermic-more often-when new interactions formed are weaker than those broken - forming new weaker and less stable solution
if the delta G of the reactants is less than the products the reaction is
endergonic
the process of forming a salt solution can be better understood in 3 steps 1. break solute into individual components 2. Make room for solute in solvent by overcoming intermolecular forces in the solvent 3. Allowing solute-solvent interactions to occur to form the solution explain enthalpy at each step
endo, endo, exothermic 1. endothermic - energy required to break molecules 2. endothermic - energy required to make room - must overcome intermolecular attractions 3. exothermic - forming new stable solution in which polar water interacts with the ions
evaporation is ________ (endo/exothermic) and the heat source is the _________
endothermic; liquid
spontaneous distribution of energy within a system or between system and surroundings is known as what law of thermodynamics does this relate to when is this distribution maximized
entropy -more commonly thought of as disorder of a system -state function - (products -reactants), indep of path -ratio of heat transfer/temp -second law max@equilibrium
1. liquid to gas is known as? is this process endothermic, exothermic, or neither? compare this process to boiling. when this happens what happens to remaining liquid temperature
evaporation or vaporization (can occur at any temp, but boiling is special form in which molecules rapidly evap above boiling point), endothermic liquid has lost high energy particle, temp of liquid decrease
if the delta G of the products is less than the reactants the reaction is
exergonic think of the free energy graphs like a slide
Kinetic Molecular Theory of Gases -This theory describes what -what are the Assumptions -what is proportional to what - explain and show the mathematical relationship
explain behavior of gases Assumptions: 1. volume of gas particles is negligible to volume of container it is in (individual particles don't have volume but gas as a whole does) 2. particles in constant, random motion that exhibit no intermolecular attractions and undergo completely elastic collisions with each other and with the walls of the container -elastic meaning momentum and KE conserved 3. Average KE of individual gas particles is directly proportional to temp of gas (K) 4. Average KE of any gas as a whole will have the same temp KE = 0.5mv^2 = 1.5*Kb*T Kb is constant v is the average molecular speed of gas as whole - impossible to isolate single gas particle velocity
true or false the rate of the rxn is related to gibbs free E
false
true or false gibbs free energy is temperature independant
false - temp dependent
true or false in an equilibrium reaction: increasing k1 (reaction rate constant) results in decrease of k-1
false if the forward reaction constant increases so will the reverse constant equilibria baby!
What is vapor pressure? How is it effected by temp? How is it related to BP?
force exerted by a gas on liquid at equilibrium as temperature increases, vapor pressure increases - higher temp = more molecules entering gas phase and increasing the pressure gas exerts on a liquid when VP = external or (ambient, incident) P - this is BP
ΔTb = i*Kb*molality what is the formula for and what does Kb depend on
formula to measure changing BP relative to a pure solvent (BP of non-pure solvent will be greater than a pure one) Kb depends on the solvent properties
when a system is at equilibrium explain what is occuring between forward and reverse rates? what happens to entropy? to gibbs free E?
forward rate = reverse rate gibbs free energy is at a minimum = 0 entropy is at maximum
reaction coefficient (Q) > Keq reactants or products favored? reverse rate or forward rate will increase?
forward reaction exceeded equilibrium - more products reverse rate will increase to restore equilibrium --> positive delta G
ΔTf = i*Kf*m what is the formula for and what does Kf depend on
freezing point depression - how the freezing point will change compared to a pure solvent (solute particles interrupt solid lattice structure - more E must be removed in order to freeze) Kf depends on the solvent properties, don't confuse with formation constant for complex ions - this is a proportionality constant to measure change in freezing point
characteristics that make the gas phase unique? what is assumed of an ideal gas?
gas phase characteristics: -easily (not infinitely) compressible -weak intermolecular forces due to large intermolecular distances allow expansion -fluid that flows and fills shape of given container -molecules move rapidly ideal gas: -no intermolecular forces between molecules -individual molecules occupy no volume -PV=nRT
understanding of g equivalents? what is gram equivalent of h2s04?
