1 Fossum-Reyes

Chapter 4 Aqueous Reactions and Solution Stoichiometry

4.1 General Properties of Aqueous Solutions What is a solution? How do you identify the following two?

• Solvent.

• Solute(s). Dissociation. What is it? KCl(aq) = K+ (aq) + Cl- (aq) CuSO4(aq) = Cu+2(aq) + SO4

2-(aq) K2SO4(aq) = 2 K+ (aq) + SO4

2-(aq) Electrolytes. What are they? Nonelectrolytes. What are they and how do you recognize them? Electrolyte’s Strength – Strong vs Weak. Difference? Strong Electrolyte Weak Electrolyte Nonelectrolyte Ionic All soluble (or liquid) None None Molecular Strong acids (See Table 4.2) Weak acids Weak bases All other compounds

4.2 Precipitation Reaction AX(aq) + BZ(aq) → AZ + BX

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Solubility Rules (Handout) Problem 1 Soluble or Insoluble? K2CO3 BaSO4 NaOH CuSO4 FePO4 Ca(C2H3O2) 2 Precipitation Reaction

KI(aq) + AgNO3(aq) → KNO3(aq) + AgI(s)

A solid (precipitate) forms out of two aqueous solutions. Process for Predicting the Products of a Precipitation Reaction

• For Reactants: Determine what ions each aqueous reactant has Ex. Mix: AgNO3 (aq) + KCl (aq) Ions present: Ag+ (aq), NO3

- (aq), K+ (aq), Cl- (aq)

• To get the Products: Exchange ions (+) ion from one reactant with (-) ion from other AgCl and KNO3 (Write correct formulas: charges!! Neutral overall.)

• Determine solubility of each product in water and write their phases. – Solubility rules – If product is insoluble or slightly soluble, it will precipitate.

AgCl (s), KNO3(aq)

• Balance the equation – Count atoms

AgNO3(aq) + KCl(aq) AgCl(s) + KNO3(aq) Example 1

Mix: K2CrO4 (aq) + Ba (NO3)2 (aq)

K2CrO4 (aq) + Ba(NO3)2 (aq) → BaCrO4 (s) + 2 KNO3 (aq) Example 2

Mix: KNO3 (aq) + BaCl2 (aq)

KNO3 (aq) + BaCl2 (aq) → Ba(NO3)2 (aq) + 2 KCl (aq)

All the dissolved ions remain in solution. NR! (No chemical change takes place.)

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Ionic Equations What are the three types of equation we can write for reactions in aqueous solution? How are they different? Example 1:

• Molecular:

Na2CO3(aq) + BaCl2(aq) 2NaCl(aq) + BaCO3(s)

• Ionic:

2Na+(aq) + CO3

2-(aq) + Ba2+

(aq) + 2Cl-(aq) 2Na+(aq) + 2Cl-(aq) + BaCO3(s)

• Net ionic:

CO32-

(aq) + Ba2+(aq) BaCO3(s)

(cancel: Na+ and Cl- - unchanged) What is a spectator ion and how does it relate to net ionic equations? Example 2:

• Molecular: 2AgNO3(aq) + Na2CrO4(aq) → Ag2CrO4(s) + 2NaNO3(aq)

• Ionic:

2Ag+(aq) + 2NO3

-(aq) + 2Na+

(aq) + CrO42-

(aq) → Ag2CrO4(s) + 2Na+(aq) + 2NO3

-(aq)

• Net ionic:

2Ag+(aq) + CrO4

2-(aq) → Ag2CrO4(s) (cancel: Na+ and NO3

- - unchanged) Problem 2 Aqueous nickel (II) nitrate is added to aqueous potassium carbonate, is there a reaction? If so, write the molecular, total ionic and net ionic equations for it.

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Problem 3 Write the net ionic equation for: a. FeCl3 (aq) + NaOH (aq) b. Ba(NO3)2 (aq) + (NH4)3PO4 (aq) c. NaCl (aq) + Cu(NO3)2 (aq)

4.3 Acids, Bases, and Neutralization Reactions

Acids – How do we recognize them? Difference between Arrhenius and Brønsted-Lowry definitions? Bases – T How do we recognize them? Difference between Arrhenius and Brønsted-Lowry definitions? What is the difference between strong and weak acids and bases? Memorize the strong acids and bases. Neutralization Reactions. How do we identify them?

