Solutions

Question 1

Define:

a solution

Answer 1

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

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

Question 2

Define:

a solvent

Answer 2

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

For example, dissolving a small amount of solid NaCl into liquid water produces a liquid solution. In this case, water is the solvent.

Question 3

Define:

a solute

Answer 3

It is the substance whose phase is lost after solvation (or the substance of which less is present).

For example, dissolving a small amount of solid NaCl into liquid water produces a liquid solution. In this case, NaCl is the solute.

Question 4

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

Answer 4

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.

Question 5

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

Answer 5

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.

Question 6

What charge do Li, K, and Na usually take when forming ions in solution?

Answer 6

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.

Question 7

What charge do Br, Cl, and F usually take when forming ions in solution?

Answer 7

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

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

Question 8

Give the molecular formula and charge for these common ions:

1. ammonium
2. chloride
3. dichromate
4. mercury (II)
5. silver

Answer 8

1. ammonium, NH₄⁺, +1 charge
2. chloride, Cl^-, -1 charge
3. dichromate, Cr₂O₇^2-, -2 charge
4. mercury(II), Hg²⁺, +2 charge
5. silver, Ag⁺, +1 charge

Question 9

Give the molecular formula and charge for these common ions:

1. hydroxide
2. barium
3. sodium
4. permanganate
5. sulfite

Answer 9

1. hydroxide, OH^-, -1 charge
2. barium, Ba²⁺, +2 charge
3. sodium, Na⁺, +1 charge
4. permanganate, MnO₄^-, -1 charge
5. sulfite, SO₃^2-, -2 charge

Question 10

Give the molecular formula and charge for these common ions:

1. hydrogen phosphate
2. magnesium
3. calcium
4. bromide
5. copper (I)

Answer 10

1. hydrogen phosphate, HPO₄^2-, -2 charge
2. magnesium, Mg²⁺, +2 charge
3. calcium, Ca²⁺, +2 charge
4. bromide, Br^-, -1 charge
5. copper (I), Cu⁺, +1 charge

Question 11

Give the molecular formula and charge for these common ions:

1. sulfate
2. nitrate
3. peroxide
4. hydronium
5. iron (II)

Answer 11

1. sulfate, SO₄^2-, -2 charge
2. nitrate, NO₃^-, -1 charge
3. peroxide, O₂^2-, -2 charge
4. hydronium, H₃O⁺, +1 charge
5. iron (II), Fe²⁺, +2 charge

Question 12

Define:

solvation

Answer 12

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

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

Question 13

Define:

hydration

This term has multiple uses, but assume it refers to solutions here.

Answer 13

It is the process of solvation, specifically when water is used as the solvent.

For example, 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.

Question 14

Explain the “like dissolves like” rule.

Answer 14

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

For example, a nonpolar solute such as naphthalene is insoluble in water, slightly soluble in methanol, and highly soluble in nonpolar benzene.

Question 15

Explain why the MCAT will rarely refer to an H⁺ ion in aqueous solution. What ion will be used instead?

Answer 15

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

More commonly, the hydronium 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 one proton) and H₇O₃⁺ (three water molecules).

Question 16

Define:

solubility

Answer 16

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

Question 17

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

1. Nitrates (NO₃^-)
2. Sulfites (SO₃^–)
3. Acetates (CH₃COO^-)
4. Chlorides (Cl^-)
5. Bromides (Br^-)

Answer 17

1. Nitrates are always soluble
2. Sulfites are insoluble (except in Group I and ammonium compounds)
3. Acetates are soluble (except in silver compounds)
4. Chlorides are soluble (except in silver compounds)
5. Bromides are soluble (except in silver, lead, copper, or mercury compounds)

Question 18

Describe what is meant by a saturated solution.

Answer 18

It 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 a precipitate will form).

Question 19

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

Answer 19

unsaturated

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

Question 20

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

Answer 20

supersaturated

A solution containing more solute than needed for saturation is said to be supersaturated; it has exceeded the saturation equilibrium. This is often accomplished by heating a solution, dissolving more solvent, and then cooling it quickly.

Question 21

Define:

precipitation

Answer 21

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

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

Question 22

Define:

molarity

(M)

Answer 22

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

Shown above is the equation for calculating a solution’s molarity.

Question 23

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

Answer 23

10 mol NaCl

Molarity = mol/L

5 mol/L = x mol / 2L
x = 10 moles

Question 24

Define:

molality

(m)

Answer 24

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

Shown above is the equation for calculating a solution’s molality.

Question 25

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

Answer 25

0.5 m ## Footnote molality = mol/kg The molecular weight of NaCl is 58.5 g, so 1 mole is present. molality = 1mol / 2kg = 0.5m

Question 26

# Define: mole fraction | (x)

Answer 26

It is a measure of the concentration of a solution, given as a ratio of moles of one component to total moles of all components. ## Footnote Shown above is the equation for calculating the mole fraction of a solute.

