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Chemistry 101 > Solutions > Flashcards

Solutions Flashcards

From a refresher on complex ions to calculating the molar solubility of silver chloride, use these cards to master the topic of chemical solutions as tested in most introductory undergrad chemistry courses and even on the AP Chemistry exam.

1
Q

Define:

a solution

A

A solution is a homogeneous mixture of two or more substances in a single phase.

Ex: NaCl dissolved into water creates a solution of Na+ ions, Cl- ions, and H2O all in one phase (liquid).

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2
Q

Define:

a solvent

A

A solvent is the substance whose phase remains after solvation (or the substance in excess).

Ex: dissolving a small amount of solid NaCl into liquid water produces a liquid solution. Water was the solvent.

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3
Q

Define:

a solute

A

A solute is the substance whose phase is lost after solvation (or the substance in scarcity).

Ex: dissolving a small amount of solid NaCl into liquid water produces a liquid solution. NaCl was the solute.

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4
Q

Ionic compounds dissolve readily in polar solvents; what ion form do metals usually take?

A

Metals typically become cations in solution.

Metals, which are found on the left side of the periodic table, mostly form cations by losing electrons to a nonmetal so that both can form stable octets.

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5
Q

Ionic compounds dissolve readily in polar solvents; what ion form do nonmetals usually take?

A

Nonmetals typically become anions in solution.

Nonmetals, which are found on the right side of the periodic table, mostly form anions by gaining electrons from a metal so that both can form stable octets.

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6
Q

What charge do the elements below usually take when forming ions in solution?

  • Li
  • Na
  • K
A

The alkali metals form cations with a single positive charge, +1.

Li, Na, and K are all alkali metals from column 1 of the periodic table.

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7
Q

What charge do the elements below usually take when forming ions in solution?

  • Br
  • Cl
  • F
A

The halides form anions with a single negative charge, -1.

Br, Cl, and F are all halides from column 7 of the periodic table.

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8
Q

Give the molecular form and charge for these common ions:

  1. ammonium
  2. chloride
  3. dichromate
  4. mercury(II) mercuric
  5. silver
A
  1. ammonium, NH4+, +1 charge
  2. chloride, Cl-, -1 charge
  3. dichromate, Cr2O7–, -2 charge
  4. mercury(II) mercuric, Hg++, +2 charge
  5. silver, Ag+, +1 charge
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9
Q

Give the molecular form and charge for these common ions:

  1. hydroxide
  2. barium
  3. sodium
  4. permanganate
  5. sulfite
A
  1. hydroxide, OH-, -1 charge
  2. barium, Ba++, +2 charge
  3. sodium, Na+, +1 charge
  4. permanganate, MnO4-, -1 charge
  5. sulfite, SO3–, -2 charge
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10
Q

Give the molecular form and charge for these common ions:

  1. hydrogen phosphate
  2. magnesium
  3. calcium
  4. bromide
  5. copper(I) cuprous
A
  1. hydrogen phosphate, HPO4–, -2 charge
  2. magnesium, Mg++, +2 charge
  3. calcium, Ca++, +2 charge
  4. bromide, Br-, -1 charge
  5. copper(I) cuprous, Cu+, +1 charge
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11
Q

Give the molecular form and charge for these common ions:

  1. sulfate
  2. nitrate
  3. peroxide
  4. hydronium
  5. iron (II) ferrous
A
  1. sulfate, SO4–, -2 charge
  2. nitrate, NO3-, -1 charge
  3. peroxide, O2–, -2 charge
  4. hydronium, H3O+, +1 charge
  5. iron (II) ferrous, Fe++, +2 charge
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12
Q

Define:

solvation

A

Solvation occurs when oppositely charged ends of polar solvent molecules surround solute ions.

Ex: water solvates the Na+ ion in the image below, creating a “solvation shell” around it.

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13
Q

Define:

hydration for solutions

A

Hydration is the process of solvation, where water is specifically used as the solvent.

Ex: because water is being used as the solvent in the image below, this can also be called a “hydration shell” of water molecules surrounding the Na+ ion.

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14
Q

Explain the “like dissolves like” rule.

A

Polar solvents readily dissolve polar solutes, while nonpolar solvents readily dissolve nonpolar solutes.

Ex: a non-polar solute such as naphthalene is insoluble in water, slightly soluble in methanol, and highly soluble in non-polar benzene.

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15
Q

Explain why exams like the AP Chemistry exam will rarely refer to an ion of H+ in aqueous solution? What ion will be used instead?

A

Because H+ is simply a proton in solution, it represents a very strong positive ion that water will form a hydration shell around.

Most commonly the ion H3O+ is used to represent the fact that a water molecule has bound to the free proton.

Rarely seen, but also possible, is H5O2+ (two water molecules sharing the proton) and H7O3+ (three water molecules).

