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
4.1
Define “electrolyte”.
A substance (such as NaCl) whose aqueous solutions contain ions.
4.1
Define “nonelectrolyte”.
A substance (such as C12H22O11) that doesn’t form ions in solution.
4.1
The ____ process helps stabilize the ions in solution & prevents cations & anions from recombining.
Solvation
4.1
Define “strong electrolyte”.
Solutes that exist in solution completely or nearly completely as ions.
4.1
Define “weak electrolyte”.
Solutes that exist in solution mostly in the form of molecules with only a small fraction in the form of ions.
4.1
What importance does chemical equilibrium have in solutions?
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
4.2
Define “solubility”.
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.
4.2
Cl- solubility
Soluble except:
Ag+, Hg2 2+, Pb 2+
4.2
Br- solubility
Soluble except:
Ag+, Hg2 2+, Pb 2+
4.2
I- solubility
Soluble except:
Ag+, Hg2 2+, Pb 2+
4.2
SO4 2- solubility
Soluble except:
Sr 2+, Ba 2+, Hg2 2+, Pb 2+
4.2
S 2- solubility
Insoluble except:
Alkali’s, NH4+
Ca 2+, Sr 2+, Ba 2+
4.2
CO3 2- solubility
Insoluble except:
Alkali’s, NH4+
4.2
PO4 3- solubility
Insoluble except:
Alkali’s, NH4+
4.2
OH- solubility
Insoluble except:
Alkali’s, NH4+
Ca 2+, Sr 2+, Ba 2+
4.2
Define “precipitation reaction”.
Reactions that result in the formation of an insoluble product.
Ex. Pb(NO3)2(aq) + 2 KI(aq) —> KNO3(aq) + PbI2(s)
4.2
Define “exchange (metathesis) reaction”.
AX + BY –> AY + BX
Ex. AgNO3(aq) + KCl(aq) —> AgCl(s) + KNO3(aq)
4.2
Define “molecular equation”.
An equation showing the complete chemical formulas of the reactants & products.
Ex. Pb(NO3)2(aq) + 2 KI(aq) —-> KNO3(aq) + PbI2(s)
4.2
Define “complete ionic equation”.
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)
4.2
Define “spectator ion”.
Ions that appear in identical forms among both the reactants & products of a complete ionic equation.
Ex. 2 K+, 2 NO3-
4.2
Define “net ionic equation”.
The ionic equation left when spectator ions are omitted.
Ex. Pb 2+ (aq) + 2 I- (aq) –> PbI2 (s)
4.2
If every ion in a complete ionic equation is a spectator, then _____.
no reaction occurs.
4.2 Samp. Ex.
Write the net ionic equation for the precipitation reaction that occurs when solutions of calcium chloride & sodium carbonate are mixed.
CaCl2 (aq) + Na2CO3 (aq) –> CaCO3 (s) + 2 NaCl (aq) = molecular equation
Ca 2+ (aq) + CO3 2- (aq) —-> CaCO3(s) = ionic equation
4.3
Define “acids”.
Substances that ionize in aqueous solutions to form H ions, thereby increasing the concentration of H+ (aq) ions.
H+ donor.
4.3
Define “monoprotic acids”.
Acids that ionize to yield one H+ per molecule of acid.
Ex. HCl, HNO3
4.3
Define “diprotic acids”.
Acids that ionize to yield two H+ per molecule of acid.
Ex. H2SO4
4.3
Describe the ionization of H2SO4 and other diprotic acids.
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).
4.3
Define “bases”.
Substances that react with H+ ions.
H+ acceptor.
4.3
What do bases produce when they dissolve in water?
OH-
4.3
Define “strong acids/bases”.
Acids/bases that are strong electrolytes.
4.3
Define “weak acids/bases”.
Acids/bases that are weak electrolytes.
4.3
Name the 7 strong acids.
- HCl
- HBr
- HI
- HClO3
- HClO4
- HNO3
- H2SO4
4.3
Name the strong bases
Group 1A metal hydroxides
LiOH, NaOH, KOH, RbOH, CsOH
Heavy group 2A metal hydroxides
Ca(OH)2, Ba(OH)2, Sr(OH)2
4.3
Strong acids are more reactive than weak acids when the reactivity depends only on….
the concentration of H+ (aq).
4.3
Explain how the reactivity of an acid can depend on the anion as well as on H+ (aq).
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).
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
Strong: CaCl2 (ionic), HNO3 (strong acid), KOH (strong base)
Weak: HCOOH (weak acid)
Non: C2H5OH (molecular compound - not a metal hydroxide, doesn’t ionize)
4.3
Define “neutralization reaction”.
When a solution of an acid and a solution of a base are mixed.
4.3
A neutralization reaction between an acid and a metal hydroxide produces ______.
H2O & a salt.
Ex. HCl(aq) + NaOH(aq) —-> H2O(l) + NaCl(aq)
4.3
Write the net ionic equation for HCl(aq) + NaOH(aq) —-> H2O(l) + NaCl(aq)
H+ (aq) + Cl- (aq) + Na+ (aq) + OH- (aq) —> H2O(l) + Na+ (aq) + Cl- (aq) = complete ionic
H+ (aq) + OH- (aq) –> H2O(l)
4.3
Write the net ionic equation for Mg(OH)2 (s) + 2 HCl (aq) —> MgCl2 (aq) + 2 H2O(l)
Mg(OH)2 (s) + 2 H+ (aq) —–> 2 H2O(l) + Mg 2+ (aq)
- 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.