Question 27

# Define: percent composition by mass (%)

Answer 27

It is a measure of the amount of one component of a compound in grams compared to the total number of grams of all components. ## Footnote Shown above is the equation for calculating the percent composition of a particular solute in a solution.

Question 28

# Define: parts per million | (ppm)

Answer 28

It 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. ## Footnote Shown above is the equation for calculating parts per million of a particular solute in a solution.

Question 29

What equation can be used to calculate the change in molarity of a diluted solution?

Answer 29

M₁V₁ = M₂V₂ ## Footnote Where: * M₁ = initial molarity (mol/L) * V₁ = initial volume (L) * M₂ = final molarity (mol/L) * V₂ = final volume (L)

Question 30

# Define: normality | (N)

Answer 30

It is the concentration of the reactive species in mole-equivalents compared to the volume of solvent in liters. ## Footnote Normality (N) = Molarity \* (equivalents of reactive species per mole) Normality is generally used to quantify acidic or basic equivalents.

Question 31

What is the acid normality (N) of a 1 M solution of sulfuric acid?

Answer 31

2N ## Footnote Normality (N) = Molarity \* (equivalents of reactive species per mole) In this case, equivalents of acid are being asked for. For H₂SO₄, two H⁺ ions dissociate in solution for every one H₂SO₄. 1M \* (2 equivalents) = 2N.

Question 32

# Define: solubility product constant | (K_sp)

Answer 32

It is the equilibrium constant for the solvation reaction of a salt. ## Footnote K_sp = [anions]^a[cations]^c Just like other equilibrium constants, K_sp 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.

Question 33

If the temperature of a solution is increased, what happens to the solute's K_sp?

Answer 33

increases ## Footnote A warmer solvent can dissolve more solutes. This leads to a higher concentration of ions in solution, so K_sp increases.

Question 34

What is the solubility product constant formula for AgCl, which dissociates as written below? AgCl(s) ⇔ Ag⁺(aq) + Cl^-(aq)

Answer 34

K_sp = [Ag⁺][Cl^-] ## Footnote The coefficient in front of both silver and chloride is 1, so the exponents will also be 1.

Question 35

What is the solubility product constant formula for BaF₂, which dissociates as written below? BaF₂(s) ⇔ Ba⁺⁺ (aq) + 2F^-(aq)

Answer 35

K_sp= [Ba⁺⁺][F^-]² ## Footnote The coefficient in front of barium is 1, so the exponent will be 1 for [Ba²⁺]. The coefficient in front of fluoride is 2, so the exponent will be 2 for [F^-].

Question 36

# Define: common ion effect

Answer 36

It states that a salt will be less soluble in a solution containing product ions than it is in pure solvent. A solution containing ions that a salt would separate into when dissolved is referred to as having **common ions**. ## Footnote For example, AgCl will be less soluble in chlorinated water (with Cl^- ions present) than in pure water. The added Cl^- ion (product) shifts equilibrium to the side containing the solid salt (reactant).

Question 37

What will happen if NaCl is added to a saturated aqueous solution of AgCl? ## Footnote Note that NaCl has a higher solubility than AgCl.

Answer 37

* NaCl will still dissolve. The presence of Cl^- ions from AgCl will not prevent the highly soluble NaCl from dissociating. * AgCl will precipitate. The additional Cl^- ions now present exceed the saturation levels for AgCl in solution, and will force those ions back into solid AgCl salt form.

Question 38

A given solution contains PbCl₂ and PbSO_4. How could lead (II) chloride be selectively harvested from the solution?

Answer 38

If a highly soluble chloride salt such as NaCl were added to the solution, the common ion effect of additional Cl^- ions would lead to the selective precipitation of lead (II) chloride.

Question 39

# Define: complex (or coordination complex)

Answer 39

It is formed when a metallic atom or ion is bonded to a surrounding array of molecules or anions; these are usually written with brackets as [complex]. ## Footnote Complexes cannot be solvated back into their original atoms or ions. For example, cisplatin [PtCl₂(NH₃)₂] is formed when the Pt²⁺ cation is bonded to two Cl^- anions and two NH₃ groups.

Question 40

How is the solubility of a salt affected if: 1. common ions are present? 2. complexes form from the product ions?

Answer 40

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.

Question 41

Suppose that a solution is saturated with both K₂[PtCl₄] and NH₃. What must be true of the solubility of both species if the following reaction occurs? PtCl₄^2- + NH₃ ⇒ PtCl₂(NH₃)₂ + 2 Cl^-

Answer 41

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 PtCl₄^2- ions and NH₃ molecules from solution. According to Le Chatelier's principle, the system then shifts towards the product ions, increasing the solubility of both species.

Question 42

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.

Answer 42

Base molecules already present in the basic solution will bond to the H⁺ ions, removing them from solution, shifting equilibrium to the right, and increasing the solubility of the acid (HA). HA + H₂O ⇔ H⁺ + A^- Acid molecules already present in the acidic solution will bond to the OH^- ions, removing them from solution, shifting equilibrium to the right, and increasing the solubility of the base (B^-). B^- + H₂O ⇔ HB + OH^-