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Flashcards in Chemistry 3 Deck (68):
1

  • Phase Changes
    • Define:
      • melting point
      • boiling point
      • volatile
      • non-volatile

  • Melting point is
    • the temperature at which a substance changes state from solid to liquid.
    • However, it is very important to realize that melting point and freezing point are exactly the same thing.
    • You might think of freezing point as the temperature at which a liquid changes into a solid, but the value measured for mp or fp is simply a temperature, which indicates no direction of progress.
    • For any substance, mp = fp.
  • Boiling point is the temperature at which a substance changes state from liquid to gas.
    • Liquids boil when the vapor pressure of the liquid equals atmospheric pressure.
  • Volatile
    • a term used to describe the relative tendency of a substance to form a vapor.
    • How readily a substance vaporizes is primarily a function of its vapor pressure.
      • Therefore, if one substance is said to be “more volatile” than another, this indicates that the former has a higher vapor pressure than the later at the same temperature.
  • Non-volatile 
    • indicates that the substance does not form a vapor, or has an extremely low vapor pressure, at room temperature.
      • This usually refers to solutes such as sodium chloride that do not contribute to the vapor pressure of a solution when dissolved in a solvent.
      • By contrast, something like methanol would have its own vapor pressure that would add to the vapor pressure of the solvent into which it is dissolved.

1

  • Phase changes
    • How is vapor pressure affected by temperature?

  • Increased temperature increases vapor pressure.
  • This makes logical sense for a few reasons.
    • First, increased temperature means the molecules of the liquid have a higher average kinetic energy.
      • This indicates that a larger fraction of those molecules will have the energy necessary to escape the intermolecular forces between liquid molecules to enter the gas phase.
  • One could also remember that liquids boil when their vapor pressure increases to the point that it equals atmospheric pressure.
  • It obviously requires an increase in temperature to cause a liquid to boil, therefore increasing temperature must increase vapor pressure. 

2

  • Solution Chemistry
    • Ways of measuring & describing solution [ ]
      • Compare Molarity with Molality
        • How do they both change with temperature?

  • Molarity, M
    • =moles solute/Liter solution
  • Molality, m
    • =moles solute/Kg solvent
  • Molarity (M) changes w/ temperature, but molality (m) DOES NOT!
  • Provide a possible explanation...
    • Hint: what happens to volume when temperature increases? (increases. Mass doesnt change though)

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3

  • Galvanic Cells
    • What is used to create a Galvanic Cell?
    • What will a correctly set up galvanic cell always produce?
    • Describe direction of electron flow in a galvanic cell

A functioning Galvanic cell can be created using any two metals, regardless of their reduction potentials.

  • If a galvanic cell is properly set up, it will always produce current.
  • Electrons will automatically flow from the species with the lower reduction potential to the species with the higher reduction potential. 

LO⇒HI

3

  • Solution Chemistry
    • Ions in Solution
      • What do Na+ and Cl- look like in water?
      • In specific, how are the dipoles oriented? 

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4

RAOULT'S LAW 

(Solving for vapor pressure of a solvent + impure solute)

VpTOTAL=XVp

How do you solve for the mole fraction,"X?" 

Given: grams of solvent and solute

Once you find X, how do you solve for VpTOTAL?

Given: grams of solvent &solute

  • Find respective # of MOLES of both

Add them together=TOTAL # of MOLES

  • Divide # of INDIVIDUAL MOLES of solvent & solute
    • ...by TOTAL # of MOLES

THIS IS YOUR MOLE FRACTION, X

SOLVENT & SOLUTE COMBINED MUST EQUAL 1.0!!

  • Now, simply multiply X's for solvent & solute by their respective vapor pressures, Vp

ADDING THESE TOGETHER GIVES: VpTOTAL 

5

  • Solution Chemistry
    • "like dissolves like" refers to? 
    • Describe Ksp (3 points)

  • “Like dissolves like” refers to the fact that polar substances are soluble in polar solvents and non-polar substances are soluble in non-polar solvents.
    • Polar and non-polar substances DO NOT form solutions.
  • The Solubility Product Constant, Ksp
    • Exactly the same thing as Keq, Ka, and Kb.
    • Like all the other K's, remember the following:
    1. Leave out pure liquids and pure solids (this will make all Ksp equations only have a numerator - if you have something in the denominator of a Ksp equation, you’ve made a mistake). 
    2. Temperature is the ONLY THING that changes Ksp 
    3. Ksp can ONLY be observed in a saturated solution.
      • This is because saturation is the point at which the dissolution reaction has reached equilibrium.
      1. In other words, it’s just like all other equilibrium constants—you cannot measure them anywhere other than at equilibrium.

