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AP Chemistry > Chapter 4 - Aqueous Reactions & Solution Stoichiometry > Flashcards

Chapter 4 - Aqueous Reactions & Solution Stoichiometry Flashcards

-Recognize compounds as acids or base, & as strong, weak, or nonelectrolytes -Recognize reactions as acid-base, precipitation, metathesis, or redox -Be able to calculate moles or grams of substances in solution using molarity -Understand how to carry out a dilution to achieve a desire solution concentration -Understand how to perform & interpret the results of a titration

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1
Q

4.1

Define “electrolyte”.

A

A substance (such as NaCl) whose aqueous solutions contain ions.

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2
Q

4.1

Define “nonelectrolyte”.

A

A substance (such as C12H22O11) that doesn’t form ions in solution.

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3
Q

4.1

The ____ process helps stabilize the ions in solution & prevents cations & anions from recombining.

A

Solvation

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4
Q

4.1

Define “strong electrolyte”.

A

Solutes that exist in solution completely or nearly completely as ions.

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5
Q

4.1

Define “weak electrolyte”.

A

Solutes that exist in solution mostly in the form of molecules with only a small fraction in the form of ions.

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6
Q

4.1

What importance does chemical equilibrium have in solutions?

A

Weak electrolytes have the tendency to ionize and then recombine. The balance between these opposing processes determines the relative numbers of ions and neutral molecules, which produces a state of chemical equilibrium in which the relative numbers of each type of ion or molecule in the reaction are constant over time.

  • Varies from one electrolyte to another
  • Strong electrolytes don’t recombine
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7
Q

4.2

Define “solubility”.

A

The solubility of a substance at a given temperature is the amount of the substance that can be dissolved in a given quantity of solvent at the given temperature.

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8
Q

4.2

Cl- solubility

A

Soluble except:

Ag+, Hg2 2+, Pb 2+

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9
Q

4.2

Br- solubility

A

Soluble except:

Ag+, Hg2 2+, Pb 2+

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10
Q

4.2

I- solubility

A

Soluble except:

Ag+, Hg2 2+, Pb 2+

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11
Q

4.2

SO4 2- solubility

A

Soluble except:

Sr 2+, Ba 2+, Hg2 2+, Pb 2+

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12
Q

4.2

S 2- solubility

A

Insoluble except:

Alkali’s, NH4+
Ca 2+, Sr 2+, Ba 2+

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13
Q

4.2

CO3 2- solubility

A

Insoluble except:

Alkali’s, NH4+

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14
Q

4.2

PO4 3- solubility

A

Insoluble except:

Alkali’s, NH4+

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15
Q

4.2

OH- solubility

A

Insoluble except:

Alkali’s, NH4+
Ca 2+, Sr 2+, Ba 2+

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16
Q

4.2

Define “precipitation reaction”.

A

Reactions that result in the formation of an insoluble product.

Ex. Pb(NO3)2(aq) + 2 KI(aq) —> KNO3(aq) + PbI2(s)

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17
Q

4.2

Define “exchange (metathesis) reaction”.

A

AX + BY –> AY + BX

Ex. AgNO3(aq) + KCl(aq) —> AgCl(s) + KNO3(aq)

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18
Q

4.2

Define “molecular equation”.

A

An equation showing the complete chemical formulas of the reactants & products.

Ex. Pb(NO3)2(aq) + 2 KI(aq) —-> KNO3(aq) + PbI2(s)

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19
Q

4.2

Define “complete ionic equation”.

A

An equation with all the soluble strong electrolytes shown as ions.

Ex. Pb 2+(aq) + 2 NO3-(aq) + 2 K+(aq) + 2 I-(aq) —> PbI2(s) + 2 K+ (aq) + 2 NO3- (aq)

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20
Q

4.2

Define “spectator ion”.

A

Ions that appear in identical forms among both the reactants & products of a complete ionic equation.

Ex. 2 K+, 2 NO3-

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21
Q

4.2

Define “net ionic equation”.

A

The ionic equation left when spectator ions are omitted.

Ex. Pb 2+ (aq) + 2 I- (aq) –> PbI2 (s)

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22
Q

4.2

If every ion in a complete ionic equation is a spectator, then _____.

A

no reaction occurs.

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23
Q

4.2 Samp. Ex.

Write the net ionic equation for the precipitation reaction that occurs when solutions of calcium chloride & sodium carbonate are mixed.

A

CaCl2 (aq) + Na2CO3 (aq) –> CaCO3 (s) + 2 NaCl (aq) = molecular equation

Ca 2+ (aq) + CO3 2- (aq) —-> CaCO3(s) = ionic equation

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24
Q

4.3

Define “acids”.

A

Substances that ionize in aqueous solutions to form H ions, thereby increasing the concentration of H+ (aq) ions.