CH3COOH (aq) + NaOH (aq) CH3COONa (aq) + H2O (l)

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The net ionic equation for the reaction of a strong acid and strong base is always the same, why?

HCl (aq) + NaOH (aq) NaCl (aq) + H2O (l)

H+ (aq) + OH- (aq) H2O (l) What is a polyprotic acid and how do they behave chemically?

Mix equal # of moles: 1 mol NaOH + 1 mol H2SO3 → H2O + NaHSO3

Using excess base: 2 NaOH + H2SO3 → H2O (l) + Na2SO3 (aq)

Write the molecular, complete and net ionic equations for the reaction of acetic acid and potassium hydroxide. HC2H3O2 (aq) + KOH (aq) → KC2H3O2 (aq) + H2O (l)

HC2H3O2 (aq) + K +(aq) + OH-

(aq) → K +(aq) + C2H3O2

-(aq) + H2O (l)

HC2H3O2 (aq) + OH-(aq) → C2H3O2

-(aq) + H2O (l)

Problem 4 Write the molecular, total ionic, and net ionic equations for the reactions of H3PO4 with excess NaOH. (Hint: This is a weak acid.) Gas-Forming Reactions Memorize

H2CO3 (aq) H2O (l) + CO2 (g) ; H2SO3 (aq) H2O (l) + SO2 (g) ; NH4OH (aq) H2O (l) + NH3 (g)

• Carbonates and bicarbonates with acids produce H2CO3.

CaCO3 (s) + HCl (aq) CaCl2 (aq) + CO2 (g) + H2O (l)

NaHCO3 (aq) + HBr (aq) NaBr (aq) + CO2 (g) + H2O (l) H2CO3 (aq) Baking soda is used to:

- Neutralize acid spills. - Make baked goods rise.

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• Here, the expected product H2SO3 (aq), but it decomposes.

SrSO3 (s) + 2 HI (aq) SrI2 (aq) + SO2 (g) + H2O (l)

• Ammonium salt + strong base. Here, NH4OH (aq), decomposes:

NH4Cl (aq) + NaOH (aq) H2O (l) + NH3 (g) + NaCl (aq)

• Hydrogen sulfide (H2S) has a low solubility in water; therefore, reactions that produce it will form a gas (memorize this fact); the gas smells like rotten egg.

Na2S (aq) + H2SO4 (aq) Na2SO4 (aq) + H2S (g) Problem 5 Predict the products. Write the molecular, total and net ionic equations. a. b.

4.4 Oxidation-Reduction Reactions Describe the processes of oxidation and reduction. How do we use oxidation numbers?

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Oxidation Numbers

• Elements in natural elemental state have oxidation number = 0.

• The oxidation number of a monatomic ion is the same as its charge.

• Nonmetals tend to have negative oxidation numbers, although some are positive in certain compounds or ions.

➢ Oxygen has an oxidation number of −2, except in the peroxide ion in which it has an oxidation number of −1.

➢ Hydrogen is −1 when bonded to a metal, +1 when bonded to a nonmetal. ➢ Fluorine always has an oxidation number of −1. ➢ The other halogens have an oxidation number of −1 when they are negative; they can have positive

oxidation numbers, however, most notably in oxyanions.

• The sum of the oxidation numbers in a neutral compound is 0.

• The sum of the oxidation numbers in a polyatomic ion is the charge on the ion. Problem 6 Determine the oxidation numbers: Li2O H3PO4 MnO4

– Cr2O7

2– C7H8 Is this a redox reaction? CH4(g) + 2 O2(g) → CO2(g) + 2 H2O(g) Is this a redox reaction? CO2(g) + H2(g) → 2 CO(g) + H2O(g)

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Problem 7 Identify what’s oxidized and what’s reduced: Cu(s) + 2 NO3

– + 4H+ → Cu2+ + 2 NO2 + 2 H2O Single Displacement Reactions

A + BZ → AZ + B

There are two possible outcomes to this type of reaction, which are they and what id the meaning behind each? Example 1 Cu (s) + HgCl2 (aq) CuCl2 (aq) + Hg (l) BUT Hg (l) + CuCl2 (aq) NO Reaction!! Orange colorless pale blue silvery or black

Which one is more active? Example 2

Cu (s) + 2 Ag+ (aq) Cu2+ (aq) + 2 Ag (s) (net ionic eq.) Displacement Reactions

1. Metal and aqueous solution Fe(s) + CuSO4(aq) → FeSO4(aq) + Cu(s) Fe(s) + Cu2+(aq) → Fe2+(aq) + Cu(s) (Net Ionic)

2. Metal and aqueous acid solution Fe(s) + 2 HCl(aq) → FeCl2(aq) + H2(g) Fe(s) + 2 H+(aq) → Fe2+(aq) + H2(g) (Net Ionic)

• Metals less active than (H) show no reaction: Au(s) + H2SO4(aq) → NR

3. Active metal and water Ca(s) + 2 H2O(l) → Ca(OH)2(aq) + H2(g) Use the Activity Series to Make Predictions (No need to memorize.)

