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AP Chemistry
Chapter 4:Aqueous Reactions & Solution
Stoichiometry
Solutions
• Homogeneous mixtures• Have 2 parts:– Solute• Smaller quantities• Dissolved in the solvent
– Solvent• Larger quantities• Dissolves the solute(s)• Often water
Electrolytic Properties
• Electrolytes– Aqueous solutions that contain ions– Have ionic compounds or strong acids or bases as
solute– Will conduct electricity
• Nonelectrolytes– Aqueous solutions that do not contain ions– Have molecular compounds as solute
Ionic Compounds in Water
Ions dissociate from each other in water
Ions attracted to one another to form a crystal
lattice
Ionic Compounds in WaterWater is polar
Negative end of wateris attracted to positive ion
Each ion will besurrounded bywater molecules:Process is called solvation
Water prevents ionsfrom recombiningto form a crystal
Ionic Compounds in Water
Molecular Compounds in Water
Molecules remain in tact
Ions not usually formed
Strong acids are exceptions
Strong vs. Weak ElectrolytesWeak Electrolytes:• Exist in solution
almost entirely as molecules (a small fraction may be ions)
• Molecular compounds fall into this category
Strong Electrolytes:•Exist in solution completely (or nearly completely) as ions
•Essentially all soluble ionic compounds & some molecular (some acids) fall into this category
Strong vs. Weak ElectrolytesStrong Electrolyte Example:
HCl(aq) H+(aq) + Cl-(aq)
have no tendency to recombine to form HCl
Weak Electrolyte Example:
HC2H3O2(aq) H+(aq) + C2H3O2-(aq)
reaction is significant in both directions is said to be in chemical equilibrium
Soluble Ionic Compounds are Strong Electrolytes
(metal & nonmetal or contain ammonium ion)
Strong vs. Weak ElectrolyteIS NOT the same thing as solubility….compounds can be very soluble and be a weak electrolyte and vice versa….
POINTS
TO
REMEMBER
Sample Exercise 4.1The diagram represents an aqueous solution of one of the following compounds: MgCl2, KCl, or K2SO4. Which solution does it best represent?
2-
2-
2-
2-+ +
+
+
++
+
+
Precipitation Reactions
Lead nitrate + Potassium Iodide
Precipitation Reactions
• Result in the formation of an insoluble compound
• A precipitate is an insoluble compound formed by a reaction in solution
Solubility Rules
HAS CANHalides
Except:Hg2
2+
AgPb All Alkali
metals
Sulfates
Except: Ba2+, Sr2+, Hg22+, Ag, and Pb
All AmmoniumcompoundsAll C2H3O2
- compounds
All Nitrates
Hydroxides and Sulfides are soluble if
bonded with Ca2+, Ba2+, or Sr2+, alkali
metals or ammonium
Precipitation ExampleWill a precipitate form when solutions of Mg(NO3)2 and NaOH are mixed?
Sample Exercise 4.2Classify the following ionic compounds as soluble or insoluble in water (a) sodium carbonate (b) lead sulfate
Metathesis Reactions
• Cation and anion appear to “change partners”• Precipitation reactions conform to this pattern
AY + BX AX + BY
Completing and Balancing Metathesis:
1. Use chemical formulas of reactants to determine the ions that are present
2. Write the chemical formulas for the products by combining the cation from one reactant with the anion form the other reactant
3. Balance the equation
* If all products are soluble, we say NO REACTION has occurred.
Sample Exercise 4.3
(a) Predict the identity of the precipitate that forms when solutions of BaCl2 and K2SO4 are mixed.
(b) Write the balanced equation for this reaction.
Ionic Equations
• Molecular equations show the chemical formulas for the reactants and the products
• Complete ionic equations show the ions in solution
Pb(NO3)2 (aq) + 2KI (aq) PbI2 (s) + 2KNO3 (aq)
Pb2+ (aq) +2NO3- (aq) + 2K+ (aq) + 2I- (aq) PbI2 (s) +2K+ (aq) +2NO3
- (aq)
Ionic Equations
• Spectator ions– appear in both the reactants and the products in
the equation– play no part in the reaction– can be omitted (cancelled out)
Ionic Equations
• Net ionic equations show only the ions and molecules that are directly involved in the reaction
Pb2+ (aq) +2NO3- (aq) + 2K+ (aq) + 2I- (aq) PbI2 (s) +2K+ (aq) +2NO3
- (aq)
Pb2+ (aq) + 2I- (aq) PbI2 (s)
NET IONIC EQUATION:
Reminders
Charge is
conserved: Sum
of charges of
reactants must
equal sum of
charges of
products
If every ion is a spectator, no reaction occurs
To write net ionic equations:
1. Write a balanced molecular equation for the reaction
2. Rewrite to show ions that form in solution when each soluble, strong electrolyte dissociates into its component ions (Only strong electrolytes written in ionic form)
3. Identify and cancel the spectator ions
Sample Exercise 4.4Write the net ionic equation for the precipitation reaction that occurs when solutions of calcium chloride and sodium carbonate are mixed.