7

  • Electrochemistry
    • Electrical Potentials (Eº)
      • What do they tell us? 
      • Units for Eº
      • What is the most common "half-reaction" on the MCAT? 
      • What are the only half-rxns on the MCAT that DONT begin with ^this?^

  • tell us the degree to which a species "wants electrons," or "wants to be reduced"
  • These potentials are given in Volts are are always presented in what is called a "half-reaction"
    • ex: Ag2+(aq) + 2e- ⇒ Ag(s) 
  • Most common half-reaction is an aqueous metal ion being reduced to form the associated solid metal
    • ONLY half-rxns that DONT begin with a metal cation: 
      • O2, H2O, H+

7

  • Ideal Gas Law
    • Kinetic Theory of Gases
      • What are the 2 main ideal gas "assumptions?" (describe them)
      • Name the other 6 
      • What is the only measurement of the molecules THEMSELVES that we ever consider?

  1. Gas molecules themselves are of negligible volume compared to the volume occupied by the gas.
  2. All intermolecular forces between gas molecules are negligible.
    • Focus on the first two assumptions; they are responsible for most of the differences between what PV=nRT predicts and how real gases actually behave.
    • To simplify things even further: THINK OF IDEAL GAS MOLECULES AS HAVING:
      • NO volume and NO intermolecular forces.
  3. All collisions between gas molecules are perfectly elastic 
  4. Gases are made up of a large number of molecules that are very far apart from one another 
  5. Pressure is due to collisions between gas molecules and the walls of the container 
  6. All molecular motion is random 
  7. All molecular motion follows Newton’s laws of motion
  8. The average kinetic energy (KE) of gas molecules is proportional to temperature
  • Only measurement of the molecules we consider: THE NUMBER OF MOLES !!!!

7

  • Gases
    • The Ideal Gas Law
      • What is "STP?" 
      • Give the values for P,V,n,R,&T

  • Standard Temperature & Pressure (STP) 
  • A set of standard conditions true of any ideal gas said to be “at STP.”
    • For the MCAT, unless you are specifically told otherwise, assume that all gases are ideal and start out at STP.
  • At STP the variables in the Ideal Gas Law are defined as follows: 
    • P = 1 atm 
    • V = 22.4 L 
    • n = 1 mole 
    • R = 0.0821 L*atm/mol*K or 8.31 J/mol*K 
    • T = 273 K (0°C)

8

  • Gases
    • The Ideal Gas Law=?

  • PV = nRT
    • R = 0.0821 L*atm/mol*K or 8.314 J/mol*K

8

  • Gases
    • Real Gases
      • Van der Waals Equation
      • What are the only 2 things you need to know about the VDWs equation? 

  • [P+a'(n/2V)2]*[(V/n)-b']=RT
    • a'
      • is a constant that represents the actual strength of the IM attractions
    • b'
      • is a constant that represents the actual volme of the molecules 
  1. Increased intermolecular attractions (a’)
    1. decrease pressure in real gases.
    2. The larger a’ is, the larger the second term will become and therefore the smaller P will be.
  2. Increased molecular volume (b’) increases volume in real gases.

9

  • Solution Formation
    • Describe the difference b/t a (+) and (-) ΔHsolution 

NEGATIVE  

exothermic and heat will be released

POSITIVE ΔHsolution=

A positive heat of solution means that energy must be ADDED to the system to make the solute dissolve.

 

For a solution to form, the intermolecular forces between the solute particles must first be broken;

  • then any intermolecular forces between the solvent particles must be broken
    • (to make room for the solute) 
  • Finally, new intermolecular forces are formed between the solute particles and the solvent particles.
  • If the new intermolecular forces formed are greater (i.e., stronger, more stable) than the sum of the intermolecular forces that had to be broken, net energy is released and the solution is said to have a negative Heat of Solution (ΔHsolution
  • This means that the dissolution process is exothermic and heat will be evolved.
  • If the new intermolecular forces are not more stable than the old ones, the solution has a positive ΔHsolution.
    • A positive heat of solution means that energy must be added to the system to make the solute dissolve

11

  • Electrochemical cells
    • Galvanic Cells
      • What does the salt bridge do?

  • There will be a buildup of negative charge in the copper vessel due to continual loss of copper cations
  • and a buildup of positive charge in the zinc vessel due to the continual production of zinc cations. This polarity resists the flow of electrons and would eventually shut down the cell if a salt bridge were not present.
  • Within the salt bridge sodium ions can flow toward the copper vessel and nitrate ions can flow toward the zinc vessel, neutralizing the buildup of charge and allowing electron flow to continue.
  • The metal cations themselves, as well as any other ions in the solutions, can also flow through the salt bridge.
  • In an electrical sense, the salt bridge connects the circuit, allowing continual flow of electrons from electrode to electrode and then back through the salt bridge via ion diffusion.

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11

  • Phase changes
    • Heating Curves 
      • Describe them

  •  A graph of temperature (T) in Kelvin or Celsius vs. heat (q) in Joules.
  • Occasionally, time is graphed on the x-axis instead of heat
    • (if heat is added at a constant rate the temperature vs. heat graph and the temperature vs. time graph look approximately the same).  

