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MCAT GC WORKING > Solution Chem > Flashcards

Solution Chem Flashcards

1
Q
A
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2
Q

Define and give an example of:

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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3
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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4
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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5
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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6
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 nonmetal so that both can form stable octets.

Exception: The noble gases, which already have a full octet, typically go into solution as neutral atomic species.

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7
Q

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

  1. Li
  2. Na
  3. 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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8
Q

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

  1. Br
  2. Cl
  3. 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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9
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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10
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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11
Q

Give the molecular form and charge for these common anions:

  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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12
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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13
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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14
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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15
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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16
Q

Explain why the MCAT 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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17
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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18
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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19
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).

20
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.

21
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.

22
Q

Define:

precipitation

A

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

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

23
Q

Identify 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.

24
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

25
# Identify and give the equation for: Molality (m)
Molality is a measure of the concentration of a solution, given in units of moles of solute dissolved per kilogram of solvent.
26
58.5 grams of NaCl are dissolved in 2.0 kg of water at room temperature. What is the molality of the solution?
0.5 m solution of NaCl ## Footnote molality = mol/kg 58.5 grams is the molecular weight of NaCl, so there is 1 mole present. molality = 1mol / 2kg = .5m
27
# Identify and give the equation for: Mole fraction (x)
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.
28
# Identify and give the equation for: Percent composition by mass (%)
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.
29
# Identify and give the equation for: Parts per million (ppm)
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.
30
# Identify and give the equation for: Normality (N)
Normality is the concentration of the reactive species in mole-equivalents compared to the volume of solvent in liters. Normality (N) = Molarity \* (equivalents of reactive species per mole)
31
What is the acid Normality (N) of a 1M solution of sulfuric acid?
2N solution. Normality (N) = Molarity \* (equivalents of reactive species per mole) In this case, equivalents of acid are being asked for. H2SO4 has two H+ ions dissociate in solution for every one H2SO4. Hence: 1M \* (2 equivalents) = 2N
32
# What is the equation to calculate: the change in molarity of a diluted solution
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
33
# Define: the solubility product constant
The solubility product constant (Ksp) is the equilibrium constant for a solvation reaction. Ksp = [anions]a[cations]c ## Footnote Just like other equilibrium constants, Ksp is equal to the concentration of products over reactants raised to the power of their coefficients. However, since pure liquids and solids are omitted from equilibrium calculations, the equation simplifies to just the concentration of ions in solution, raised to the power of their coefficients.
34
If the temperature of a solution is increased, what happens to the solute's Ksp?
Ksp increases. ## Footnote A warmer solvent can dissolve more solutes. This leads to a higher concentration of ions in solution, thus Ksp increases.
35
What is the solubility product constant for AgCl, which dissociates as below: AgCl(s) ⇔ Ag+(aq) + Cl-(aq)
Ksp = [Ag+][Cl-] ## Footnote The coefficient in front of both Silver and Chloride is 1, so the exponents will be 1.
36
What is the solubility product constant for BaF2, which dissociates as below: BaF2(s) ⇔ Ba++ (aq) + 2F-(aq)
Ksp= [Ba++][F-]2 ## Footnote The coefficient in front of Barium is 1 so the exponent will be 1 over [Ba++]. The coefficient in front of Floride is 2, so the exponent will be 2 over [F-].
37
# Define: **the common ion effect**
If a solution already contains ions that a salt would separate into when dissolved, they are referred to as **common ions**. The salt will be less soluble than if dissolved in a pure solvent. ## Footnote Ex: NaCl dissolved into chlorinated water (Cl- ions present) will be less soluble than NaCl in pure water. The added Cl- ion (product) shifts equilibrium to the side containing the solid salt (reactant).
38
Suppose NaCl were added to a saturated aqueous solution of AgCl. What will happen? Note: NaCl has a higher solubility than AgCl.
1. NaCl will still dissolve. The presence of Cl- ions from AgCl will not prevent the highly soluble NaCl from dissociating. 2. AgCl will precipitate. The additional Cl- ions now present from NaCl exceed the levels for AgCl in solution, and will force those ions back into solid AgCl salt form.
39
A given solution contains PbCl2 and PbSO4. How could lead (II) chloride be selectively harvested from the solution?
If a highly soluble chloride salt such as NaCl were added to the solution, the common ion effect would lead to the selective precipitation of lead (II) chloride.
40
# Define: **Complex (or coordination complex)**
A **complex** is formed when a metallic atom or ion is bonded to a surrounding array of molecules or anions; usually written with brackets as: [complex]. Complexes cannot be solvated back into their original atoms/ions. ## Footnote Ex: Cisplatin [PtCl2(NH3)2] is formed when the Pt++ cation is bonded to two Cl- anions and two NH3 groups.
41
What is the affect on solubility of a salt if: 1. there are common ions present 2. complexes form from the product ions
1. Solubility decreases. Common ions prevent salts from fully dissociating, since the solution already contains some concentration of product ions. 2. Solubility increases. Complexes formed from product ions effectively remove those ions from solution, allowing more salt to dissociate and replace the lost product ions.
42
Suppose that a solution is saturated in both K2[PtCl4] and NH3. What must be true of the solubility of both species if the following reaction occurs? PtCl4-- + NH3 ⇒ PtCl2(NH3)2 + 2 Cl-
Solubility for both will be increased. ## Footnote The solution will accept more K2[PtCl4] and NH3, since the [PtCl2(NH3)2] complex that forms effectively removes both PtCl4-- ions and NH3 molecules from solution. According to LeChatelier's principle, the system then shifts towards the product ions, increasing the solubility of both species.
43
Explain why: 1. acids are more soluble in an aqueous base than neutral water. 2. bases are more soluble in an aqueous acid than neutral water.
1. Base molecules already present in a basic solution will bond to the H+ ions, removing them from solution, shifting equilibrium to the right, and increasing solubility of HA. HA + H2O ⇔ H+ + A-  2. Acid molecules already present in an acidic solution will bond to the OH- ions, removing them from solution, shifting equilibrium to the right, and increasing solubility of B-. B- + H2O ⇔ BH + HO-
44
Give the molecular form and charge for these common anions: 1. hydroxide 2. chloride 3. nitrate 4. bromide 5. permanganate
1. hydroxide, OH- 2. chloride, Cl- 3. nitrate, NO3- 4. bromide, Br- 5. permanganate, MnO4- All of these anions have a -1 charge in solution.
45
Give the molecular form and charge for these common anions: 1. ammonium 2. sodium 3. silver 4. hydronium 5. copper(I) cuprous
1. ammonium, NH4+ 2. sodium, Na+ 3. silver, Ag+ 4. hydronium, H3O+ 5. copper(I) cuprous, Cu+ All of these cations have a +1 charge in solution.
46
Give the molecular form and charge for these common anions: 1. sulfate 2. dichromate 3. peroxide 4. hydrogen phosphate 5. sulfite
1. sulfate, SO4-- 2. dichromate, Cr2O7-- 3. peroxide, O2-- 4. hydrogen phosphate, HPO4-- 5. sulfite, SO3-- All of these anions will form a -2 charge in solution.
47
Give the molecular form and charge for these common cations: 1. Iron (II) ferrous 2. Magnesium 3. Calcium 4. mercury(II) mercuric 5. Barium
1. Iron (II) ferrous, Fe++ 2. Magnesium, Mg++ 3. Calcium, Ca++ 4. mercury(II) mercuric, Hg++ 5. Barium, Ba++ All of these cations will form a +2 charge in solution.

MCAT GC WORKING (11 decks)

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