Page 1 CHAPTER 4: R EACTIONS 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 CH 3 CH 2 OH stir Water CH 3 CH 2 OH (l) CH 3 CH 2 OH (l) CH 3 CH 2 OH (aq) stir Water NaCl (s) Na + Cl – NaCl (aq) NaCl (s)
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Page 1
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)
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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.
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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.
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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
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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.
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Sample Problem:
Write the balanced formula equation for this reaction. Include phase descriptors (s, l, g, aq).
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
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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)
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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?
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)
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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:
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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
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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)
What is oxidized? What is reduced?
What is the oxidizing agent? What is the reducing agent?