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12

Phase Changes

  • Describe Osmotic Pressure
    • What side RECEIVES water? 
    • More___=More OSMOTIC PRESSURE

  • A measure of the tendency of water to move from one solution to another across a semi-permeable membrane
  • Usually represented by the capital Greek letter pi, Π.
  • It is the side that will receive the water via osmosis that has the higher osmotic pressure.
  • In other words, more solute means more osmotic pressure.

Π= iMRT

  • i = # of ions formed in solution
  • M is the solute molarity
  • R is the gas constant
  • T is the absolute temperature

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13

The Ideal Gas Law

  • STP

Define Absolute Temperature

  • What is it measure relative to?
    • What's happening on a molecular level at the point Abs. Temp. is relative to?

An absolute temperature is any temperature measured relative to absolute zero

  • The Kelvin scale is measured relative to absolute zero
    • where absolute zero is defined as 0 degrees Kelvin
      • ∴ all Kelvin temperatures are absolute
  • Absolute zero is a theoretical temperature limit where all molecular motions cease

14

  • Solution Formation
    • What happens to entropy when a solution forms from a solute? 

  • The dissolution of a solute into a solution is accompanied by a very large, positive change in entropy.
  • A solid or crystal is highly ordered and the break-up and solvation of that solid into individual molecules represents a significant increase in disorder

15

  • Redox Rns
    • Define Oxidation State

  • Is the apparent charge that an atom takes while in a molecule
  • The sum of oxidation states for all the atoms in a molecule must equal the charge of that molecule
    • or equal zero if the molecule is neutral

15

  • Gases
    • The Ideal Gas Law
      • The Kinetic Theory of Gases
        • Describe it in general

  •  Theoretical model used by scientists to study and predict the “ideal” behavior of gases. 

16

  • Phase Changes
    • Define a phase change

  • In common use, the term “phase” is used to distinguish between the solid, liquid, and gas forms of a substance.
  • Solid, liquid and gas are more correctly called “states” (a.k.a., “states of matter”).
      •  

17

  • Gases
    • Dalton's Law of Partial Pressures 

  • Ptotal=P1+P2+P3....
  • This seems pretty straightforward; the sum of the partial pressures equals the total pressure.
  • However, if you gloss over this you will miss a very important point.
    • If we add more of Gas 1 (P1) to an existing mixture of three gases, we have increased the partial pressure of Gas 1 and the total pressure, but have had zero effect on the partial pressure of the other gases.
    • Partial pressure is NOT similar to mole fraction or mass percent.
    • By adding more of Gas 1 we did decrease the mole fraction and the mass fraction of Gases 2 and 3, but we did NOT decrease their partial pressures.

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18

  • Phase Changes
    • Heating Curves
      • What do horizontal portions of the graph represent?
      • What does the slope of the various non-horizontal sections of the graph represent? 

  • The horizontal sections of the graph represent phase changes.
    • The first flat section will represent the phase change between solid and liquid and the second will represent the phase change between liquid and gas.
    • If heat is on the x-axis then the length of the first horizontal section represents the heat of fusion and the length of the second horizontal section represents the heat of vaporization.
  • The slope of the lines between these horizontal sections represents the inverse (∆T/Q) of heat capacity (Q/∆T) for that particular phase of the substance.
  • One should observe, therefore, that different phases of the same substance usually have different heat capacities—as indicated by the differing slopes of those sections of the following graph:

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18

Phase Changes

  • Vapor Pressure

Raoult's Law

  • What does Raoult's Law solve for?
  • Give the formula for vapor pressure w/ a NON-Volatile solute

SOLVES FOR VAPOR PRESSURE OF A SOLVENT THAT HAS A SOLUTE ADDED TO IT

VpTotal = XV  

 

Vapor Pressure w/ a Non-Volatile Solute =

(mole fraction of the pure solvent, X)

X

 (Vp of the pure solvent, Vp°)

 

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19

  • Phase Changes
    • Draw a phase diagram. Label:
      • axes
      • solid, liquid,  gas regions
      • critical point
      • triple point
    • Describe what's going on at the critical & triple points wrt phases

  • The triple point 
    • the precise temperature and pressure at which all three phases (i.e., states) exist simultaneously in equilibrium with each other.
  • The critical point 
    • the precise temperature and pressure above which liquid and gas phases become indistinguishable.
      • At this point liquid and gas phases cease to exist, merging into a single phase called a supercritical fluid.
        • This supercritical fluid cannot be compressed back into the liquid phase by increasing pressure, nor can it be turned into a gas by increasing temperature.
  • The critical temperature and the critical pressure are simply the temperature and pressure at the critical point.
  • At the triple point all three phases are present.
  • At the critical point none of the original three phases are present, only the new “supercritical fluid” phase.

