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# Aqueous Reactions and Solutions Stoichiometry

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```									Chapter 4 – Aqueous Reactions and
Solutions Stoichiometry
• Types of Aqueous Reactions
– Precipitation
– Acid/Base
– Redox
• Solution Stoichiometry
• Concentrations of Solutions
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 – you guessed it
   Try sugar solution vs. salt solution
   animation
Ionic compounds

   Why do they dissolve in water?
   Called dissociation
   Animation
Molecular Compounds

   Polar – stay in tact but are separated by
polar water molecules (sugar)
   Nonpolar – do not dissolve (oil)
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
because they ionize completely
   HCl equation
Weak Electrolytes

   Weak acids form weak electrolytes because
they form an equilibrium and ionize very little
   Acetic acid equation
   Animation
Equations

   Dissociation of ionic compounds must reflect
the number of ions in the formula
   Try the dissociation of Al2(SO4)3
Precipitation Reactions

   Result in the formation of an insoluble product
   Demo – write the equation
   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
   Need to know rules! Masterson list is easier.
Rules

   “Insoluble” means there is less than 0.01
moles of the substance dissolved in a liter of
solution
   Memorize rules on page 118
   NOTE: All compounds of the alkali metal
ions and NH4+ are soluble
   NOTE: All nitrates are soluble
Metathesis or Exchange Reactions

   General term used when ions appear to
exchange and reform
   Comes from Greek word for transpose
   PE on page 119
Ionic Equations

 Molecular:
Pb(NO3)2(aq) + 2KI(aq) → PbI2(s) +
2KNO3(aq)
 Ionic:
Pb2+ + 2NO3- + 2K+ + 2I- → PbI2(s) + 2NO3- + 2K+
 Net Ionic: Show only the ones that react!
Pb2+ + 2I- → PbI2(s)
PE page 121
Acid –Base Reactions

   Important in the body and in the environment
   Aqueous solution
Acids

   Substances that ionize in water to form H+
ions
   This is actually a proton
   Acids are called proton donors
   Can be monoprotic (HCl) or diprotic (H2SO4)
Bases

   Substances that accept protons or H+ ions.
   Two ways –
dissociate to form OH- (NaOH)
ionize water to leave OH- (NH3)
Equations?
   Learn Table 4.2 on page 122
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
   Demo
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
from an acid and a cation from a base
   Net ionic?
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?
   What is the fizz in Alka seltzer?
Net Ionic Equations for each acid-base
combination

   SA + SB: HCl + NaOH
   WA + SB: CH3COOH + NaOH
   SA + WB: HCl + NH3
   Learn these general types.
Oxidation-Reduction Reactions

 Transfer of electrons
 Oxidation – Loss of electrons
Mg → Mg2+ + 2e-
 Reduction – gain of electrons
O2 + 4e- → 2O2-
2Mg(s) + O2(g) → 2MgO(s)
Leo says grrrrrr!
You need both in a reaction.
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.
5. Fluorine is -1 in all compounds.
6. Sum of all oxidation numbers is “0” for a compound
or equal to the charge of an ion.
7. Practice!
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.
What is reduced? This is called the oxidizing
agent.
Try for Copper in silver nitrate.
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.
   Predict if Cu can be oxidized by HCl.
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
   Try a problem
Preparation of Solutions

   Volumetric flask is used
   Calculated mass is put in flask
   Water is added to fill line
(animation)
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 [ ]
   [Na+] = 0.2 M
   Try one
Interconverting

   Molarity can be used as a conversion factor
   Liters X (mol/liter) = moles
   Moles X (liters/mol) = liters
   Try a problem
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
Solution Stoichiometry

   Calculate moles by solution calculation
   Look at balanced equation
   Do final calculation by regular stoichiometry
   Complete to volume if necessary
   Try one
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
   Indicator changes color at first sign of excess
of one reagent – this is called the end point
Two Kinds of Titration

   Acid-Base reaction
   Redox Titration
   Try some problems

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