Problem 8 Predict whether a single displacement reaction will occur. Write the equation for the reaction. (Use the activity series) a. Cu(s) + Pb(NO3)2(aq) b. Pb(s) + HCl(aq)

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4.5 Concentrations of Solutions

𝑀𝑜𝑙𝑎𝑟𝑖𝑡𝑦 (𝑀) =𝑚𝑜𝑙𝑒𝑠 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑒

𝑣𝑜𝑙𝑢𝑚𝑒 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛 𝑖𝑛 𝑙𝑖𝑡𝑒𝑟𝑠

Problem 9 If you dissolve _____________ g NaCl in _____________ mL of solution, what is the molarity? Using Molarity Three values, simply solve for the unknown. Problem 10 What volume of _____________ M HCl solution contains _____________ g HCl? Problem 11 How many grams of K2Cr2O7 are there in _____________ mL of _____________ M K2Cr2O7? How are molarity and dissociation related? Problem 12 If the concentration of a solution of Cr2(SO4)3 is 0.300 M, what are the concentrations of all ions in solution? Dilution What does it refer to? What equation can we derive from this concept?

Molarity (M) =Moles of solute

Liters of solution

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Problem 13 What volume of __________ M NaOH needs to be diluted to prepare 5.00 L if 0.10 M NaOH?

4.6 Solution Stoichiometry and Chemical Analysis

a A + b B → c C + d D

Problem 14 What mass of silver bromide is produced from the reaction of _____________ mL of ____________ M aluminum bromide with excess silver nitrate solution?

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Titration What do we use it for? Why should we try to get the lightest color change from the indicator? Problem 15 In a titration, ______________ mL of ______________ M H2SO4 requires 32.59 mL of NaOH to reach the endpoint. What is the concentration of the NaOH? Problem 16 How many mL of _____________ M HI solution are needed to reduce 22.5 mL of a 0.374 M KMnO4 solution?

16 HI + 2 KMnO4 → 5 I2+ 2 MnI2 + 2 KI + 8 H2O Problem 17 An impure sample containing oxalic acid (H2C2O4) weighs ______________ g. If 32.18 mL of 0.3526 M NaOH is needed to reach the endpoint, what was the mass-% of oxalic acid in the mixture? Problem 18 An unknown monoprotic acid is titrated with NaOH. It requires ____________ mL of 0.1003 M NaOH to titrate 0.8359 g of the acid to the endpoint. What is the molar mass of the acid?

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Problem 19 In the precipitation reaction between ____________ mL of ___________ M lead(II) nitrate and 45.0 mL of 0.287 M potassium chloride, how much product is produced? Problem 20 50.00 mL of 1.500 M NaOH and _____________ mL of 1.000 M FeCl3 are mixed. What mass of ppt forms? What are the ion concentrations after the reaction?

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13.4 Expressing Solution Concentration Mass Percentage

𝑚𝑎𝑠𝑠 % 𝐴 =𝑚𝑎𝑠𝑠 𝑜𝑓 𝐴 𝑖𝑛 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛

𝑡𝑜𝑡𝑎𝑙 𝑚𝑎𝑠𝑠 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛× 100 ; 5.0 % 𝑜𝑓 𝑁𝑎𝐶𝑙 =

5.0 𝑔 𝑜𝑓 𝑁𝑎𝐶𝑙

100 𝑔 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛

Parts per Million (ppm)

𝑝𝑝𝑚 ==𝑚𝑎𝑠𝑠 𝑜𝑓 𝐴 𝑖𝑛 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛

𝑡𝑜𝑡𝑎𝑙 𝑚𝑎𝑠𝑠 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛× 106

2 𝑝𝑝𝑚 𝑃𝑏2+ =2 𝑔 𝑃𝑏2+

106 𝑔 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛 𝑜𝑟

2 𝑚𝑔 𝑃𝑏2+

𝑘𝑔 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛

Since a dilute solution is mostly water (d = 1g/mL).