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21

  • Gases
    • Solving Ideal Gas Law Problems
      • Describe the 2 methods for approaching these problems 

  1. Manipulating Equations:
    • Compare the variables for any system to the variables outlined above for STP.
    • For any variation found, estimate the factor by which that variable has changed and use the manipulating equations skills to predict the effect that change will have on the unknown variable
  2. P1V1/T1 = P2V2/T2 
    • (a.k.a., The Combined Gas Law) 
    • Because PV/T = nR and R is a constant, for the same number of moles of gas the ratio of PV/T must remain constant regardless of the changes made to the system.
    • You can choose the first set of data as being STP, or as any other point where P, V and T are known.
  • The second set of data will be different, but the ratio will always be the same.
    • Plug in the data and solve for the unknown.
  • Conceptually, you’re probably better off if you understand and can apply the first method.
  • The first method is also much faster!

21

  • Solution Chemistry
    • Ions in solution
      • Define:
        • Solvation
        • Hydration
        • Hydration Number
        • Hydrate
        • Anhydrous
        • Aqueous 

  • Solvation
    • a general term for the process wherein solvent molecules surround a dissolved ion or other solute particle creating a shell.
  • Hydration
    • a specific kind of solvation wherein water is the participating solvent.
      • Water molecules, being polar, can surround both negatively and positively charged solutes by directing either their partially-negative oxygen, or partially positive hydrogen, moieties toward the ion.
  • The hydration number
    • is the number of water molecules an ion can bind via this solvation process, effectively removing them from the solvent and causing them to behave more like an extension of the solute.
  • A hydrate
    • an inorganic compound in which water molecules are permanently bound into the crystalline structure.
      • The nomenclature of a hydrate is altered to reflect the presence of water molecules.
      • For example, anhydrous cobalt(II)chloride contains no water, but cobalt(II)chloride hexahydrate [CoCl2∙6H2O] contains six water molecules complexed with each cobalt.
        • As we see in these two names, the term anhydrous is often applied to a compound that can form complexes with water to differentiate molecules that do not contain water from those that do.
  • Aqueous
    • efers to any solution for which water is the solvent.

22

Phase Changes

  • Freezing Point Depression
    • Freezing point of a liquid is depressed when a ____ _____ is added according to WHAT FORMULA?

  • The freezing point of a liquid is depressed when a non-volatile solute is added according to: 

∆T = kfmi 

  • kis a constant (different than kb

23

  • Electrochemical cells
    • Concentration cells
      • Nernst Equation
      • For a concentration cell, Eo will always be...?

  • E=Eº-(.06/n)*log [lower]/[higher]
    • n=moles of electrons transferred
      • ex: Ag+(aq)⇒Ag(s)
        • 1 electron transferred 
  • Eo will always be ZERO

24

  • Phase changes
    • How does the addition of a non-volatile solute affect vapor pressure?
    • What about a volatile solute?

BOTH DECREASE Vp

Non-volatile solute

Addition of a non-volatile solute DECREASES vapor pressure.

  • The non-volatile solutes in solution occupy a portion of the limited surface area available for vaporization.
  • Liquid molecules must be at the surface of the liquid in order to escape into the gas phase

Volatile solute

  • When a volatile solute is added to a solvent it usually decreases vapor pressure for the same reason that a non-volatile solute decreases vapor pressure.
  • As long as the vapor pressure of the solute is LESS THAN the vapor pressure of pure solvent:
    • addition of the volatile solute will decrease vapor pressure.

However, if a solute is added that has a vapor pressure greater than that of the pure solvent:

then the vapor pressure of the solution will actually be higher than that of the pure solvent

26

  • The ratio given by PV/nRT tells us which of two assumptions is the major cause of the deviation from Ideal Gas Law behavior
    • What are the 2 assumptions? 
    • Hint: If PV/nRT is > or

  1. If PV/nRT > 1 it is due mostly to the molecular volume assumption
  2. If PV/nRT it is due mostly to the intermolecular forces assumption

28

  • Gases
    • The greatest deviation between ideal gas behavior and real gas behavior occurs when either: (2)
      • Describe why they deviate from the IGL

  1. the temperature is extremely low, or
    • At very low temperature, the interaction between gas molecules (which is not really zero, as assumed in the ideal gas law) starts to become more important. 
  2. the pressure is extremely high
    • ​At very high pressure the molecules are pushed close together and their actual size (assumed to be zero) becomes comparable to the distance between them.
  • In either case, there is a deviation from ideal behavior, but with opposite effects:
    • at high pressure, gas molecules will occupy a greater volume;
    • and at low temperature, they will produce a smaller pressure than predicted by the ideal gas law.
    • These effects are better represented in the Van der Waals equation.