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16
Q

Define:

solubility

A

Solubility is a measure of how much solute can be dissolved in the given solvent at a specific temperature and pressure.

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17
Q

In general, are the following soluble or insoluble in water?

  1. Nitrates (NO3-)
  2. Sulfites (SO3–)
  3. Acetates (CH3COO-)
  4. Chlorides (Cl-)
  5. Bromides (Br-)
A
  1. Nitrates, always soluble
  2. Sulfites, insoluble (except in group I and ammonium compounds)
  3. Acetates, soluble (except in silver compounds)
  4. Chlorides, always soluble
  5. Bromides, soluble (except in Silver, Lead, Copper, or Mercury compounds)
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18
Q

How soluble are the salts of alkali metal cations such as Li+, Na+, K+, etc.?

A

All salts of alkali metals are highly soluble.

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19
Q

How soluble are the salts of the ammonium cation, NH4+?

A

All ammonium salts are highly soluble.

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20
Q

What determines solubilities of the salts of alkaline earth and transition metal cations such as Ca+2, Mg+2, Fe+3, etc.?

A

Salts of alkaline earth metals and transition metals have solubilities that vary based on the anion that the cation is paired with in the salt.

Highly soluble anions like chloride (Cl-) will make soluble salts with alkaline earth and transition metals. Insoluble anions like phosphate (PO4-3) will make insoluble salts with these cations.

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21
Q

Which anions make salts that are always fully soluble?

A

The following anions make salts that are always soluble:

  • Nitrate (NO3-)
  • Chlorate (ClO3-)
  • Perchlorate (ClO4-)
  • Acetate (CH3COO-)
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22
Q

What is the solubility of salts made with the halogen anions Cl-, Br-, and I-?

A

Halogen anion salts are soluble unless the cation is Ag+, Pb2+, or Hg22+.

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23
Q

What is the solubility of salts made with the sulfate anion, SO42-?

A

Sulfate anion salts are soluble unless the cation is Ag+, Pb2+, Hg22+, Ca2+, Sr2+, or Ba2+.

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24
Q

Which anions make salts that are generally insoluble?

A

The following anions make salts that are generally insoluble:

  • Hydroxide (OH-)
  • Carbonate (CO3-2)
  • Phosphate (PO4-3)
  • Sulfite (SO3-2)
  • Chromate (CrO42-)
  • Sulfide (S2-)

Exception: salts with these anions paired with alkali metals (such as NaOH) or ammonium cations (such as (NH4)2CO3) will be soluble.

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25
Q

Describe what is meant by a saturated solution.

A

A saturated solution contains the maximum amount of solute that can be dissolved into the solvent at a particular temperature and pressure.

The solution is at equilibrium when fully saturated, so if more solute is added it will not dissolve (or there will be a precipitate formed).

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26
Q

A solute is being added to a solvent, and the solute is readily dissolving. During that time, what do we call the solution?

A

unsaturated

A solution containing less solute than needed for saturation is said to be unsaturated, and not yet at saturation equilibrium.

27
Q

A solution is somehow formed that contains more solute particles per solvent than should be possible at that temperature and pressure. What is that solution termed?

A

supersaturated

A solution containing more solute than needed for saturation is said to be supersaturated; it has exceeded the saturation equilibrium.

28
Q

Define:

precipitation

A

Precipitation is the reverse reaction of dissolution. Previously dissolved (solvated) salt ions bond together to form the original salt (solid).

Precipitation indicates that saturation has been exceeded at that temperature and pressure.

29
Q

Define and give the equation for:

molarity (M)

A

Molarity is a measure of the concentration of a solution, given in units of moles of solute dissolved per liter of solvent.

30
Q

How many moles of sodium chloride would 2 liters of a 5.0 M solution contain?

A

10 mol NaCl

Molarity = mol/L
5M = x?mol / 2L
x? = 10 moles

31
Q

Define and give the equation for:

molality (m)

A

Molality is a measure of the concentration of a solution, given in units of moles of solute dissolved per kilogram of solvent.

32
Q

58.5 grams of NaCl are dissolved in 2.0 kg of water at room temperature. What is the molality of the solution?

A

0.5 m solution of NaCl

molality = mol/kg
58.5 grams is the molecular weight of NaCl, so there is 1 mole present.
molality = 1mol / 2kg = .5m

33
Q

Define and give the equation for:

mole fraction (x)

A

Mole fraction is a measure of the concentration of a solution, given in units of moles of one component to total moles of all components.

34
Q

Define and give the equation for:

percent composition by mass (%)

A

Percent composition by mass is a measure of the amount of one component of a compound in grams compared to the total number of grams of all components.

35
Q

Define and give the equation for:

parts per million (ppm)

A

Parts per million is a measure of the concentration of one component of a mixture in kilograms compared to the total number of kilograms of all components of the mixture.