gram equivalents = molar mass/ parts equivalent "n" (# mols, H+, etc) h2s04 MM = 98 and there are 2 mol H+/ H2so4 so 92/2 = 49 g
solubility rule regarding halides *know generally
halides - Cl-, Br-, I- (except fluorine) are water soluble unless in a compound with Ag+, Pb2+, or Hg22+
at what temperature and pressure conditions are real gases similarly approximated to ideal gases in what ways do real gases deviate from ideal gases
high temps, low pressures real gases have intermolecular forces and their volume isn't negligible - especially at high pressures and low temps
which statement(s) describe kinetics of this reaction rate = k[a][b] i. reaction is second order ii. a consumed = b consumed iii. the rate will not be affected by the addition of a third compound i ii iii i and ii i and iii ii and iii all none
i only ii is wrong because they may have any stoichiometric ratios- which is not told in rate kinetics
what will occur in the following scenarios if IP>Ksp if IP<Ksp if IP=Ksp
if IP>Ksp - supersaturated, precipitation is thermodynamically favored if IP<Ksp - dilute, dissolution will occur (therm favored) if IP=Ksp -saturated at equilibrium
as the pressure of a real gas increases - explain what occurs in terms of intermolecular forces explain how this would also deviate from volume estimated using ideal gas calculations? explain how these calculations would differ at extremely high pressures?
increase pressure - more intermolecular forces until gas condenses to liquid the volume calculated in ideal gas calculations will be much larger than the actual volume (real volume is less) - due to intermolecular attraction However, at very high pressures the real volume will be underestimated by ideal gas calculations because the particles will take up a larger volume (particles are large relative to the distance between them)
increasing the partial of pressure of a gas above a liquid will effect solubility how?
increase solubility (henry's law)
how does increasing temperature effect the solubility of a solid vs gas
increase temp: solid increase solubility, gas decrease solubility
formation of complex ions has what effect on solubility? relative Ksp to other reactions?
increases solubility (these reactions will have higher Ksp values in the end --> Kf driving forward) dissolution stabilized by dipole-dipole
How will increasing and decreasing temperature effect non-gas solutes Ksp?
increasing temp will increase Ksp note: gas solutes are also dependent on pressure
what is ion product and how does it relate to Ksp and other equilibrium
ion product is the equivalent of "Q" for other equilibria Represents ionic compound ion constituents that are not at equilibrium and will predict whether dissolution or precipitation will occur if IP>Ksp - supersaturated, precipitation occurs if IP<Ksp - dilute, dissolution will occur if IP=Ksp -saturated at equilibrium
Reactions that precede in one direction only, and reaction goes to completion Maximum amount of product is formed and determined by limiting reagent
irreversible
which derivations of the ideal gas law assume: -isothermal conditions -isobaric -isovolumetric
isothermal: boyles law p1v1=p2v2 - inversely proportional isobaric: charles law V1/T1=V2/T2 - directly proportional isovolumetric: gay lussacs P1/T1=P2/T2
in pH problems if given the concentration of reacting compounds containing nitrate (not redox) or sodium --> these will translate to the ion concentrations
just know this
removing heat / (aka reaction at lower temperature) will result in formation of kinetic or thermodynamic product? is product more or less stable? does this product have a higher or lower free energy relative to other product (if kinetic other product is thermo, if thermo other product is kinetic) does the transition state require more or less E to be reached? form faster or slower?
kinetic product product is less stable - higher free E - relative to thermo product transition state requires less E, form faster
As related to kinetic molecular theory of gases, what will happen to the speed of gas particles: - if mass increases -if temp increases
larger = slower speed / diffusion higher temp = faster
expressing concentration as percent composition by mass
mass of solute / mass of SOLUTION *100
solubility rules regarding metal oxides *know generally
metal oxides are INSOLUBLE, with exception of those formed with alkali metals, ammonium, CaO, SrO, and BaO --> will hydrolyze to form solutions of corresponding metal hydroxide
a gas dissolved in another gas is a
mixture (a solution)
How to find normality
molarity * parts equivalent = equivalents per liter
What is molar solubility? How is it useful in calculations
molarity of a solute in a saturated solution at equilibrium, no more solute will dissolve can use this and a balanced equation to determine concentrations of solute to calculate Ksp
expressing concentration as molality when is this notation important
mols solute / kg solvent freezing point depression and boiling point elevation
expressing concentration as molarity
mols solute/ L SOLUTION usually how solution concentrations are expressed --. rate laws, law of mass action, osmotic pressure, pH, pOH, nernst
reaction coefficient (Q) < Keq reactants or products favored? reverse rate or forward rate will increase? what does this mean for gibbs free energy?
more reactants forward will increase --> negative delta G
when a solution is diluted - adding more solute will cause what thermodynamically favored process to occur
move forward towards equilibrium - fully saturated (rate of precipitation is slower)
do catalysts effect Keq? do they effect reaction rate k?