H+ donor.

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25
Q

4.3

Define “monoprotic acids”.

A

Acids that ionize to yield one H+ per molecule of acid.

Ex. HCl, HNO3

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26
Q

4.3

Define “diprotic acids”.

A

Acids that ionize to yield two H+ per molecule of acid.

Ex. H2SO4

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27
Q

4.3

Describe the ionization of H2SO4 and other diprotic acids.

A

It occurs in two steps.

H2SO4 (aq) —> H+ (aq) + HSO4- (aq)
HSO4- (aq) —> H+ (aq) + SO4 2- (aq)

Thus, aqueous solutions of H2SO4 contain a mixture of H+ (aq), HSO4- (aq), and SO4 2- (aq).

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28
Q

4.3

Define “bases”.

A

Substances that react with H+ ions.

H+ acceptor.

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29
Q

4.3

What do bases produce when they dissolve in water?

A

OH-

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30
Q

4.3

Define “strong acids/bases”.

A

Acids/bases that are strong electrolytes.

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31
Q

4.3

Define “weak acids/bases”.

A

Acids/bases that are weak electrolytes.

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32
Q

4.3

Name the 7 strong acids.

A
  1. HCl
  2. HBr
  3. HI
  4. HClO3
  5. HClO4
  6. HNO3
  7. H2SO4
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33
Q

4.3

Name the strong bases

A

Group 1A metal hydroxides
LiOH, NaOH, KOH, RbOH, CsOH

Heavy group 2A metal hydroxides
Ca(OH)2, Ba(OH)2, Sr(OH)2

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34
Q

4.3

Strong acids are more reactive than weak acids when the reactivity depends only on….

A

the concentration of H+ (aq).

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35
Q

4.3

Explain how the reactivity of an acid can depend on the anion as well as on H+ (aq).

A

For example, HF is a weak acid, but it’s very reactive & vigorously attacks many substances, including glass. This reactivity is due to the combined action of H+ (aq) & F- (aq).

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36
Q

4.3 Samp. Ex.

Classify each of the following dissolved substances as a strong electrolyte, weak, electrolyte, or nonelectrolyte: CaCl2, HNO3, C2H5OH (ethanol), HCOOH (formic acid), KOH

A

Strong: CaCl2 (ionic), HNO3 (strong acid), KOH (strong base)

Weak: HCOOH (weak acid)

Non: C2H5OH (molecular compound - not a metal hydroxide, doesn’t ionize)

37
Q

4.3

Define “neutralization reaction”.

A

When a solution of an acid and a solution of a base are mixed.

38
Q

4.3

A neutralization reaction between an acid and a metal hydroxide produces ______.

A

H2O & a salt.

Ex. HCl(aq) + NaOH(aq) —-> H2O(l) + NaCl(aq)

39
Q

4.3

Write the net ionic equation for HCl(aq) + NaOH(aq) —-> H2O(l) + NaCl(aq)

A

H+ (aq) + Cl- (aq) + Na+ (aq) + OH- (aq) —> H2O(l) + Na+ (aq) + Cl- (aq) = complete ionic

H+ (aq) + OH- (aq) –> H2O(l)

40
Q

4.3

Write the net ionic equation for Mg(OH)2 (s) + 2 HCl (aq) —> MgCl2 (aq) + 2 H2O(l)

A

Mg(OH)2 (s) + 2 H+ (aq) —–> 2 H2O(l) + Mg 2+ (aq)

41
Q
  1. 3 Samp. Ex.
    (a) Write a balanced molecular equation for the reaction between aqueous solutions of acetic acid (CH3COOH) & barium hydroxide, Ba(OH)2.
    (b) Write the net ionic equation for this reaction.
A

Ba(OH)2 (aq) + 2 CH3COOH(aq) —-> 2 H2O(l) + Ba(CH3COO)2(aq)

Ba 2+ (aq) + 2 OH- (aq) + 2 CH3COOH (aq) —-> 2 H2O (l) + Ba 2+ (aq) + 2 CH3COO- (aq) = complete ionic

2 OH- (aq) + 2 CH3COOH(aq) —> 2 H2O(l) + 2 CH3COO- (aq)

Reduce to 1. This is the net ionic.

42
Q

4.3

When does H2S(g) form?

A

When an acid reacts with a metal sulfide.

43
Q

4.3

What does CO3 always decompose into?

A

CO2(g) and H2O(l)

44
Q

4.1

When asked what causes electrolyte solutions to conduct electricity, a student responds that it is due to the movement of electrons through the solution. Is the student correct? If not, what is the correct response?

A

No. Electrolyte solutions conduct electricity because the dissolved ions carry charge through the solution from one electrode to the other.