28

  • Solution Chemistry
    • Calculating Solubility
      • 4 steps

  1. Write out the Ksp expression 
  2. Substitute into the expression the value given for Ksp
  3. Substitute a factor of x into the equation for the concentration of each ion, using 2x, 3x, etc., if more than one mole of each ion is produced
    • Hint: Ask yourself, “If x moles of the reactant are dissolved, how many moles of each ion will be produced?”
  4. Solve for x
    • Your answer, “x” is the “solubility” of that particular specie.

29

Oxidation-Reduction (REDOX) Rxns

  • Define
    • What does it mean when something is "oxidized?"
    • Define Reducing Agent
    • Define Oxidizing Agent 

  • a rxn where one or more electrons are transferred from one atom to another 
  • When an atom loses e's, it is "oxidized" (Gains=Reduced) 
  • Reducing Agent
    • is an atom or molecule that donates electrons to another atom or molecule and is itself oxidized in the process
  • Oxidizing Agent
    • is an atom or molecule that accepts electrons and is itself reduced in the process

30

  • Solution Chemistry
    • ALL cpds containing ___ are soluble?
    • ALL cpds containing ___ are insoluble?

  1. All compounds containing the following are SOLUBLE:
    1. nitrate, NO3-
    2. ammonium, NH4+
    3. all alkali metals (Group IA)
  2. All compounds containing the following are INSOLUBLE:
    • (unless paired with something from the “always soluble” list above)
      1. carbonate, CO32-
      2. phosphate, PO43-
      3. silver, Ag
      4. mercury, Hg
      5. lead (Pb)

31

  • Phase Changes

    • Molecules of the same “phase" are/have?  (3) 

       

  • a) are in the same state, 
  • b) have the same chemical composition, and 
  • c) are structurally homogenous 
    • e.g., In its solid state, carbon can exist as either diamond 

32

  • Gases
    • Ideal Gas Law
      • What are Standard Conditions?
      • What is the difference b/t STP and "Standard Conditions?" 

  • Standard Conditions is a set of agreed-upon conditions at which thermodynamic data, reduction potentials, or other standardized data are measured.
  • STP and standard conditions are NOT interchangeable.
  • As an example, temperature at STP is 0°C and temperature at standard conditions is 25°C.

33

  • The relationship b/t Gibb's Free Energy & Chemical energy
    • Gibb's Free Energy eqn=?
      • what do the components of this eqn teach us? 
    • What does Faraday's Constant represent? 

  • ∆G° = -nFE°
    • n is the number of moles of electrons transferred in the balanced redox reaction
    • F is Faraday’s constant.
  • From the above EQN we learn:
    • + E˚ = negative
    • ∆G = spontaneous reaction
  • Faraday’s Constant = the charge on one mole of electrons. 

34

  • Ways of measuring & describing solution [ ]
    • Define:
      • mole fraction
      • mass %
      • ppm
      • normality

  • Mole fraction
    • moles solute/total moles solution (solute + solvent) o
  • Mass %
    • mass solute/total mass of solution * 100 o
  • ppm
    • mass solute/total mass solution * 106 (for ppb multiply by 109 )  
      • Parts per million (ppm) is NOT a measure of how many solute particles there are per 1 million total particles.
      • This is how most students erroneously think of it
      • It is nothing more than mass percent multiplied by 104 , or “mass fraction” multiplied by 106 .
      • The purpose of multiplying by 1 million is to make very, very small concentrations easier to work with.  
      • ppm = mg/Kg = mg/L (since 1 L of water has a mass of 1 Kg) 
  • normality
    • # of moles of equivalents/Liter solution. 
    • For example, A 1M solution of H2SO4 can be referred to as a “2 Normal” solution because it produces two moles of hydronium ions per Liter of solution.
    • By the same token, a 2M solution of H3PO4 3- would be a “6 Normal” solution with respect to hydronium ions.
    • This concept ignores the decreasing acidity of each proton.

35

  • Phase Changes
    • Liquid⇒Gas=?
    • Gas⇒Liquid=?
    • Describe ΔHformation

  • Liquid⇒Gas
    • Evaporation
  • Gas⇒Liquid
    • Condensation
  • ΔHformation
    • The amount of energy (in Joules/mol) required to go from liquid to gas OR
    • the energy that must be removed to go from gas to liquid.
    • Again, it describes both evaporation and condensation.

36

  • Solution Chemistry
    • The Common-Ion Effect
      • What will happen if a common ion is added to a saturated solution?
      • What would happen to a spectator ion in the same solution?

