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CHAPTER 4: REACTIONS

AQUEOUS SOLUBILITY

DISSOLUTION IN WATER

Soluble molecular compounds (with oxygen or nitrogen) in solution…

Soluble ionic compounds (“salts”) in solution…

IONIC SOLUBILITY TRENDS

Grams of solute that dissolve in 100 mL water (25 ˚C)

RbNO3 KCl Li2CO3 Ca(OH)2 BaSO4 Mg3(PO4)2

Solubility g/100 mL H2O 65.0 35.5 1.30 0.160 0.00031 0.00009

CH3CH2OHstir

Water

CH3CH2OH (l)

CH3CH2OH (l)

CH3CH2OH (aq)

stir

Water

NaCl (s)Na+

Cl–

NaCl (aq)

NaCl (s)

Is each compound soluble (aq) or insoluble (s) in water?

CuCl2 (NH4)3PO4 CaSO4

KOH Ba(C2H3O2)2 NiCO3

ELECTROLYTIC SOLUTIONS

CONDUCTIVITY TESTS

IONS ALLOW FOR CONDUCTIVITY

AQUEOUS SOLUBILITY

Salts are Soluble with: (1) Group IA metal ions; (2) Ammonium ions; (3) Nitrates, Nitrites, Chlorates, and Perchlorates; (4) Acetates (except with aluminum and silver); (5) Halogens (except with silver, lead(II), and copper(I) ions); (6) Sulfates (except with barium, calcium, strontium, and lead(II) ions). Salts are Insoluble with (applies only when situations 1-6 are absent): (7) Hydroxides (except with barium); (8) Carbonate, Chromate, Phosphate, and Sulfite ions; (9) Sulfide ions (except with Group IIA metals); (10) Other ions not previously mentioned.

STRONG, WEAK, NON-ELECTROLYTES

Strong Electrolytes Weak Electrolytes Non Electrolytes

Sample Problem:

When sodium carbonate is added to water, the solution is highly conductive. However, when calcium carbonate is added to water, the mixture does not conduct electricity. Explain.

STRONG + WEAK ACIDS

PROBLEMS

Sample Problems:

Draw a beaker representation of the major solute component(s) in each aqueous solution. For acids, show several solute particles so as to distinguish between weak and strong acids.

HClO4 (aq) HC3H5O2 (aq) C6H12O6 (aq)

Strong Acids:

HCl, HBr, HI

HNO3, H2SO4

HClO3, HClO4

Weak Acids:

HF, HC2H3O2

H2CO3, H3PO4

many others

Sample Problem: Solutions of the following compounds are made by dissolving 0.5 mole solute in 500 mL water. Rank the resulting aqueous solutions in order of increasing conductivity. NaHCO3 CH3OH H3PO4 BaCl2

PRECIPITATION REACTIONS

WRITING FORMULA EQUATIONS

Formula Equation:

Na2CrO4 (aq) + AgNO3 (aq) →

→

Na2CrO4 (aq) AgNO3 (aq) Product mixture

Yellow Na2CrO4

Red solid

Clear AgNO3

“Black Smokers” Dissolved iron(II) sulfide emitted underwater from hydrothermal vents, which solidify when they cool to form large black chimneys.

Sample Problem:

Write the balanced formula equation for this reaction. Include phase descriptors (s, l, g, aq).

FeCl3 (aq) + Ca(OH)2 (aq) →

COMPLETE IONIC AND NET IONIC EQUATIONS

For this balanced equation:

Pb(NO2)2 (aq) + 2 NH4I (aq) → PbI2 (s) + 2 NH4NO2 (aq)

Complete Ionic Equation:

Net Ionic Equation:

Sample Problems:

Write the balanced formula equation and net ionic equation for this reaction:

Formula Eqn: AgNO3 (aq) + MgBr2 (aq) →

Net Ionic Eqn:

Write these equations for this reaction (done previously): FeCl3 (aq) + Ca(OH)2 (aq) →

Complete Ionic Eqn:

Net Ionic Eqn:

ACID-BASE REACTIONS

Acid Reflux Effect of acid rain on marble statues Acids and bases reacting

EQUATIONS

STRONG ACIDS

Formula Equation: HCl (aq) + KOH (aq) →


Net Ionic Equation:

WEAK ACIDS

Formula Equation: H3C6H5O7 (aq) + NaOH (aq) →


Net Ionic Equation:

Sample Problems:

Write the balanced formula equation for this reaction:

H2SO4 (aq) + NaOH (aq) →

Which is the net ionic equation for: HNO2 (aq) + CsOH (aq) → CsNO2 (aq) + H2O (l)

A: H+ (aq) + OH– (aq) → H2O (l)

B: HNO2 (aq) + OH– (aq) → H2O (l)

C: HNO2 (aq) + OH– (aq) → NO2– (aq) + H2O (l)

D: HNO2 (aq) + Cs+ (aq) + OH– (aq) → Cs+ (aq) + NO2– (aq) + H2O (l)

CARBONATES

Formula Equation: HCl (aq) + NaHCO3 (aq) →

Formula Equation: HNO3 (aq) + K2CO3 (aq) →

SOLUTION CONCENTRATION

GENERAL IDEA

Aqueous solutions of Cu(NO3)2

Chalk (CaCO3) reacts with vinegar to form bubbles. Limestone, marble (CaCO3) reacts with acidic solutions.

MOLARITY CALCULATIONS

Sample Problems:

Into a volumetric flask is placed 2.62 g BaCl2 and enough water to make 25.0 mL of solution. What is the molarity of the BaCl2 solution?

Molarity is also a conversion factor:

How many grams of H2SO4 are present in 50.00 mL of 6.00 M H2SO4?

What volume of 1.09 M NaCl contains 0.075 mol of NaCl?

Mass Moles Volume

DILUTION

Diluting Liquids: Which has more alcohol?

Shot of whiskey

Whiskey sour:

1 shot of whiskey 1 tsp. lime juice

3.75 oz. sweet + sour mix

Sample Problems:

When 30.0 mL of blood is diluted to 1.000 L, the concentration of NaCl is 0.00474 M. What is the concentration of NaCl in the original blood sample?

250.0 mL of water is added to 25.0 mL of a 0.750 M MgF2 solution. What is the concentration of fluoride ions in the resulting solution?

SOLUTION STOICHIOMETRY

ACID-BASE INDICATORS

Indicator reaction: HA (aq) + OH– (aq) → A– (aq) + H2O (l)

Goldenrod Indicator Phenolphthalein indicator

LABORATORY TITRATIONS

Titration: use of a burette to find precise amounts of reactants.

Ideally at the “endpoint” (indicator color change), a reaction is complete.

Sample Problem:

What volume of a 0.0200 M aqueous potassium hydroxide (in mL) is required to neutralize (completely react with) 26.50 mL of a 0.0150 M aqueous solution of oxalic acid (H2C2O4)?

Moles A Moles B Mass BVolume A

Sample Problems:

A 1.20 g sample of a monoprotic acid is dissolved in 25.0 mL water. A titration of this solution requires 15.27 mL of a 0.500 M NaOH solution to neutralize it. Calculate the molar mass of the acid.

A titration requires 35.70 mL of a 0.205 M KOH solution to neutralize a solution of sulfuric acid. How many moles of H2SO4 were present?

H2SO4 (aq) + 2 KOH (aq) → K2SO4 (aq) + 2 H2O (l)

OTHER SOLUTION REACTIONS

Sample Problems:

How many grams of MgCrO4 (MM = 140.31) can be produced from reaction of 150.0 mL of 0.0167 M MgCl2 with 250.0 mL of 0.0085 M K2CrO4?

MgCl2 (aq) + K2CrO4 (aq) → MgCrO4 (s) + 2 KCl (aq)

What volume (in mL) of 1.500 M Ca(OH)2 reacts completely with 43.5 mL of 0.1694 M HNO3?

2 HNO3 (aq) + Ca(OH)2 (aq) → Ca(NO3)2 (aq) + 2 H2O (l)

Sample Problem:

350.0 mL of 0.520 M lead(II) nitrate is reacted with excess potassium iodide (375.0 mL of a 2.50 M solution) to form a solid. Calculate the concentration of Pb2+ and I– present in solution after the precipitation is complete. Assume the volumes are additive when the solution is made.