45
Q

4.3

Which of the following solutions has the largest concentration of solvated protons: (a) 0.2 M LiOH, (b) 0.2 M HI, (c) 1.0 M methyl alcohol (CH3OH)? Explain.

A

HI.

LiOH is a strong base, HI is a strong acid, and CH3OH is a molecular compound & nonelectrolyte. The H+ in acids are protons.

46
Q

4.4

Define “redox reaction”.

A

Electrons are transferred between reactants.

47
Q

4.4

Define “oxidation”.

A

Loss of electrons by a substance.

Ex. 2 Ca(s) + O2(g) —> 2 CaO(s)

Oxygen goes from neutral to a -2 oxidation state.

48
Q

4.4

Define “reduction”.

A

Gain of electrons by a substance.

49
Q

4.4

For an atom in its elemental form, the oxidation number is always…

A

zero.

Ex. Each H atom in H2 has an oxidation # of 0
Ex. Each P atom in the P4 molecule has an oxidation # of zero.

50
Q

4.4

Define “oxidation number”.

A

The actual charge for a monatomic ion.

51
Q

4.4

For any monatomic ion, the oxidation number equals …

A

the charge on the ion.

Ex. K+ has an oxidation # of +1.
Ex. S 2- has an oxidation # of -2.

52
Q

4.4

Nonmetals usually have ___ oxidation numbers.

A

Negative

53
Q

4.4

The oxidation # of oxygen is usually ___ in ionic & molecular compounds. The major except is in peroxides, which contain the O2 2- ion, giving each oxygen an oxidation number of ____.

A

-2 , -1

54
Q

4.4

The oxidation number of hydrogen is usually ___ when bonded to nonmetals and _____ when bonded to metals.

A

+1, -1

55
Q

4.4

The oxidation number of fluorine is ____ in all compounds. The other halogens have an oxidation number of ___ in most binary compounds; however, when combined with oxygen as in oxyanions, they have positive oxidation states.

A

-1, -1

56
Q

4.4

The sum of the oxidation numbers of all atoms in a neutral compound is ___. The sum of the oxidation numbers in a polyatomic ion equals the _____.

A
  1. charge of the ion.

Ex. In H3O+,

Each H has an oxidation number of +1, O has an oxidation number of -2.

(1) + (1) + (1) + (-2) = +1 … the charge of H3O.

57
Q

4.4

Alkali metal ions always have an oxidation number of _____.

A

+1

58
Q

4.4

Alkaline earth metal ions always have an oxidation number of _____.

A

+2

59
Q

4.4

Group 3A ions always have an oxidation number of _____.

A

+3

60
Q

Determine the oxidation # of sulfur in each of the following:

(a) H2S
(b) S8
(c) SCl2
(d) Na2SO3
(e) SO4 2-

A

a. -2. Each H is +1, so 1 + 1 + x = 0
b. 0. It’s in its elemental state.
c. +2. Each Cl is -1, so -1 + (-1) + x = 0
d. +4. Each Na is +1 & each O is -2 , so 1 + 1 + (-2) + (-2) + (-2) + x = 0
3. +6. Each O is -2, so -2 + (-2) + (-2) + (-2) + x = -2

61
Q

4.4

Define “displacement reaction”.

A

A + BX —> AX + B

Ex. Mg (s) + 2 HCl (aq) —> MgCl2 (aq) + H2 (g)

Mg is oxidized (loses 2 electrons), H is reduced (gains 1)
Mg goes from ox. state of 0 to +2, hydrogen goes from state of +1 to 0.

62
Q

4.4

Whenever one substance is oxidized, some other substance must be ____.

A

Reduced.

63
Q

4.4

Define “activity series”.

A

A list of metals arranged in order of decreasing ease of oxidation.

64
Q

4.4

Any metal on an activity series can be ____ by the ions below it.

A

Oxidized.

Ex. Copper is above silver.

Cu(s) + 2 Ag+ (aq) —> Cu 2+ (aq) + 2 Ag (S)

Copper loses electrons, while silver gains electrons.

65
Q

4.4 Give It Some Thought (p. 142)

Which is more easily reduced, Mg 2+ (aq) or Ni 2+ (aq)?

A

Ni 2+. It’s lower on the activity series.

66
Q

4.4

Only metals ____ hydrogen in the activity series are able to react with acids to form H2.

A

Above.

Ex. Ni (s) + 2 HCl (aq) —> NiCl2 (aq) + H2 (g)

67
Q

4.4

Ex…

A

Cu(s) + 4 HNO3 (aq) —> Cu(NO3)2 (aq) + 2 H2O (l) + 2 NO2 (g)

Cu is oxidized by nitrate and is accompanied by the formation of brown NO2 (g).

68
Q

4.5

Scientists use the term concentration to _____

A

designate the amount of solute dissolved in a given quantity of solvent or quantity of solution.