Acid – Base Reactions• Acids:• Substances that ionize in solutions to from H+ ions• H+ is simply a proton• Acids often called proton donors• Monoprotic acids• Have only one H • Examples: HCl, HNO3
• Diprotic acids• Ionized in 2 steps• Examples:H2SO4
H2SO4(aq) H+(aq) + HSO4-(aq)
HSO4-(aq) H+(aq) + SO4
2-(aq)
Acid-Base Reactions
• Bases– Substances that accept protons (H+)– Produce hydroxide ions (OH-) when dissolved in water– Include ionic hydroxide compounds• NaOH, MgOH, etc.
– NH3(ammonia) is also a base…it accepts H+ from water leaving OH-
• NH3(aq) + H2O(l) NH4+(aq) + OH-(aq)
Only a small fraction will form NH4+ ammonia is a
weak electrolyte
Acid-Base Reactions
• Acids and bases that are strong electrolytes are called strong acids and bases
• Reactivity can be determined by strength of acid or on the strength of the anion– Example: HF is a weak acid, but is very reactive
because of the F- ion
STRONG ACIDSH2SO4
HNO3
HClO3
HClO4
HCl
HBr
HI
STRONG BASESGroup I Metal
Hydroxides
Ca, Sr, & BaHydroxides
Sample Exercise 4.5The following diagrams represent aqueous solutions of three acids (HX, HY, HZ) with water molecules omitted for clarity. Rank them from strongest to weakest.
+
-
-
+
HX HY HZ
+
+
+
+
+
+
+
+--
-
-
--
-
-
-
--+
+
+
Type of Compound
StrongElectrolyte
Weak Electrolyte Nonelectrolyte
Ionic All None None
Molecular Strong Acids
Weak Acids & Bases
All others
Electrolytic Behavior of Soluble Compounds
Sample Exercise 4.6Classify each of the following substances as strong electrolyte, weak electrolyte, or nonelectrolyte: CaCl2, HNO3, C2H5OH (ethanol), HCHO2 (formic acid), KOH
Acids vs. Bases
• Taste sour• Turn blue litmus paper red
Acids
• Taste bitter• Turn red litmus paper blue
Bases
NeutralizationReaction between acidic and basic solutions
If the base is a metal hydroxide
Products are a salt and water
Compound whose cation comesfrom a base and anion comesfrom an acid
Neutralization Example
HCl(aq) + NaOH(aq) NaCl(aq) + H2O(l)
H+(aq) + Cl-(aq) + Na+(aq) + OH-(aq) Na+(aq) + Cl-(aq) + H2O(l)
H+(aq) + OH-(aq) H2O(l)
Molecular Equation:
Complete Ionic Equation:
Net Ionic Equation:
Sample Exercise 4.7(a) Write a balanced molecular equation for the reaction between aqueous solutions of acetic acid (HC2H3O2) and barium hydroxide [Ba(OH)2]. (b) Write the net ionic equation for this reaction
Gas Forming Reactions• Acids will form gases when they react with– Sulfides…H2S (g) forms
– Carbonates & bicarbonates…CO2 (g) forms• First to form is the unstable carbonic acid (H2CO3) which
decomposes to form carbon dioxide gas and water vapor
• Example: Alka Seltzer: NaHCO3 reacts with stomach acid
• Sodium bicarbonate & sodium carbonate are used to clean up acid spills
Oxidation-Reduction Reactions
• Reactions in which electrons are transferred between substances
• Oxidation– Loss of electrons– Substance becomes more positive
• Reduction– Gain of electrons– Substance becomes more negative
Oxidation – Reduction Reactions
• Oxidation of one substance is ALWAYS accompanied by the reduction of another
Oxidation Numbers
• Assigned to each atom in a neutral substance or in a charged species
• This is the charge of the monatomic ions or the hypothetical charge assigned to the atom assuming the electrons are completely held by one atom or another
Oxidation Numbers
• In redox reactions, the ox. #’s change from products to reactants
• Oxidation = increase in oxidation number
• Reduction = decrease in oxidation number
Assigning Oxidation Numbers…the Rules
1. Oxidation number is 0 for atoms in elemental form H2 H = 0 S8 S = 0 P4 P = 0
2. Oxidation number = the charge of monatomic ionsTo distinguish between oxidation number and
actual charge: Ca2+ oxidation number = +2
3. Nonmetals (usually negative ox. nos.)a) Oxygen is usually -2 (both ionic & molecular)
Exception: in peroxide, each oxygen is -1b) Hydrogen is +1 when bonded with nonmetals and -1
when bonded with metalsc) Fluorine is ALWAYS -1d) Other halogens are usually -1
Exception: When bonded with oxygen, they are +
4. Sum of oxidation numbers in a compound is always 0: in a polyatomic ion the sum = the charge of the ion
Assigning Oxidation Numbers…the Rules
Sample Exercise 4.8Determine the oxidation number of sulfur in each of the following: (a) H2S (b) S8 (c) SCl2 (d) Na2SO3 (e) SO4
2-
Oxidation of Metals by Acids & Salts
• When a metal is oxidized, it appears to be “eaten away” as it reacts
• Called displacement reactions because the ion in solution is displaced (replaced) by the oxidation of an element
Oxidation of Metals by Acids and Salts• Many metals will react with acids to form hydrogen
gas and a salt
• For example: Mg(s) + 2 HCl(aq) MgCl2(aq) + H2(g)
ox. #’s 0 +1 -1 +2 -1 0
Mg is oxidizedH is reduced
Cl is a spectator
Oxidation of Metals by Acids & Salts
Fe (s) + Ni(NO3)2 (aq) Fe(NO3)2 (aq) + Ni (s)
One more time:
If something is oxidized….Something else must
be reduced!