Pb(NO3)2 (aq) + 2 KI (aq) → 2 KNO3 (aq) + PbI2 (s)

REDOX REACTIONS

Burning space shuttle fuel Burning natural gas Browning bananas Many violent reactions

GENERAL IDEA

“Redox” = Reduction – Oxidation Reaction

Rusting of nails: 4 Fe (s) + 3 O2 (g) → 2 Fe2O3 (s)

Each Fe atom and is . Fe undergoes .

Each Ox atom and is . Ox undergoes .

Redox reactions involve:

LEO the lion says GER

Loss of electrons = oxidation Gain of electrons = reduction

Oxidizing agent (Oxidizer):

Reducing agent:

OXIDATION STATES (TO TRACK REDOX)

Oxidation States / Oxidation Numbers: theoretical charge on an atom.

Assigning Oxidation States / Numbers Examples

Ionic compounds: oxidation state for monatomic ion is its charge.

Ba2+ = Cr in CrCl3 =

Elemental forms: oxidation state is zero. Na (s) = F2, H2, O2, N2 =

Molecular compounds or polyatomic ions: determine oxidation number relative to given ox # of F, O, or H.

Fluorine has an oxidation state of –1. F in HF =

Oxygen has an oxidation state of –2. (FYI, exception O is –1 in a peroxide, O22–) O in NaOH =

Hydrogen has an oxidation state of +1. H in H2S =

Sample Problems: Assign the oxidation states for each atom.

SF2 Cl2

Ox states

NH4+ NO2–

Ox states

Na2Cr2O7

Ox states

ANALYZING REDOX REACTIONS

Redox reactions involve changing charge or changing oxidation number.

Copper Plating on Iron

Fe (s) + Cu2+ (aq) → Fe2+ (aq) + Cu (s)

What is oxidized? What is reduced?

What is the oxidizing agent? What is the reducing agent?

Zinc metal in Acid Zn (s) + 2 HCl (aq) → ZnCl2 (aq) + H2 (g)



Aqueous Redox

5 Fe2+ (aq) + MnO4– (aq) + 8 H+ (aq) → Mn2+ (aq) + 5 Fe3+ (aq) + 4 H2O (l)

MnO4

– is deep purple; Fe2+ is pale green; Fe3+ is orange.



A total of ________ electrons are transferred from ________ to ________.

Concentrated Nitric Acid on a Copper Penny

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



Electrolysis of Salt (NaCl) solution

BALANCING REDOX REACTIONS – IN ACTIVITY

Growing Silver on Copper

Cu (s) + Ag+ (aq) → Cu2+ (aq) + Ag (s)

Half-Reaction Method:

• Separate the reaction into “half-reactions,” separating different elements (except O, H).

• Balance all atoms except hydrogen and oxygen.

• Balance oxygen by adding H2O to the necessary side.

• Balance hydrogen by adding H+ to the necessary side.

• Balance the charge by adding electrons (e–) to the necessary side.

• Equalize the (e–) lost by one half-reaction to the (e–) gained by the other half-reaction.

• Combine the half-reactions and simplify the reaction if possible (cancel terms).

• If in basic conditions, add OH– to neutralize any H+ appearing in the balanced reaction

(OH– + H+ → H2O) and cancel terms.

What reaction occurs to cause the paper to turn red?

A: 2 H2O (l) → O2 (g) + 4 H+ (aq)

B: 2 H2O (l) → H2 (g) + 2 OH– (aq)

C: Na+ (aq) → Na (s)

D: Cu (s) → Cu2+ (aq)

Is this reaction an oxidation or reduction?

Sample Problem:

Balance the following redox reaction under acidic and basic conditions.

Mn2+ (aq) + S2O82– (aq) → MnO4– (aq) + HSO4– (aq) peroxydisulfate

Half-Reaction 1 Half-Reaction 2

1/2

Elem

Ox

H

Chg

e–

Basic conditions: Add enough OH– to both sides to neutralize the H+ (H+ + OH– → )

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