69
Q

4.5

What does Molarity (M) express?

A

The concentration of a solution as the number of moles of solute in a L of solution.

70
Q

4.5

What’s the equation for Molarity?

A

Mol solute / L soln

71
Q

4.5

A 1.0 M solution of NaCl is ___ M in Na+ ions and ___ M in Cl- ions.

A

1.0, 1.0

72
Q

4.5

A 1.0 M solution of Na2SO4 is ___ M in Na+ ions and ___ M in SO4 2- ions.

A

2.0, 1.0

73
Q

4.5 Samp. Ex

What are the molar concentrations of each of the ions present in 0.025 M aqueous solution of calcium nitrate?

A

Ca(NO3)2

.025 M Ca 2+, .05 M NO3-.

74
Q

4.5 Prac. Ex.

What is the molar concentration of K+ ions in a 0.015 M solution of K2CO3?

A

.03 M K+

75
Q

4.5 Samp. Ex.

How many grams of Na2SO4 are required to make 0.350 L of 0.500 M Na2SO4?

A

.350 L (.500 mol Na2SO4 / 1 L) = .175 mol Na2SO4

mol to gram = 24.9g Na2SO4

76
Q

4.5

Moles solute before dilution = …
Moles solute in concentrated soln = …

A

mol solute after dilution

mol solute in diluted soln

77
Q

4.5

M concentrated x Volume concentrated =

A

M dilute x Volume dilute

78
Q

4.5 Give It Some Thought (p. 148)

How is the molarity of a 0.50 M KBr solution changed when water is added to double its volume?

A

(.5 mol)(1 L) = x(2 L)

x = .25 mol

79
Q

4.5

How many mL of 3.0 M H2SO4 are needed to make 450 mL of 0.10 M H2SO4?

A

(3.0 M)(x mL) = (.10 M)(450 mL)

x = 15 mL

80
Q

4.6 Samp. Ex.

How many grams of Ca(OH)2 are needed to neutralize 25.0 mL of 0.100 M HNO3?

A

.025 L (.100 mol / 1 L) = .0025 mol HNO3
Divide by 2. .00125 mol Ca(OH)2.

.00125 mol Ca(OH)2 (74.1g / 1 mol) = .093g

81
Q
  1. 6 Give It Some Thought (p. 150)
  2. 00 mL of a .100 M HBr solution is titrated with a 0.200 M NaOH solution. How many mL of the NaOH solution are required to reach the equivalence point?
A

12.50 mL

82
Q

4.6 Samp. Ex

Ag+ (aq) + Cl- (aq) —> AgCl (s)

(a) How many grams of Cl- are in a sample of H2O if 20.2 mL of 0.100 M Ag+ is needed to react with all the Cl- in the sample?
(b) If the sample has a mass of 10.0g, what % Cl- does it contain?

A

(a) .0202 L Cl- (.100 mol / 1 L) = .00202 mol ( 35.45g / 1 mol) = .072g Cl-
(b) (.072g / 10.0g) x 100 = .72%

83
Q

4.6 Prac. Ex.

What is the molarity of an NaOH solution if 48.0 mL is needed to neutralize 35.0 mL of 0.144 M H2SO4?

A

.035 L (.144 mol / 1 L) = .00504 mol H2SO4

.00504 mol H2SO4 (2 mol NaOH / 1 mol H2SO4) = .01008 mol NaOH

.01008 mol NaOH / .048 = .210 M

84
Q

4.1

Define “solvation”.

A

The clustering of solvent molecules around a solute particle.

85
Q

4.1 - Give It Some Thought (p. 122)

What dissolved species are present in a solution of (a) KCN, (b) NaClO4?

A

(a) K+(aq), CN-(aq)

b) Na+(aq), ClO4-(aq

86
Q

4.1

Why can’t a substance be classified as strong or weak based on its solubility?

A

Strong or weak electrolytes are categorized by whether or not they dissociate completely, or whether they dissociate into molecules + ions/just molecules, not to the extent that they dissolve.

Ex. Ba(OH)2 is pretty insoluble, but it’s a strong electrolyte because the portion that does dissolve dissociates completely into Ba 2+ and OH-.

87
Q

4.1 - Give It Some Thought (p. 123)

Which solute will cause the lightbulb in the experiment to glow more brightly, CH3OH or MgBr2?

A

MgBr2. It’s a strong electrolyte, while CH3OH is a weak electrolyte.

88
Q

4.1 - Practice Ex. 4.1

How many anions would you show if the diagram contained 6 cations?

(a) NiSO4
(b) Ca(NO3)2
(c) Na3PO4
(d) Al2(SO4)3

A

(a) 6, (b) 12, (c) 2, (d) 9

AP Chemistry (20 decks)

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