no yes - increase rxn rate
adiabatic process: 1. what is constant, if anything 2. explain if and how it alters the first law of thermodynamics 3. how is this curve represented on PV graph
no heat enters the system so Q=0 so: delta internal energy = -W -internal energy is = work done ON the system
to determine if concentration of a reactant converted to product is negligible what is a rule of thumb
not negligible if it's starting concentration is within 2 orders of magnitude of Keq
in a certain equilibrium process, the activation energy of the forward reaction is greater than the reverse. this reaction is endothermic exothermic spontaneous nonspontaneous
not spontaneous --> endergonic *thermic relates to enthalpy
Ideal Gases which are true I. have no volume II. have particles with no attractive forces between them III. have no mass I II III I and II I and III II and III all none
only II each particle has negligible volume to the container it is in, but as a whole gas has measurable volume
the following system obeys second order kinetics 2no2 ---> no3 + no (slow) no3 + co --> no2 + co2 (fast) what is the rate law?
only depends on reactants of rate limiting (slow) squared because first line tells us it's second order kinetics rate = k[no2]^2
Typical products in a combustion reaction
oxidant and hydrocarbon (or sulfur or sugar)
loss of electrons is called? when it is in the form H+? catalyzed by
oxidation? dehydrogenation? dehydrogenases
concentrations are commonly expressed as percent composition, mol fraction, molarity, molality, normality what are the differences in formulas?
percent composition: (mass of solute/ mass of SOLUTION)*100 (%) mol fraction: mols of solute/ mol of all species molarity: mols solute/ L SOLUTION (M) molality: mols solute/kg SOLVENT (m) normality: # of equivalents/ L SOLUTION (N)
what is osmotic pressure?
pressure that counteracts osmotic movement across a selectively permeable membrane "sucking pressure" pi=iMRT
increasing temperature can alter Keq of rn. why might this be unfavorable indefinitely? keq has defined limit, which can't be surpassed products or reactants may decompose increasing temp decreases P - which could alter the rxn conditions if the rxn is reversible Keq resists changes in temp
products or reactants may decompose wrong answers explained A- yes rxn has limit but this doesn't answer W B - increasing temp will increase P if volume is constant irreversible rxns (which do not exhibit equilibria) are rxns which resist changes in temp
first order reaction what is rate equal to, what are the units of k, how does this relate to slope on a 1/[concentration] vs. time graph
rate is proportional to the concentration of only one reactant, does not require another molecule or chemical/physical interaction rate = k [a] or (- delta[a]/t) k units: s-1 1/[concentration] vs. time - makes linear k= - slope
what is the Collision Theory of Chemical Kinetics? How does it differ from transition state theory? Understand how different variables effect rate such as: Ea, temp, # of molecules (concentrations)
rate of a reaction is directly proportional to the number of collisions that take place between reactants per second --> products only form if the collision takes place in a specific orientation and if minimum (activation) energy is obtained *diff from transition state theory because it claims all or nothing forward reaction (where as transition state can form products or reform reactants) increases with temperature, decreases if Ea is higher (more energy needed), more collisions - by adding more molecules makes faster other factors such as medium reaction takes place in and the state of the reactants is important
salt KCl dissolves in a beaker and feels cool as it dissolves what can we conclude about enthalpy, entropy, gibbs free E
reaction has -deltaG - because it is proceeding towards dissolution spontaneously the cooling tells us that there is a positive enthalpy and heat has been used to break the bonds knowing this about enthalpy, entropy must be pretty large and positive to overcome enthalpy and yield a negative delta G
if heat is added to an endothermic reaction the equilibrium will shift
right
what will make the following reaction shift right in regards to -temp -products and reactants -temperature -pressure/volume A+2B <-- --> C+heat
shift right towards products - if C is removed - if A or B is added - if pressure is increased (or volume decreased) (right has less mols) -if temp decreases think: le chatelier will make the reversible reaction proceed in favor of counterbalancing
if the pressure in a system is increased it will shift in which direction this is according to
side containing less moles -decrease gas moles and therefore decrease pressure *make sure you only look at mols of gas!!! le chatelier principle and ideal gas law
maximum amount of substance that can be dissolved in a particular solvent at a given temp is known as the solution is considered what is delta G
solubility saturated at max (no more solute will dissolve instead it will precipitate) saturation point = equilibrium -delta G is zero -temp dependent
Henry's Law
solubility of a gas in a liquid is directly proportional to the partial pressure (vapor pressure) that gas exerts on the surface of the liquid equilibrium between condensation and evap: [A]1/P1 = [A]2/P2 = henry's constant "k" or [A] = K * Pa
how can solubility of solids be changed? solubility of gases?