Sample Exercise 4.9Write the balanced molecular and net ionic equations for the reaction of aluminum with hydrobromic acid.
The Activity Series
• A list of metals arranged in order of decreasing oxidation
• Can be used to predict reactions between metals and metals salts or acids
• Any metal on the list can be oxidized by ions of elements below it
• Most easily oxidized = most reactive is at the top of the list
• List is on page 143 of your textbook
Sample Exercise 4.10Will an aqueous solution of iron (II) chloride oxidize magnesium metal? If so, write the balanced molecular and net ionic equations for the reaction.
Concentrations of Solutions
• Designates the amount of solute dissolved in a given amount of solvent or solution
• Molarity– Quantitative expression of concentration– M = mol/L – Moles of solute/ Liter of solution
Sample Exercise 4.11Calculate the molarity of a solution made by dissolving 23.4 g of sodium sulfate in enough water to form 125 mL of solution.
Expression of Concentration
• Can be made in terms of the compound:– 1.0 M Na2SO4
• Can be made in terms of the ions:– 2.0 Na+ and 1.0 M SO4
2-
Sample Exercise 4.12What are the molar concentrations of each of the ions present in a 0.025 M aqueous solution of calcium nitrate?
Sample Exercise 4.13How many grams of Na2SO4 are required to make 0.350 L of 0.500 M Na2SO4?
Dilution
• Preparation of solutions of lower concentrations from solutions of higher concentrations by adding water
• The addition of water (solvent) does not change the # of moles of solute
• Equation:M1V1 = M2V2
Sample Exercise 4.14How many milliliters of 3.0 M H2SO4 are needed to make 450 milliliters of 0.10 M H2SO4?
Sample Exercise 4.15How many grams of Ca(OH)2 are needed to neutralize 25.0 mL of 0.100 M HNO3?
Titration
• Method used to determine the amount of solute in a solution
• Uses the addition of a solution of known concentration called a standard solution
• Can involve different types of reactions– Acid base– Redox– precipitation
Titration• Makes use of equivalence point– The point where stoichiometrically equivalent quantities
are brought together• Use indicators– Indicates when equivalence point is reached– Changes color in response to change in pH– Ex. Phenolphthalein is colorless in acid, but pink in a
base– End point of titration is when solution changes color– Must choose an indicator that’s end point coincides with
the equivalence point
Sample Exercise 4.16The quantity of Cl- in a municipal water supply is determined by titrating the sample with Ag+. The reaction taking place during the reaction is: Ag+(aq) + Cl-(aq) AgCl(s)The end point in this type of titration is marked by the color change of a special type of indicator. (a) How many grams of chloride ion are in a sample of the water if 20.2 mL of 0.100 M Ag+ is needed to react with all the chloride in the sample? (b) If the sample has a mass of 10.0 grams, what percent Cl- does it contain?
Sample Exercise 4.17One commercial method used to peel potatoes is to soak them in a solution of NaOH for a short time, remove them from the NaOH, and spray off the peel. The concentration of NaOH is normally in the range of 3 to 6 M. The NaOH is analyzed periodically . In one such analysis, 45.7 mL of 0.500 M H2SO4 is required to neutralize a 20.0 mL sample of NaOH solution. What is the concentration of the NaOH solution?
Integrative SampleA sample of 70.5 mg of potassium phosphate is added to 15.0 mL of 0.050 M silver nitrate, resulting in the formation of a precipitate. (a) Write the molecular equation for the reaction (b) What is the limiting reactant in the reaction? (c) Calculate the theoretical yield, in grams, of the precipitate that forms.
That’s all Folks!