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.
4.3
When does H2S(g) form?
When an acid reacts with a metal sulfide.
4.3
What does CO3 always decompose into?
CO2(g) and H2O(l)
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?
No. Electrolyte solutions conduct electricity because the dissolved ions carry charge through the solution from one electrode to the other.
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.
HI.
LiOH is a strong base, HI is a strong acid, and CH3OH is a molecular compound & nonelectrolyte. The H+ in acids are protons.
4.4
Define “redox reaction”.
Electrons are transferred between reactants.
4.4
Define “oxidation”.
Loss of electrons by a substance.
Ex. 2 Ca(s) + O2(g) —> 2 CaO(s)
Oxygen goes from neutral to a -2 oxidation state.
4.4
Define “reduction”.
Gain of electrons by a substance.
4.4
For an atom in its elemental form, the oxidation number is always…
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.
4.4
Define “oxidation number”.
The actual charge for a monatomic ion.
4.4
For any monatomic ion, the oxidation number equals …
the charge on the ion.
Ex. K+ has an oxidation # of +1.
Ex. S 2- has an oxidation # of -2.
4.4
Nonmetals usually have ___ oxidation numbers.
Negative
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 ____.
-2 , -1
4.4
The oxidation number of hydrogen is usually ___ when bonded to nonmetals and _____ when bonded to metals.
+1, -1
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.
-1, -1
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 _____.
- 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.
4.4
Alkali metal ions always have an oxidation number of _____.
+1
4.4
Alkaline earth metal ions always have an oxidation number of _____.
+2
4.4
Group 3A ions always have an oxidation number of _____.
+3
Determine the oxidation # of sulfur in each of the following:
(a) H2S
(b) S8
(c) SCl2
(d) Na2SO3
(e) SO4 2-
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
4.4
Define “displacement reaction”.
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.
4.4
Whenever one substance is oxidized, some other substance must be ____.
Reduced.
4.4
Define “activity series”.
A list of metals arranged in order of decreasing ease of oxidation.
4.4
Any metal on an activity series can be ____ by the ions below it.
Oxidized.
Ex. Copper is above silver.
Cu(s) + 2 Ag+ (aq) —> Cu 2+ (aq) + 2 Ag (S)
Copper loses electrons, while silver gains electrons.
4.4 Give It Some Thought (p. 142)
Which is more easily reduced, Mg 2+ (aq) or Ni 2+ (aq)?
Ni 2+. It’s lower on the activity series.
4.4
Only metals ____ hydrogen in the activity series are able to react with acids to form H2.
Above.
Ex. Ni (s) + 2 HCl (aq) —> NiCl2 (aq) + H2 (g)
4.4
Ex…
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).
4.5
Scientists use the term concentration to _____
designate the amount of solute dissolved in a given quantity of solvent or quantity of solution.
4.5
What does Molarity (M) express?
The concentration of a solution as the number of moles of solute in a L of solution.
4.5
What’s the equation for Molarity?
Mol solute / L soln
4.5
A 1.0 M solution of NaCl is ___ M in Na+ ions and ___ M in Cl- ions.
1.0, 1.0
4.5
A 1.0 M solution of Na2SO4 is ___ M in Na+ ions and ___ M in SO4 2- ions.
2.0, 1.0
4.5 Samp. Ex
What are the molar concentrations of each of the ions present in 0.025 M aqueous solution of calcium nitrate?
Ca(NO3)2
.025 M Ca 2+, .05 M NO3-.
4.5 Prac. Ex.
What is the molar concentration of K+ ions in a 0.015 M solution of K2CO3?
.03 M K+
4.5 Samp. Ex.
How many grams of Na2SO4 are required to make 0.350 L of 0.500 M Na2SO4?
.350 L (.500 mol Na2SO4 / 1 L) = .175 mol Na2SO4
mol to gram = 24.9g Na2SO4
4.5
Moles solute before dilution = …
Moles solute in concentrated soln = …
mol solute after dilution
mol solute in diluted soln
4.5
M concentrated x Volume concentrated =
M dilute x Volume dilute
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?
(.5 mol)(1 L) = x(2 L)
x = .25 mol
4.5
How many mL of 3.0 M H2SO4 are needed to make 450 mL of 0.10 M H2SO4?
(3.0 M)(x mL) = (.10 M)(450 mL)
x = 15 mL
4.6 Samp. Ex.
How many grams of Ca(OH)2 are needed to neutralize 25.0 mL of 0.100 M HNO3?
.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
- 6 Give It Some Thought (p. 150)
- 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?
12.50 mL
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) .0202 L Cl- (.100 mol / 1 L) = .00202 mol ( 35.45g / 1 mol) = .072g Cl-
(b) (.072g / 10.0g) x 100 = .72%
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?
.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
4.1
Define “solvation”.
The clustering of solvent molecules around a solute particle.
4.1 - Give It Some Thought (p. 122)
What dissolved species are present in a solution of (a) KCN, (b) NaClO4?
(a) K+(aq), CN-(aq)
b) Na+(aq), ClO4-(aq
4.1
Why can’t a substance be classified as strong or weak based on its solubility?
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-.
4.1 - Give It Some Thought (p. 123)
Which solute will cause the lightbulb in the experiment to glow more brightly, CH3OH or MgBr2?
MgBr2. It’s a strong electrolyte, while CH3OH is a weak electrolyte.
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) 6, (b) 12, (c) 2, (d) 9