solubility of solids can be increased by increasing temp solubility of gases can be increased by decreasing temp (decrease temp) or according to henry's law increasing partial pressure of gas
co2 gas dissolved in water -what is formed -overall enthalpy -favorable temps
solution exothermic gas dissolution is favorable at lower temps according to le chat
homogenous mixtures of 2 or more substances - solute and solvent what phases
solutions -can be any phases mixed, usually solid in aqueous -gases in gases = mixtures, all mixtures are solutions but not vice versa
pure sodium metal spontaneously combusts upon being placed in room temp water. what is true about the Keq of the combustion rxn at 25C? Keq < 0 0 < keq < 1 Keq = 1 Keq > 1
spontaneous means ΔG is negative equation containing keq and ΔG: ΔGrxn = -RTln(Keq) for ΔG to be negative ln of a # >1 is positive and therefore will be the right answer keq>1
Graham's Law -what is it and what conditions are assumed -note can describe effusion or diffusion using the same equation but what is the difference
states that the rate (avg speed) of effusion/diffusion for a gas is inversely proportional to the square root of its molar mass isobaric and isothermal conditions r1/r2 = sq rt of: M2/M1 diffusion - flow of gas down concentration gradient, think in this case - gases mixing effusion - flow of gas under pressure
solid directly to gas? gas directly to solid?
sublimation deposition
what is the critical point of a phase diagram indicate
temp & pressure showing termination between distinct gas and liquid phases, densities become equal and NO distinction between phases, heat of vap is 0 "supercritical fluids"
temp and heat are not the same - explain
temp - related to avg. KE, state function heat - energy entering or leaving a system due to difference in temp, process function, Q=MCAT related but different
gas solubility product constant dependent on
temp and pressure increasing pressure increases solubility - henrys law
explain how deviations in temperature effect a real gases behavior and deviation from calculations/predictions using ideal gas laws / assumptions
temp decrease - average speed slower and intermolecular attraction is greater as condensation point is reached volume predicted by ideal gas laws is actually smaller (same as if pressure increased) At extremely low temps (same as extremely high pressure) the gas will occupy more volume than predicted - since gas law assumes 0 volume
Ksp depends on
temperature
solubility product constant (Ksp) -what does it describe -what is it's formula -what is a high vs low Ksp
the equilibrium expression for a chemical equation representing the dissolution of a slightly to moderately soluble ionic compound (temperature dependant) uses ionic constituents concentration at equilibrium Ksp = [A^n+]^m[N^m-]^n EX: AgCl --> Ag+ + Cl- --> Ksp = [Ag+][Cl-] *think formula similar to Keq, and because working with pure solids or liquids and don't include them there won't be a denom representing products high Ksp means solute is very soluble
part per million (ppm)
the number of particles of one substance in one million particles of a mixture 1 ppm of substance in water would mean --> 1mg/L or 10^-6 g/L
Hess's Law
the overall enthalpy change in a reaction is equal to the sum of enthalpy changes for the individual steps in the process =products - reactants =bonds broken (endothermic) - bonds formed (exothermic) =E absorbed - E released make sure to multiply by stoich coeff
Raoult's Law -theoretically and mathematical representations -what conditions are necessary for this law to hold true
the vapor pressure of a solution is directly proportional to the mole fraction of solvent present accounts for vapor pressure depression caused by solutes added to a solution Pa=XaPa* Pa = VP solvent A when solutes present Xa = mol fraction of solvent in solution Pa*= VP solvent A in pure state Conditions: must be ideal solutions (enthalpy = 0, strength of attraction of diff components of mixture = same as attraction between any molecules of any 1 component in pure state) *basically want a solution formed of similar components
adding more heat / (aka reaction at higher temperature) will result in formation of kinetic or thermodynamic product? is product more or less stable? does this product have a higher or lower free energy relative to other product (if kinetic other product is thermo, if thermo other product is kinetic) does the transition state require more or less E to be reached? form faster or slower?
thermodynamic favored product is more stable - lower gibbs free E relative to kinetic product transition state require more E to be reached - which is why more heat is needed - form slower
which of the following best describes ionic compounds formed from molecules containing 2+ atoms formed of charged particles, measured by molecular weight formed of charged particles sharing e- equally three dimensional rays of charged particles, determined by formula weight
three dimensional rays of charged particles, determined by formula weight -don't share e- so instead of measuring using molecular weight --> formula weight (calculate same process by adding sum of atomic weights just different nomenclature!)
