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Flashcards in Gen. Chem 4 Deck (38):
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Ideal Gas Law

PV= nRT

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

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Ideal Gas Assumptions for MCAT

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

THINK OF IDEAL GASES AS HAVING: NO VOLUME AND NO INTERMOLECULAR FORCES

2

Number of moles (n) of gas

The number of moles of gas is the only measurement of the molecules themselves we consider.

3

STP (Standard temperature and pressure)
- Assume that all gases are ideal AND start out at STP

Variables in the IDEAL GAS LAW at STP:

- 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 degrees C)

4

How to solve IDEAL GAS problems

1. Manipulating equations
2. P1V1/T1 = P2V2/T2 for the name number of moles of gas the ratio of PV/T must remain constant regardless of the changes made to the system.

5

Boyle's Law

P1V1=P2V2
Assumes constant temperature

6

Charles' Law

V1/T1=V2/T2
Assumes constant pressure

7

Van der Waals Equation

[P + a'(n/V)^2] * [(V/n) - b']= RT

a' is a constant that represent the actual strength of the intermolecular attractions
b' is a constant that represents the actual volume of the molecules

8

PV/nRT ratio tells us which of two assumptions is the major cause of the deviation from IDEAL GAS LAW behavior:

PV/nRT > 1 it is due mostly to MOLECULAR VOLUME assumption
PV/nRT < 1 it is due mostly to INTERMOLECULAR FORCES assumption

9

Dalton's Law of Partial Pressures

Ptotal=P1+P2+P3...
Adding more of gas 1 (P1) to an existing mixture of 3 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.

10

Effusion and Diffusion (Graham's Law)

Diffusion: process by which gas molecules spread from areas of high concentration to areas of low concentration due to random motion imparted to them as a result of their KE and collisions w/ other molecules

Effusion: diffusion of gas particle through a pin hole. Pin hole is defined as a hole smaller than the average distance a gas molecule travels between collisions.

11

E1/E2 = sqrt(MW2)/sqrt(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 (MW) of the gas

12

Phase

Molecules of the same "phase"
a) Are in the same state (solid, gas, liquid)
b) have the same chemical composition
c) Are structurally homogenous (e.g., in the solid state, carbon can exist as either diamond or graphite. These meet the first two criteria, but not the third, so they are different "phases" of carbon.)

13

Heating Curves

-Horizontal sections represent phase changes
-If heat on x-axis, then length of first horizontal section represents the heat of fusion and the length of the second horizontal section represents the heat of vaporization.
- Slope of the lines between horizontal sections represents the inverse (DELTA T/Q) of heat capacity (Q/DELTA T) for that particular phase.

14

Calculating DELTA H from a heating curve

There is no change in temperature during a phase change (that is why it's horizontal), all energy goes into breaking intermolecular forces and none goes toward an increase in temperature.

15

Vapor pressure

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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How is vapor pressure affected by temperature?

By increasing temperature, we increase vapor pressure. KE is higher in higher temperature so more molecules will have energy required to break intermolecular forces.

17

What two quantities are equal when a liquid boils?

When vapor pressure of the liquid equals atmospheric pressure

18

Raoult's Law

Vapor Pressure w/ a NONVOLATILE solute= (mole fraction of the pure solvent, X) * (Vp of the pure solvent, VP knot)
VP=XVPknot

Total Vapor Pressure w/ a VOLATILE SOLUTE= (mole fraction of solvent * Vpknot of the solvent) + (mole fraction of the solute * Vpknot of the solute). Vp total = Vp solvent + Vp solute = Xsolvent*Vpknotsolvent + Xsolute*Vpknot solute

19

Henry's Law

The solubility of a gas in a liquid is directly proportional to the partial pressure of the gas over that liquid

20

Gas solubility

Opposite of solubility of solids in liquids.
-For gases dissolved in liquids, increased temperature DECREASES solubility and decreased temperature INCREASES solubility
-Increasing the vapor pressure of a gas X over a liquid increases the solubility of gas X in that liquid
- Polar and non-polar gases easily form homogenous mixtures.

21

Boiling Point Elevation

Boiling point of a liquid is ELEVATED when a NON-VOLATILE solute is added according to
DELTA T = kbmi
kb= constant
m=molality
i=number of ions formed per molecule (VAN'T HOFF FACTOR)
T'T

22

Freezing Point Depression

The freezing point of a liquid is depressed when a NON-VOLATILE solute is added according to
DELTA T= kfmi
kf= constant (different than kb)

23

Osmotic Pressure

Tendency of water to move from one solution to another across a semi-permeable membrane. It is the side that will RECEIVE the water via osmosis that has the highest osmotic pressure.
Pi (greek symbol)= iMRT
i=Van Hoft factor M= molarity of solute R= gas constant T= absolute temperature

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Solvation

general term for the process wherein solvent molecules surround a dissolved ion or other solute particle creating a shell.

25

Hydration

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.

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Hydration Number

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.

27

Hydrate

inorganic compound in which water molecules are permanently bound into the crystalline structure.

28

Solution Formation

For a solution to form, the intermolecular forces between the solute particles must be first broken, then any intermolecular forces between the solvent particles must be broken. Finally, new intermolecular forces are formed between solute and solvent particles.

29

DELTA Hsolution

If the new intermolecular forces FORMED are greater (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 (DELTAHsolution < 0). This means that the dissolution process is exothermic and heat will evolve. If DELTAHsolution > 0 (new intermolecular forces NOT more stable), means that energy must be added to the system to make the solute dissolve.

30

Solubility

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.

31

Precipitate

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.

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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.

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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.

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Supersaturated solution

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.

35

The Ion Product

"Solubility Product", it 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. 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 the ion product happens to be exactly equal to Ksp, then the solution MUST be exactly saturated.

36

Calculating Solubility

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.
4) Solve for x. Answer "x" is the "solubility" of that particular species.

37

Common Ion Effect

specific application of Le Chatelier’s principle to solution chemistry. 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 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.”