36
Q

What is the equation to calculate:

the change in molarity of a diluted solution

A

M1V1 = M2V2

Where:
M1 = initial molarity in mol/L
V1 = initial volume in L
M2 = final molarity in mol/L
V2 = final volume in L

37
Q

Define:

a colligative property

A

A colligative property of a mixture is one that only depends on the number of molecules of solute dissolved in the solvent, not the solute’s chemical nature.

Ex: A solvent’s vapor pressure, boiling point, freezing point, and osmotic pressure are all colligative properties.

38
Q

Explain the relationship between a liquid’s vapor pressure and its boiling point.

A

Boiling occurs when the vapor pressure above a liquid equals the surrounding pressure of the environment.

39
Q

Define and give the equation for:

mole fraction

A

Mole fraction is the ratio of moles of a particular substance present to the total moles of all substances present.

xA = nA / ntotal

where:
xA = mole fraction of A
nA = number of moles of A
ntotal = total number of moles of all substances present

40
Q

Define:

Raoult’s Law

A

Raoult’s Law: Ptotal = P0AxA + P0BxB + …

The partial pressure of a liquid above a mixture is equal to the vapor pressure of the pure liquid times its mole fraction in the mixture.
The total pressure above the mixture is the sum of the partial pressures of all compounds in the mixture.

41
Q

What is the partial pressure of acetone above a 50/50 mixture of water and acetone at STP, if the pure vapor pressure of acetone at STP is 100 torr?

A

The partial pressure of acetone above the mixture is 50 torr.

Raoult’s Law: Ptotal = P0AxA + P0BxB + …

The partial pressure of a liquid above a mixture is equal to the vapor pressure of the pure liquid times its mole fraction in the mixture. In a 50/50 mixture of water and acetone, each one will only contribute 1/2 the vapor pressure of the pure compound; hence acetone’s partial pressure will be 1/2 its pure vapor pressure.

42
Q

Define:

Non-ideal solutions

A

A non-ideal solution is one that does not follow Raoult’s Law. Non-ideal solutions can have pressures either greater or less than those predicted by Raoult’s Law.

Ex: A solution with a large attraction between solute and solvent, such as salt water, will have a lower pressure than that predicted by Raoult’s Law.

43
Q

Define:

the van’t Hoff factor

A

The van’t Hoff factor, i, is the number of particles that an ionic solute yields upon dissociation.

Ex: NaCl dissociates into Na+ and Cl- ions, a total of 2 per equivalent of NaCl. The van’t Hoff factor for NaCl is i=2.

44
Q

What are the van’t Hoff values for:

  1. KOH
  2. C6H12O6
  3. H2SO4
  4. CaCl2
  5. H3[CuNH3Cl5]
A
  • KOH: i = 2 (K+ and OH-)
  • C6H12O6: i = 1 (glucose doesn’t dissociate)
  • H2SO4: i = 3 (2H+ and SO42-)
  • CaCl2: i = 3 (2Cl- and Ca2+)
  • H3[CuNH3Cl5]: i = 4 (3H+ and [CuNH3Cl5]3-)

Note: on the AP Chem exam you are expected to know that a complex ion (in brackets) will not further dissociate.

45
Q

define:

molality (m)

A

Molality is defined as the number of moles of solute dissolved into a certain mass of solvent, with units mol/kg.

Molality uses the lowercase letter m in equations, be careful not to confuse this with mass.

46
Q

Give the equation for:

boiling point elevation of a mixture

A

ΔTb = Kbmi

where:
Kb= is the solvent’s ebullioscopic constant in kg*K/mol
m = the mixture’s molality in mol/kg
i = the solute’s van’t Hoff factor
ΔTb = boiling point temperature change in K

47
Q

Which will have a higher boiling point: a 1m NaCl solution or a 1m glucose solution?

A

The NaCl solution will have the higher boiling point.

ΔTb = Kbmi.
i(NaCl) = 2, while i(glucose) = 1, so the NaCl solution will have twice the boiling point elevation.

48
Q

When can the boiling point elevation equation not be used?

A

The boiling point elevation equation cannot be used when:

  • the mixture is volatile
  • the solute does not fully dissolve or dissociate
  • there is a significant density difference between the mixture components such that one always sits on top of the other
49
Q

How does addition of a solute affect the freezing point of a liquid?

A

The freezing point of a liquid will be lowered (colder), as the solute hinders proper crystallization of the solvent molecules.

50
Q

Give the equation for:

freezing point depression of a mixture

A

ΔTf = Kfmi

where:
Kf= is the solvent’s cryoscopic constant in kg*K/mol
m = solution molality in mol/kg
i = solute’s van’t Hoff factor
ΔTf = change in freezing point temperature in K

51
Q

Which will have a lower freezing point, 1m CaCl2 solution, or a 2m CaCl2 solution?