  • Addition of a common ion will cause precipitation

 

  • ​If a spectator ion is added no precipitation will result

37

  • Solution Chemistry
    • Define: 
      • Solvent
      • Solute
      • Colloid
        • How do they differ from solutions? 
        • Give 2 examples of colloids

  • Solution 
    • a homogenous mixture of two or more compounds in the same phase.
    • (We usually think of all solutions as being in the liquid, or “aqueous” phase; however, a homogenous mixture of gases is also called a “solution”)
    • is the substance dissolved INTO the solvent
    • Thus, the SOLVENT is more abundant than the SOLUTE 

  • Colloids
    • are NOT solutions!
    • are solvents containing undissolved solute particles that are:
      • too small to be separated by filtration, but are
      • much larger than the solute particles in a true solution.
    • Colloids scatter light, while true solutions do not.
    • Examples of colloids include paint (a suspension of solid crystals in a solvent) and dust floating in air

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38

  • Solution Chemistry
    • Solubility
      • Define:
        • solubility
        • precipitate
        • saturated solution
        • unsaturated solution
        • super-saturated solution

  • Solubility
    • the amount of a solute that will dissolve in a given solvent at a given temperature.
    • Temperature is usually specified because for most solids dissolved in liquids, solubility is directly related to temperature.
    • On the MCAT, solubility is usually measured in either g/mL, g/100mL, or mol/L. A
  • Precipitate
    • a solid formed inside of a solution as the result of a chemical reaction, such as the common ion effect.
    • Precipitates only form when the ion product exceeds the solubility product constant, Ksp.
    • For example, given the dissolution of iron(III)chloride in water [Equation: FeCl3(s)  Fe3+(aq) + 3Cl- (aq)], if NaCl is added to the solution Le Chatelier’s principle predicts that the reaction will shift to the left, reforming the solid.
  • Saturated solution
    • a solution that contains the maximum amount of dissolved solute it can hold.
    • For a saturated solution the ion product equals the Ksp.
  • Unsaturated solution
    • any solution that contains less than its maximum amount of dissolved solute.
    • For unsaturated solutions the Ksp is greater than the ion product.
  • Super-saturated solution
    • a solution that contains more dissolved solute than predicted by the solubility product constant
      • in other words, the ion product exceeds the Ksp WITHOUT a precipitate forming.
    • Supersaturated solutions usually form only when a solution is held at a higher temperature during dissolution (at which Ksp would be larger) and then slowly cooled to a temperature at which Ksp is smaller.

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39

  • Solution Chemistry
    • Clarify the difference between "solubility" and "Solubility Product Constant, Ksp"

  •  Solubility is a measure of “how much” of a solute can be dissolved in a given solute.
    • For example, the solubility of iron(III)chloride in water is 74.4g/100mL.
  • The solubility product constant, or Ksp, is defined as the product of the dissolved ions in a saturated solution (i.e., at equilibrium) raised to their coefficients in the balanced equation

41

  • Phase Changes
    • Solid⇒Gas=?
    • Gas⇒Solid=?

  • Solid⇒Gas
    • Sublimation
  • Gas⇒Solid
    • Deposition

42

Phase Changes

Gas Solubility

  • How does solubility for a GAS dissolved in a LIQUID compare to a SOLID dissolved in a LIQUID?
    • How does an increase & decrease in TEMPERATURE & Vp affect solubility of Gases dissolved in  liquids?
  • Polar & Non-polar gases easily form WHAT?
    • ​picture attached 

The solubility of gases in liquids follows a trend that is

EXACTLY OPPOSITE

of the solubility of solids in liquids

Temperature: 

  • For GASES dissolved in liquids:
    • increased temperature
      • decreases solubility
    • decreased temperature
      • increases solubility

Vapor Pressure, Vp

  • Increasing the vapor pressure of gas X over a liquid: 
    • increases the solubility of gas X in that liquid

(This is why they pressurize soda pop

cans with excess CO2)

Polar and non-polar gases easily form: 

HOMOGENOUS MIXTURES (below)

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  • Phase changes
    • what 2 quantities are equal when a liquid boils? 

  • Atmospheric Pressure 
  • Vapor Pressure 

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  • Electrochemistry
    • Cell Potential (Eº)
      • Can you add two Eº values directly from the half-reduction table?  
      • One mole of Cu2+ has ___ reduction potential as 2 moles of Cu2+

  • NO! 
    • These are all reduction 1/2 rxns and you need one of each:
      • one reduction
      • one oxidation
    • You must reverse the half rxn of the species with the lowest reduction potential and take the negative of its Eº value
  • The SAME reduction potential
    • i.e., DONT use stoichiometry 

46

  • Electrochemical Cells
    • Galvanic Cells
      • Where do reduction and oxidation occur? 
      • Cathode= +/-?
      • Anode=+/-?
      • The above 2 are true in galvanic cells, but not ___ cells

  • Reduction @ Cathode
    • Cathode = (+)

  • Oxidation at Anode
    • Anode= (-) 
  • True in galvanic cells only, NOT electrolytic cells!

