33 solution stoichiometry

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Aqueous Reactions and Solutions

Stoichiometry

•

Types of Aqueous Reactions

–

Precipitation

_{Precipitation}

–

Acid/Base

•

Solution Stoichiometry

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Definitions

 Solution – homogeneous mixture

 Solvent – part of above in most amount

 Solute – part of above in least amount

 Electrolyte – solution that conducts a current,

showing the presence of ions

 Nonelectrolyte – solution that doesn’t

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Ionic compounds

 Why do they dissolve in water?

(4)

Molecular Compounds

 Polar – stay in tact but are separated by

polar water molecules (sugar)

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Strong Electrolytes

 Describes the amount of ions in solution

 Ionic compounds produce strong electrolytes

by almost 100% dissociation

 Acids – react with water to for ions.

 Called ionization

 Strong acids form strong electrolytes

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Weak Electrolytes

 Weak acids form weak electrolytes because

they form an equilibrium and ionize very little

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Equations

 Dissociation of ionic compounds must reflect

the number of ions in the formula

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Precipitation Reactions

 Result in the formation of an insoluble product

 Pb(NO₃)_2(aq) + 2KI_(aq) → PbI₂_(s) + 2KNO₃_(aq)

 Occurred because certain pairs of oppositely charged

ions attract each other so strongly that they form an insoluble solid

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Rules

 “Insoluble” means there is less than 0.01

moles of the substance dissolved in a liter of solution

 Have your solubility rules with you.

 NOTE: All compounds of the alkali metal

ions and NH₄+ are soluble

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Metathesis or Exchange Reactions

 General term used when ions appear to

exchange and reform

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Ionic Equations

 Molecular:

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

 Ionic:

Pb2+ + 2NO

3- + 2K+ + 2I- → PbI2(s) + 2NO3- + 2K+

 Net Ionic: Show only the ones that react!

Pb2+ + 2I- → PbI

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Acid –Base Reactions

 Important in the body and in the environment

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Acids

 Substances that ionize in water to form H+

ions

 This is actually a proton

 Acids are called proton donors

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Bases

 Substances that accept protons or H+ ions.

 Two ways –

dissociate to form OH- (NaOH)

ionize water to leave OH- (NH

3)

Equations?

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Determining Strength of Electrolytes

 Ionic? Strong if yes

 Molecular?

Strong acid – strong

Not a strong acid but contains “H” – weak

No “H” and not NH₃ – nonelectrolyte

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Neutralization

 Occurs when an acid and a base are mixed

 Strong acid + strong base forms a salt and

water

 NaOH + HCl → NaCl + H₂O

 Salt – compound made between an anion

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Reaction of carbonates and sulfides

with acids

 Gas formation

 2HCl(aq) + Na₂S(aq) → H₂S(g) + 2NaCl(aq)

 _{Net ionic?}

 HCl(aq) + NaHCO₃(aq) → NaCl(aq) + H₂CO₃(aq)  H₂CO₃ decomposes rapidly to H₂O and CO₂

 Net ionic?

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Net Ionic Equations for each acid-base

combination

 SA + SB: HCl + NaOH

 WA + SB: CH₃COOH + NaOH

 SA + WB: HCl + NH₃

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Oxidation-Reduction Reactions

 Transfer of electrons

 Oxidation – Loss of electrons

Mg → Mg2+ +

2e- Reduction – gain of electrons

O₂ + 4e- → 2O

2-2Mg(s) + O2(g)→ 2MgO(s)

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Bookkeeping

 Oxidation numbers

1. All free elements are “0”

2. Monatomic ions – ox # is charge

3. Oxygen: usually -2 except in O₂2-, which is -1.

4. Hydrogen is +1 when bonded to nonmetals and -1 when bonded to metals.

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Types

 Oxidation of metals by an acid or salt – called

single replacement reactions

Zn(s) + 2HBr(aq) → ZnBr₂(aq) + H₂(g)

Net ionic?

What is oxidized? This is called the reducing agent.

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Activity Series

 Listed as ease of oxidation

 Top is the most easily oxidized, or best

reducing agents.

 Bottom is most easily reduced, or best

oxidizing agents.

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Molarity

 Symbolized with “M”

 Defined as

moles of solute/volume of solution in liters

 2M is pronounced “2 molar” and means 2

moles of solute is dissolved in 1 liter of solution

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Preparation of Solutions

 Volumetric flask is used

 Calculated mass is put in flask  Water is added to fill line

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Molar Concentrations of Electrolytes

 Calculate molarity of entire species as before

 To find molarity of each ion, multiply by

coefficient of each in the balanced equation

 Example, in a 0.1 M solution of Na₂O, the

concentration of the Na+ ion is 0.2 M (animation)

 Symbolized by [ ]

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Interconverting

 Molarity can be used as a conversion factor

 Liters X (mol/liter) = moles

 Moles X (liters/mol) = liters

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Dilution Problems

 Calculate number of moles needed by

liters X (moles/liter)

 Calculate the volume of given solution that

will yield that number of moles by moles X (liters/moles) Animation

 Try one

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Solution Stoichiometry

 Calculate moles by solution calculation

 Look at balanced equation

 Do final calculation by regular stoichiometry

 Complete to volume if necessary

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Titration

 Lab procedure for calculating an unknown molarity

using a solution with a known molarity (standard solution)

 Standard solution is added to the unknown solution

using a buret (animation)

 Equivalence point is reached when stoichiometry

says quantities are equal

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Two Kinds of Titration

 Acid-Base reaction

 Redox Titration