Download Reactions In Aqueous Solution

Reactions In Aqueous Solution
Molarity
Most reactions that take place in the lab involve
ions or molecules dissolved in water (aqueous
solutions).
The concentration of a solute in solution can be
expressed in terms of its molarity:
molarity (M) = moles of solute / liters of solution
M = nsolute / Lsolution
The symbol [ ] is commonly used to represent
molarity of a species in solution.
Solutions that contain relatively large amounts of
solute per liter are referred to as concentrated,
while those containing relatively small amounts
are called dilute solutions.
Example
A bottle of dilute nitric acid contains 6.0 moles of
HNO3 per liter of solution (M = 6.0 mols/L). How
many moles of HNO3 are in 75 mL of this
solution? What volume of dilute nitric acid must
be taken to contain one mole of HNO3?
Example
Give the concentration, in moles per liter, of each
ion in (a) 0.080 M K2SO4 (b) 0.40 M FeCl3
Precipitate Reactions
When two different aqueous solutions of ionic
compounds are mixed, an insoluble solid
(precipitate) can separate out of solution.
To predict the products of a precipitate reaction
you must know which ionic substances are
insoluble in water.
Solubility Rules (GENERAL GUIDELINES)
Solubility rules that apply to water solution:
(1) All alkali metal and ammonium compounds are
soluble.
(2) All acetate, perchlorate, chlorate, and nitrate
compounds are soluble.
(3) Silver, lead, and mercury(I) compounds are
insoluble.
(4) Chlorides, bromides, and iodides are soluble.
(5) Carbonates, hydroxides, oxides, phosphates,
silicates, and sulfides are insoluble.
(6) Sulfates are soluble except for calcium and barium.
*These rules are to be applied in the order given. AgCl is
insoluble because rule 3 takes precedence over rule 4.
Example
Predict what will happen when the following pairs
of aqueous solutions are mixed.
(a) Cu(NO3)2 and (NH4)2SO4
(b) FeCl3 and AgNO3
Equations that show all species as molecules
(formula units) are called molecular equations.
Na2CO3(aq) + CaCl2(aq) --- CaCO3(s)+ Na2CO3(aq)
When all soluble species are shown as ions we
have an ionic equation (total ionic equation).
2Na+(aq) + CO3-2 (aq) + Ca+2(aq) + 2 Cl- (aq) ----
2Na+(aq) + 2 Cl-(aq) + CaCO3(s)
If you cancel the spectator ions (ions present on
both sides of the equation), which take no part in
the reaction, you are left with the net ionic
equation.
Ca+2(aq) + CO3-2(aq) --- CaCO3(s)
Net ionic equations must be balanced with respect
to mass (the number of atoms) and charge.
Example
Write net-ionic equations for any precipitation
reaction that occurs when solutions of the
following ionic compounds are mixed.
(a) NaOH and Cu(NO3)2
(b) BaCl2 and Ag2SO4
(c) (NH4)3PO4 and K2CO3
Example
When aqueous solutions of sodium hydroxide and
iron (III) nitrate are mixed a red precipitate
forms.
(a) Write a net ionic equation for the reaction.
(b) Determine the mass of the precipitate
formed when 30.0 mL of 0.125 M Fe(NO)3
reacts.
(c) Determine the mass of the precipitate
formed when 50.0 mL of 0.200 M NaOH
and 30.00 mL of 0.125 M Fe(NO3)3 are
mixed.
Acid-Base Reactions
An acid is a species that produces H+ ions in
aqueous solution, has a sour taste, and turns blue
litmus paper red.
A base is a substance that produces OH- ions in
aqueous solution, has a slippery feel, and turns
red litmus paper blue.
In this section we are using the Arrhenius
definition.
Strong acids ionize completely in water, forming
H+ ions and anions.
There are six common strong acids: HCl, HBr, HI,
HNO3, HClO4, H2SO4.
Example
HCl(aq) --- H+(aq) + Cl-(aq)
If 0.1 mole of HCl was added to water, there is 0.1
mole of H+ and Cl- ions and no HCl molecules.
Weak acids only partially ionize to H+ ions in
water.
HB(aq)
H+(aq) + B-(aq)
HF(aq)
H+(aq) + F-(aq)
The double arrow indicates that the reaction does
not go to completion. A mixture is formed
containing significant amounts of reactants and
products.
A strong base is completely ionized to OH- ions
and cations in a water solution.
Strong bases are generally hydroxides of group IA
and II A metals (LiOH, NaOH, KOH, Ca(OH)2,
Sr(OH)2, Ba(OH)2).
Example
Na(OH) -- Na+(aq)+ OH-(aq)
If 0.1 mole of NaOH is added to water, there is 0.1
mole of Na+ ions, 0.1 mole of OH- ions, and no
NaOH molecules.
Weak bases react with water molecules to acquire
H+ ions and leave OH- ions behind.
Example
NH3(aq) + H2O(aq)
NH4+(aq) + OH-(aq)
The reaction does not go to completion. If 0.1 mole
of ammonia is added to a liter of water, there is
about 0.001 mole of NH4+, 0.001 mole of OH- , and
0.099 mole of NH3.
Amines (R-NH2) are commonly weak bases.
CH3NH2(aq) + H2O(l)
CH3NH3+(aq) + OH-(aq)
Strong acids and bases are good conductors of
electrical current (ions in solution) and are strong
electrolytes. Weak acids and bases are weak
electrolytes.
Equations for Acid-Base Reactions
(a) A strong acid-strong base reaction is a
neutralization reaction.