2 𝑚𝑔 𝑃𝑏2+

𝐿 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛

Parts per Billion (ppb)

𝑝𝑝𝑏 ==𝑚𝑎𝑠𝑠 𝑜𝑓 𝐴 𝑖𝑛 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛

𝑡𝑜𝑡𝑎𝑙 𝑚𝑎𝑠𝑠 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛× 109

2 𝑝𝑝𝑚 𝑃𝑏2+ =2 𝑔 𝑃𝑏2+

109 𝑔 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛 𝑜𝑟

2 𝜇𝑔 𝑃𝑏2+

𝑘𝑔 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛

2 𝜇𝑔 𝑃𝑏2+

𝐿 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛

ppm and ppb are used for very dilute solutions (toxins). Mole Fraction (X):

𝑋𝐴 =𝑚𝑜𝑙𝑒𝑠 𝑜𝑓 𝐴

𝑡𝑜𝑡𝑎𝑙 𝑚𝑜𝑙𝑒𝑠 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛

Moles in solution = moles of solute + moles of solvent. Mole fractions of all components add up to 1. Molarity (M) What are the problems with it?

𝑀 =𝑚𝑜𝑙 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑒

𝐿 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛 ; 3.0 𝑀 𝐻𝐶𝑙 =

3.0 𝑚𝑜𝑙 𝐻𝐶𝑙

1 𝐿 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛

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Mass of solvent

+

Mass of solute

Mass of solution

Density

Volume of solution

Moles of solute

Molar mass

Molality(mol/kg solvent)

Molarity(mol/L solution)

Molality (m)

𝑚 =𝑚𝑜𝑙 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑒

𝑘𝑔 𝑜𝑓 𝑠𝑜𝑙𝑣𝑒𝑛𝑡 ; 1.5 𝑀 𝐻𝐶𝑙 =

1.5 𝑚𝑜𝑙 𝐻𝐶𝑙

1 𝑘𝑔 𝑜𝑓 𝑠𝑜𝑙𝑣𝑒𝑛𝑡

Changing Molarity to Molality Problem 21 You dissolve ____________ g of sugar (C12H22O11) in 250.00 g of water. Calculate the mole fraction, molality and mass percent of the solution. (Why can’t we calculate M?) Problem 22 The maximum allowable concentration of As in drinking water in the US is ___________ ppm.

a. If a 2.5 L water sample is found to contain 0.030 mg As, is this within the allowable limit? b. What is the allowable mass of As in 1 large glass of water (500 mL)?

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Problem 23 A sucrose solution is ___________% sucrose (C12H22O11) by mass. Calculate the molality of the solution. Problem 24 An aqueous solution of urea ((NH2)2CO) is ___________ m. Calculate M for this solution. (dsol = 1.045 g/mL) Problem 25 An aqueous solution of urea is ___________ M and has a density of 1.029 g/mL. What is its molality? M and ppm

1 𝑝𝑝𝑚 =1 𝑔 𝑋

1 × 106 𝑔 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛=

1 × 10−6𝑔 𝑋

1 𝑔 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛=

1 𝜇𝑔 𝑋

1 𝑔 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛

𝐹𝑜𝑟 𝑤𝑎𝑡𝑒𝑟: 𝑑 =1 𝑔

1 𝑚𝐿 ; 1 𝑝𝑝𝑚 𝑋 =

1 𝜇𝑔 𝑋

1 𝑚𝐿 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛 ; 𝑀[=]

𝑚𝑜𝑙

𝐿

𝟏 𝒑𝒑𝒎 =1 𝜇𝑔 𝑋

1 𝑚𝐿 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛 (1 𝐿

1000 𝑚𝐿)=

1,000 𝜇𝑔 𝑋

1 𝐿 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛=

𝟏 𝒎𝒈 𝑿

𝟏 𝑳 𝒔𝒐𝒍𝒖𝒕𝒊𝒐𝒏

ppm ⇔ M

𝑺𝒊𝒏𝒄𝒆: 𝟏 𝒑𝒑𝒎 =𝟏 𝒎𝒈 𝑿

𝟏 𝑳 𝒔𝒐𝒍𝒖𝒕𝒊𝒐𝒏

To change ppm’s to M, convert the mg’s to moles. To change M to ppm’s, convert the moles to mass in grams, then to milligrams.