A

The freezing point of the 2m solution will be lower.

Based on the equation ΔTf = Kfmi the 2m solution will have twice the drop in freezing point of the 1m solution.

52
Q

A chemist wants to melt the ice on her driveway on a very cold day, and has equal amounts of each of the following substances to add to the ice. Based on colligative properties, which would work the best?

  1. H3PO4
  2. MgCl2
  3. C12H22O11
A

H3PO4 will melt the ice the most effectively.

Assuming equimolal addition, H3PO4 will yield the largest decrease in the water’s freezing point, due to its having the highest van’t Hoff factor.

i(H3PO4) = 4 (phosphoric acid)
i(MgCl2) = 3 (magnesium chloride)
i(C12H22O11) = 1 (sucrose)

53
Q

Which of the following will have the lowest freezing point?

  1. 1M BrCl (aq)
  2. .5M Na2SO3 (aq)
  3. 1.5M KCl (aq)
A

1.5M KCl (aq) has the lowest freezing point.

Freezing point depression is dependent on the number of solute atoms per volume:
BrCl = 1M x 2ions = 2
Na2SO3 = .5M x 3ions = 1.5
KCl = 1.5M x 2ions = 3

54
Q

Give a definition for:

hypertonic, isotonic, and hypotonic solutions

A

A solution’s tonicity can only be defined relative to another solution, often referred to as the “standard”.

  • A hypertonic solution has a higher concentration of solute than the standard.
  • An isotonic solution is the same concentration of solute as the standard.
  • A hypotonic solution has a lower concentration of solute than the standard.
55
Q

Define:

osmotic pressure

A

Osmotic pressure is the pressure that results from osmosis.

As water molecules move across a semi-permeable membrane to create isotonicity on both sides, the pressure of both sides will change; this pressure is the osmotic pressure.

Ex: If red blood cells are placed in a hypo/hypertonic environment, they will expand/contract in response.

56
Q

Give the equation for:

osmotic pressure

A

π = nRTi
V

where:
Π = osmotic pressure in atm
n = number of moles of solute
R = ideal gas constant in L-atm/K*mol
T = temperature in K
i = solute’s van’t Hoff factor
V = volume in L

57
Q

At a constant temperature, how will the osmotic pressure of a 1M NaCl solution compare to that of a 1M CaCl2 solution?

A

The osmotic pressure of the CaCl2 solution will be higher.

This is due to CaCl2 having a van’t Hoff factor of 3 (Ca+2, Cl-, Cl-) compared to 2 for NaCl (Na+, Cl-).

Since Π = iMRT, and M,R,T are all constant, the determining factor is i=3 > i=2.

58
Q

How will increasing a solution’s temperature change its osmotic pressure?

A

The osmotic pressure will be greater at a higher temperature.

Since Π = iMRT, and i,M,R are all constant, the determining factor is Tnew > Told.

59
Q

Define:

colloid

A

A colloid is a system where one substance is microscopically dispersed evenly throughout another substance.

Colloidal particles will not eventually settle to the bottom due to gravity and time, and can occur in any of the three phases of matter.

Ex: air is a gaseous colloid of many gases, milk is a liquid colloid of liquid butterfat in aqueous solution, colored glass is a solid colloid of metal oxides in a silica matrix.

60
Q

What are the common characteristics of colloids?

A
  • Colloids have a continuous phase (dispersion medium) and an internal phase (dispersed medium).
  • The internal phase particles are too small to easily extract physically (though it is possible to extract them chemically).
  • Liquid and solid colloids scatter light (opaque or translucent); only gas colloids can be colorless.
61
Q

Define:

the Tyndall effect

A

The Tyndall effect is demonstrated when light hits a colloid, and is scattered by the suspended particles.

Ex: common baking flour suspended in water will look slightly blue, because blue light scatters back more strongly than red light.

62
Q

Define and give the equation for:

Henry’s Law

A

Henry’s Law states that the solubility of a gas in a liquid is directly proportional to the partial pressure of that gas above the liquid.

Assuming constant temperature, the equation is:

Pv = c*kH

where:
Pv = partial pressure of the soluble gas above the solution in atm
c = concentration of the soluble gas in mol/L
kH = Henry’s Law constant of the liquid in L*atm/mol

63
Q

Explain why a can of soda will “hiss” when opened, but eventually go “flat”, based on Henry’s Law?

A

The soda is carbonated under high pressure of CO2. The CO2 is kept in solution by a high pressure of CO2 in the can, causing the “hiss” when the CO2 escapes quickly.

Once the CO2 dissipates, however, the partial pressure of CO2 above the soda drops to atmospheric levels, near zero. Henry’s Law predicts that this will lead to the amount of CO2 dissolved in the solution decreasing over time.

Chemistry 101 (13 decks)

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