Barometer at the top of a mountain vs under water
top of the mountain - atm pressure decreases - mercury in column drops under water - increased hydrostatic pressure - mercury in column rises
Dalton's Law of Partial Pressures -what is meant by the law and what are the conditions -equation for finding total pressure given partial P -equation for finding partial pressure given total P
total pressure of a mixture of gases is equal to the sum of the pressures of all the gases in the mixture -each gas will not interact , so pressure exerted as if it is the only gas present -key that it is a mixture of gases -related to mol fraction X Ptotal = Pa + Pb ... X = mols of given gas/ total mols of gas so any given partial pressure = X*Ptotal
point at which three phases exist in equilibrium
triple point
true or false: temperature of a substance in a given phase is related to average kinetic energy
true - all molecules will not have same instantaneous speed --> which is why average
solute and solvent
two parts of a solution solute: dissolves in solvent solvent: remains in same phase after mixing, if both in same phase whichever is in greater quantity, if same quantity whichever is more commonly a solvent
root mean square speed -what is it -equation to know
understanding average speed of gas particles as related to KE average speed = square root of: 3RT/M
if given [product] and [reactant] and keq and you are asked which direction the reaction will proceed and what is the gibbs free energy how would you go about solving this problem?
use the [product]^n / [reactant]^n to calculate the Q now compare Q to the Keq: if Q>Keq the reaction has exceeded equilibrium and will counterbalance by going in the "left" "reverse" direction which is endergonic - +G (opposite is right, forward, exergonic, -G)
Van der Waals equation of state (P + (a*n^2)/V^2 )(V-nb) = nRT -what is it and what do the variables mean -don't memorize equation just understand understand how to compare 2 gases by understanding which variables will be isolated EX 1: if 1 gas has greater attractive forces what will happen to the real pressure EX 2: if 1 molecule is larger
used to correct the ideal gas law for intermolecular attractions (a) and molecular volume (b) a - attraction b - big particles 1. real pressure will be less than ideal - Pr decreases due to greater attractive forces (think of as Pr = Pi - a(n/v)^2 2. if 1 molecule is larger it will have a greater molecular volume and make overall V smaller (V-nb) - therefore pressure will increase *go to chem notes if still don't understand
all salts containing ammonium and alkali metals (group 1) cations and nitrate and acetate anions are considered
water soluble
solubility rule regarding salts containing sulfate (SO4 2-) ions *know generally
water soluble with exception of those formed with Ca2+, Sr2+, Ba2+, Pb2+
double replacement reaction typically forms what as products when acids and bases are involved as the reactants instead it produces
weak solution of electrolytes precipitate or gas that leaves solution acid base - neutralization - produces salt and water
what is the common ion effect? how does it effect molar solubility? how does it effect Ksp?
when a slightly soluble salt is added to a solution which already contains one of its components the added salt is less soluble than if it were added to a pure solvent -molar solubility decreases -no effect on Ksp
how does defining the solubility of the solute relate to gibbs free energy when deltaG is: 1. extremely negative 2. extremely positive 3. slightly negative
when gibbs free E is negative in solvation - the reaction is spontaneous and given solute is considered soluble - molar solubility > 0.1M opposite "insoluble" when the solute equilibrium lies close to dissociation state and only small negative free E we consider them sparingly soluble salts (<0.1M)
what is an ideal solution
when overall strength of original interactions = strength of interactions formed upon dissolution enthalpy = 0
what is a sparingly soluble salt
when the delta G for dissolution is only slightly negative it is not insoluble, but will only dissolve a little EX: commonly ionic compounds in aqueous solutions
when would VP be higher than calculated by raoults law
when the solute-solute and solvent-solvent interactions are stronger than the solute-solvent (particles less likely to stay in solution and evaporate if given 2 compounds see if they are extremely dissimilar - small, hyrophobic vs large, hydrophilic
whenever given a solution problem first step is always
write out balanced equation
Be able to determine kinetic order from a table
ya hurd not the same as stoichiometric
Gibbs free energy equations to know
ΔG = ΔH - TΔS ΔG = -RTlnKeq sum of ΔGproducts - ΔGreactants if not standard ΔG = ΔG* + RTlnQ = RTln(Q/Keq)
boiling point elevation of solution compared to pure solvent - formula
ΔTb = i*Kb*molality ΔTb - increase in BP (not the BP itself) i - van hoff factor Kb - proportionality constant characteristic of solvent