H+(aq) + OH-(aq) --- H2O(l) (net-ionic equation)
Example
HNO3(aq) + NaOH(aq) ---- NaNO3(aq)+ H2O(l)
acid + base -- salt + water
(b) When a weak acid is added to a strong base
a two step process occurs.
(1) The weak acid is ionized.
HB(aq)
H+(aq) + B-(aq)
(2) The neutralization of the H+ ions from
step one by the OH- ions of the strong
base.
H+(aq) + OH-(aq) -- H2O(l)
The equation for the overall reaction is
determined by adding the two above reactions
together.
HB(aq)+ OH-(aq) --- B-(aq)+ H2O(l) (net ionic
equation)
Example
NaOH is added to HF
HF(aq)
H+(aq) + F-(aq)
H+(aq) + OH-(aq) -- H2O(l)
HF(aq)+ OH-(aq) --- F-(aq) + H2O(l)
(c) When a strong acid and a weak base are
reacted two steps take place.
Example
HCl is added to a sample of aqueous ammonia,
NH3
(1) NH3(aq) + H2O(l) --- NH4(aq) + OH-(aq)
(2) H+(aq) + OH-(aq) -- H2O(l)
H+(aq)+ NH3(aq)--- NH4+(aq)
Example
Write net-ionic equations for each of the following
reactions in dilute water solution.
(a) Hypochlorous acid (HClO) and calcium
hydroxide
(b) Ammonia with perchloric acid (HClO4)
(c) Hydroiodic acid (HI) with sodium hydroxide
Acid-base reactions in water solution are commonly
used to determine the concentration of a dissolved
species or its percentage in a solid mixture.
This is done by carrying out a titration, measuring
the volume of a standard solution (a solution of
known concentration) required to react with a
measured amount of sample.
A titration is done to identify the point at which the
reaction is complete called the equivalence point.
Acid-base indicators are used to help identify the
equivalence point.
Example
In a titration, it is found that 25.0 mL of 0.500M
NaOH is required to react with
(a) a 15.0 mL sample of HCl. What is the molarity of
the HCl?
(b) a 15.0 mL sample of a weak acid, H2A. What is
the molarity of H2A, assuming the reaction to be
H2A(aq)+ 2 OH-(aq)--- 2 H2O(l) + A-2(aq)?
(c)an aspirin table weighing 2.50 g. What is the
percentage of acetylsalicylic acid, HC9H7O4, in the
aspirin tablet? The reaction is
HC9H7O4(s) + OH-(aq) -- H2O(l)+ C9H7O4-(aq)
Oxidation-Reduction Reactions
A reaction in aqueous solution that involves the transfer
of electrons between two species is called an oxidationreduction reaction or a redox reaction.
In a redox reaction, one species loses (donates) electrons
and is oxidized. Another substance gains (receives)
electrons and is reduced.
Example
Oxidation: Zn(s) -- Zn+2(aq)+ 2eReduction: 2H+(aq) + 2e- -- H2(g)
Oxidation and reduction occur together, in the same
reaction. There are two half-reactions (an oxidation and
reduction half) in every redox reaction.
There is no change in the number of electrons in a
redox reaction. Those given off by the oxidation half are
taken by a species in the reduction half.
The two species that exchange electrons in a redox
reaction are called the oxidizing agent (accepts
electrons) and reducing agent (donates electrons).
Oxidation numbers are used to keep track of
electrons in redox reactions.
Rules
(1) The oxidation number of an element in an
elementary substance (free element) is zero.
Ex: Cl2 and P4 = 0
(2) Monatomic ions have oxidation numbers equal
to their charge.
Ex: Na+1 and Cl-1
(3) O-2 except for peroxides (O2-1) where O-1
Ex: Na2O2 (O-1) and Na2O (O-2)
(4) H+1 except for hydrides (H-1)
Ex: HCl (H+) and NaH (H-)
(5) The sum of the oxidation numbers for a neutral
species is zero. For a polyatomic group, the sum
of the oxidation numbers should equal the
charge on the ion.
Ex: SO4-2 (S+6, O-2) and CO2(C+4, O-2)
Example
What is the oxidation number of phosphorus in
sodium phosphate, Na3PO4? In the dihydrogen
phosphate ion, H2PO4-1?
Oxidation is defined as an increase in oxidation
number and reduction is defined as a decrease in
oxidation number.
Example
Zn(s) + 2H+(aq) --- Zn+2(aq) + H2(g)
Zn0 - Zn+2 (oxidized)
H+1 - H0 (reduced)
Balancing Half-Reactions (Equations)
(1) Balance the atoms of the element being oxidized
or reduced.
(2) Balance the oxidation number by adding
electrons.
(3) Balance the charge by adding H+ ions in acidic
solution and OH- ions in basic solution.
(4) Balance H by adding water.
(5) Check to make sure that the equation is
balanced with respect to mass and charge.
Example
Balance the following half-equations:
(a) MNO4-(aq) --- Mn+2(aq) (acidic solution)
(b) Cr(OH)3(s) -- CrO42-(aq) (basic solution)
To balance a full redox reaction: (1) Split the equation
into two half-equations. (2) Balance the half-equations.
(3) Combine the two half-equations to eliminate
electrons.
Example
(a) Balance the following redox reaction in acidic
solution.
Fe+2(aq) + MnO4-(aq) -- Fe+3(aq) + Mn+2(aq)
(b) Using the balanced equation what volume of 0.684
M KMnO4 solution is required to react completely with
27.50 mL of 0.250 M Fe(NO3)2?
Example
Balance the following redox reaction in basic solution.
Cl2(g) + Cr(OH)3(s) ---- Cl-(aq) + CrO4-2(aq)