47

  • The relationship b/t Gibb's Free Energy & Chemical energy
    • Faraday vs. Farad

  • A Faraday is an obsolete unit of charge equal to the charge on one mole of electrons.
    • In other words, Faraday’s constant = 1 Faraday.
    • The Faraday has since been replaced by the Coulomb.
  • A Farad is a unit of capacitance.
    • It is a “summary” unit similar to a Newton.
      • Just as we can say 1 Newton instead of saying 1 Kg*m/s2 , we can say 1 Farad instead of saying 1 C2*s2 /m2*kg.
    • A Farad is the amount of capacitance necessary to hold 1 C of charge on a capacitor with a potential difference of 1 Volt.

48

Phase Changes

  • Heating Curves
    • What does the SLOPE of the various NON-HORIZONTAL sections of the graph represent? 

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The SLOPE of the lines between these horizontal sections represents:

  • THE INVERSE (ΔT/Q) of HEAT CAPACITY (Q/ΔT
    • ...for that particular phase of the substance

Therefore,  DIFFERENT phases of the same substance usually have DIFFERENT heat capacities!

as indicated by the differing slopes of those sections of the following graph (IE GAS HAS A STEEPER SLOPE) 

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  • Phase Changes
    • What main idea should we remember about Henry's Law? 

  • the solubility of a gas IN a liquid is directly proportional to the partial pressure of that gas OVER that liquid

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  • Electrochemical cells
    • Electrolytic Cells

  • is essentially a galvanic cell to which an external voltage is applied 
  • Oxidation @ anode, Reduction @ cathode
  • Cell Potential (Eº) always NEGATIVE
  • The sum of the externally applied voltage (Vbattery) and the negative cell potential ( -E ° cell) must be positive.
  • Cathode = (-); Anode = (+)
  • Note the difference compared to galvanic cells. 

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  • Gases
    • Effusion and Diffusion (Graham's Law)
      • Compare Effusion with Diffusion
      • Give Graham's Law equation
        • explain it

  • Diffusion:
    • The process by which gas molecules spread from areas of high concentration to areas of low concentration due to the random motion imparted to them as a result of their kinetic energy and collisions with other molecules. 
  • Effusion:
    • The diffusion of gas particles through a pin hole.
      • A pin hole is defined as a hole smaller than the average distance a gas molecule travels between collisions
  • E1/E2 = √MW2/√MW1  
    • E1 and E2 can represent either the effusion rate or the diffusion rate of gases 1 and 2, respectively.
  • Notice that the rate of effusion or diffusion is inversely proportional to the molecular weight of the gas.

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  • Solution Chemistry
    • Solubility
      • The Ion Product

  • Also referred to as the “Solubility Product.”
  • The ion product has the same relationship to Ksp as Q does to Keq.

 

  • Plug in the values for the actual concentrations of each species at some point other than equilibrium (i.e., for an unsaturated or supersaturated solution).
    • If the product is greater than Ksp, you know a precipitate will form.
    • If it is less than or equal to Ksp, then you know that no precipitate will form.
    • If it is exactly equal to Ksp, then the solution must be exactly saturated (i.e., at equilibrium).

55

Phase Changes

  • Boiling Point elevation
    • The boiling point of a liquid is elevated when a ______ _____ is added according to WHAT FORMULA?

  • The boiling point of a liquid is elevated when a non-volatile solute is added according to:
    • ∆T = kbmi
      • where kb is a constant,
      • m is MOLALITY (NOT molarity
      • i is the number of ions formed per molecule
        • a.k.a. The Van’t Hoff Factor;
          • Ex: For NaCl i = 2; for CaCl2 i = 3

56

  • Phase changes
    • Vapor Pressure 
      • Define

  •  Vapor Pressure (Vp) is the partial pressure of the gaseous form of a liquid that exists over that liquid when the liquid and gas phases are in equilibrium.

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  • Electrochemical cells
    • The Galvanic (or "Voltaic") Cell
      • Define
      • How do they work?

  • Galvanic cells convert energy into electrical energy
    • ​By taking advantage of the difference in reduction potentials b/t two metals, a current can be spontaneously generated along a wire that connects 2 metal electrodes submerged in solutions that contain metal ions 

 

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  • Phase Changes
    • Solid⇒Liquid=?
    • Liquid⇒Solid=?
    • Describe ΔHfusion

  • Solid⇒Liquid=Melting
  • Liquid⇒Solid=Freezing 
  • ΔHfusion
    • The amount of energy in Joules/mole required to go from solid to liquid, OR
    • the energy that must be removed to go from liquid to solid.
    • This describes the transition in both directions (i.e., melting ⇔ freezing) 

60

Electrochemical cells

  • Concentration cells
    • How are they set up?
    • What E° cell value do they ALWAYS have?
    • They have a POSITIVE reduction potential,E, if there is a DIFFERENCE in the _______s of the two solutions

    • The Nernst Equation (below) is used to calculate WHAT?

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A special type of GALVANIC cell

  • The SAME (!) electrodes and solution are used in both beakers
  • In one beaker the metal is oxidized via its oxidation half-reaction, and in the other beaker it is reduced via its reduction half-reaction.

Because the reduction potentials of oxidation and reduction half reactions for the SAME species only differ by the SIGN of E°

E°cell = 0.00V

  • It appears that nothing would happen
  • Concentration cells are therefore NONstandard conditions by definition

They have a POSITIVE reduction potential E  if:

  • there is a difference in the MOLARITIES of the two solutions.

The Nernst equation is used to calculate the cell potential based off of the E˚ of the species and the two concentrations. 

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  • Solution Chemistry
    • The Common-Ion Effect

  • Is a specific application of Le Chatelier’s principle to solution chemistry

​Example:

  • Consider the dissolution of Iron(III)Chloride in water:

FeCl3(s)⇒Fe3+(aq) + 3Cl- (aq)

  • Suppose that enough solute is added to saturate the solution.
  • If sodium nitrate, NaNO3 is then added to this solution:
    • it would have NO effect.
  • However, if NaCl were added, the presence of extra chlorine ions from NaCl would—according to LeChatelier’s Principle—drive the reaction to the LEFT
    •  resulting in precipitation.

In this example, chloride is considered a “common ion” and the precipitation as a result of its addition is what is referred to as the “Common Ion Effect.”

  • Other ions, such as sodium and nitrate—that do not shift the equilibrium—are considered “SPECTATOR ions”

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  • Ideal Gas Law
    • More limited versions of IGL 
      • Name the 2
        • What do they assume? 
      • Just focus on ideal gas law instead of these

  • Boyle's Law
    • P1V1=P2V2
      • Assumes constant temperature
  • Charles' Law
    • V1/T1=V2/T2
      • assumes constant pressure 

63

  • Electrical Potentials
    • What is the "standard" half-rxn against which all other half-rxns are compared? 
      • What is its Eº value?
    • What does the º signify on Eº? 
    • If something has a (+) Eº, it's more likely to...?
    • What about with a (-) Eº?

  • Hydrogen Half-Cell
    • 2H+ 2e- ⇒ H
  • Eº=0.00V
  • º = happens at standard conditions

 

  • Positive Eo
    •  MORE LIKELY to gain electrons (be reduced) than are hydrogen ions
  • Negative Eo
    • LESS LIKELY to gain electrons than are hydrogen ions

64

  • Phase changes
    • heating curves
      • How do you calculate ΔHfusion and ΔHformation from a heating curve? 
    • There is  NEVER _____ during a phase change?

  • ΔHfusion
    • The change in Q (x-axis) during the phase change from solid to liquid
  • ΔHformation
    • The change in Q (x-axis) during the phase chage from liquid to gas
  • ​There is never a CHANGE IN TEMPERATURE during a phase change!!!!!
    • The fact that heating curves are flat (i.e., horizontal, slope = 0) during the actual phase change demonstrates the following frequently-tested MCAT principle:
      • Once a phase change starts, all of the energy goes into breaking intermolecular forces and none goes toward an increase in temperature.

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  • Cell Potential (Eº cell) 
    • Define
      • What 2 things should you remember about half rxns?

  • Is the sum of the two half-rxns that make up an electrochemical reaction 
    1. Half rxns always come in pairs: one reduction and one oxidation
    2. The oxidation half rxn is the reverse (negative) of the reduction half rxn

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  • Phase changes
    • Henry's Law

  •  

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  • Redox Rnxs
    • List the oxidation state for:
      • any elemental atom
      • Fluorine
      • Hydrogen
      • Hydrogen w/ a metal
      • Oxygen (except peroxides)
      • Alkali metals
      • Alkaline earth metals
      • Group V atoms (V,Nb,Ta,Db)
      • Group VI atoms (Cr,Mo,W, Sg)
      • Group VII atoms (Mn,Te,Re,Bh)

  • Any elemental atom
    • 0
  • Fluorine
    • -1
  • Hydrogen
    • +1
  • Hydrogen w/ a metal
    • -1
  • Oxygen (except peroxides)
    • -2
  • Alkali metals
    • +1
  • Alkaline earth metals
    • +2
  • Group V atoms (V,Nb,Ta,Db)
    • -3
  • Group VI atoms(Cr,Mo,W, Sg)
    • -2
  • Group VII atoms (Mn,Te,Re,Bh)
    • -1

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Phase Changes

  • Vapor Pressure

Raoult's Law

  • What does Raoult's Law solve for?
  • Give the formula for TOTAL Vapor Pressure for a Solvent mixed with a VOLATILE solute

Raoult's Law determines the vapor pressure of a mix in which a solvent has an IMPURITY present

(X=mole fraction)

 

Total Vapor Pressure w/ a Volatile Solute =

(mole fraction of solvent) X (Vp° of the solvent)

+

(mole fraction of the solute) X (Vp° of the solute)

Vptotal =

Vpsolvent + Vpsolute

aka,  (Xsolvent Vp°solvent)

+

(Xsolute Vp°solute)

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