aqueous stuff aqueous stuff. reactions between ions ionic compounds, also called salts, consist of...
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Aqueous StuffAqueous Stuff
Reactions Between Ions Ionic compounds, also called salts, consist of both
positive and negative ions When an ionic compound dissolves in water, it
dissociates to aqueous ions
What happens when we mix aqueous solutions of two different ionic compounds? if two of the ions combine to form a water-insoluble
compound, a precipitate will form otherwise no physical change will be observed
NaCl(s) +Na+(aq)H2O
Cl-(aq)
Reactions Between Ions Example:
suppose we prepare these two aqueous solutions
if we then mix the two solutions, we have four ions present; of these, Ag+ and Cl- react to form AgCl(s) which precipitates
+Ag+(aq) Cl-(aq)
AgCl(s)
NO3-(aq) + Na+(aq) +
+ Na+(aq) + NO3-(aq)
AgNO3(s)H2O
Ag+(aq) + NO3-(aq)Solution 1
NaCl(s)H2O
Na+(aq) + Cl-(aq)Solution 2
Reactions Between Ions we can simplify the equation for the formation
of AgCl by omitting all ions that do not participate in the reaction
the simplified equation is called a net ionic net ionic equationequation; it shows only the ions that react
ions that do not participate in a reaction are called spectator ionsspectator ions
Ag+(aq) Cl-(aq) AgCl(s)+Net ionic equation:
Reactions Between Ions In general, ions in solution react with each other
when one of the following can happen two of them form a compound that is insoluble in water two of them react to form a gas that escapes from the
reaction mixture as bubbles, as for example when we mix aqueous solutions of sodium bicarbonate and hydrochloric acid
an acid neutralizes a base one of the materials can oxidize another
HCO3-(aq) + H3O+(aq) +CO2(g) 2H2O(l)
Bicarbonate ion Carbon dioxide
Reactions Between Ions Following are some generalizations about which ionic
solids are soluble in water and which are insoluble all compounds containing Na+, K+, and NH4
+ are soluble in water
all nitrates (NO3-) and acetates (CH3COO-) are soluble in
water most chlorides (Cl-) and sulfates (SO4
2-) are soluble; exceptions are AgCl, BaSO4, and PbSO4
most carbonates (CO32-), phosphates (PO4
3-), sulfides (S2-), and hydroxides (OH-) are insoluble in water; exceptions are LiOH, NaOH, KOH, and NH4OH which are soluble in water
Oxidation-Reduction
Oxidation:Oxidation: the loss of electrons Reduction:Reduction: the gain of electrons Oxidation-reduction (redox) reaction:Oxidation-reduction (redox) reaction: any
reaction in which electrons are transferred from one species to another
Oxidation-Reduction Example: if we put a piece of zinc metal in a beaker containing a
solution of copper(II) sulfate some of the zinc metal dissolves some of the copper ions deposit on the zinc metal the blue color of Cu2+ ions gradually disappears
In this oxidation-reduction reaction zinc metal loses electrons to copper ions
copper ions gain electrons from the zincZn(s) +Zn2+(aq) 2e- Zn is oxidized
Cu(s)+ 2e-Cu2+(aq) Cu2+ is reduced
Oxidation-Reduction
we summarize these oxidation-reduction relationships in this way
electrons flowfrom Zn to Cu2+
+Zn(s) Cu2+(aq) + Cu(s)Zn2+(aq)
loses electrons;is oxidized
gains electrons;is reduced
gives electronsto Cu2+; is thereducing agent
takes electronsfrom Zn; is the oxidizing agent
Oxidation-Reduction Although the definitions of oxidation (loss of
electrons) and reduction (gain of electrons) are easy to apply to many redox reactions, they are not easy to apply to others for example, the combustion of methane
An alternative definition of oxidation-reduction is oxidation:oxidation: the gain of oxygen or loss of hydrogen reduction:reduction: the loss of oxygen or gain of hydrogen
CH4(g) + O2(g) CO2(g) + H2O(g)Methane
Oxidation-Reduction using these alternative definitions for the
combustion of methane
CH4(g) + O2(g) CO2(g) + H2O(g)
gains O and losesH; is oxidized
gains H; is reduced
is the reducingagent
is the oxidizingagent
electrons are transferred from carbon to oxygen
Oxidation-Reduction Five important types of redox reactions
combustion:combustion: burning in air. The products of complete combustion of carbon compounds are CO2 and H2O.
respiration:respiration: the process by which living organisms use O2 to oxidize carbon-containing compounds to produce CO2 and H2O. The importance of these reaction is not the CO2 produced, but the energy released.
rusting:rusting: the oxidation of iron to a mixture of iron oxides
bleaching:bleaching: the oxidation of colored compounds to products which are colorless
batteries:batteries: in most cases, the reaction taking place in a battery is a redox-reaction
4Fe(s) +3O2(g) 2Fe2O3(s)
Heat of Reaction In almost all chemical reactions, heat is either given
off or absorbed example: the combustion (oxidation) of carbon liberates
94.0 kcal per mole of carbon oxidized
Heat of reaction:Heat of reaction: the heat given off or absorbed in a chemical reaction exothermic reaction:exothermic reaction: one that gives off heat endothermic reaction:endothermic reaction: one that absorbs heat heat of combustion:heat of combustion: the heat given off in a combustion
reaction; all combustion reactions are exothermic
C(s) + O2(g) CO2(g) + 94.0 kcal/mole C
Properties of Acids & Bases
Reaction with metal oxides strong acids react with metal oxides to give water
plus a salt
2H3O+(aq) + CaO(s) 3H2O(l) + Ca2+(aq)Calciumoxide
Properties of Acids & Bases Reaction with carbonates and bicarbonates
strong acids react with carbonates to give carbonic acid, which rapidly decomposes to CO2 and H2O
strong acids react similarly with bicarbonates
2H3O+(aq) + CO32-(aq) H2CO3(aq) + 2H2O(l)
H2CO3(aq) CO2(g) + H2O(l)
2H3O+(aq) + CO32-(aq) CO2(g) + 3H2O(l)
H3O+(aq) + HCO3-(aq) H2CO3(aq) + H2O(l)
H2CO3(aq) CO2(g) + H2O(l)
H3O+(aq) + HCO3-(aq) CO2(g) + 2H2O(l)
Properties of Acids & Bases Reaction with ammonia and amines
any acid stronger than NH4+ is strong enough to react
with NH3 to give a salt
HCl(aq) + NH3(aq) NH4+(aq) + Cl-(aq)
Self-Ionization of Water
pure water contains a very small number of H3O+ ions and OH- ions formed by proton transfer from one water molecule to another
the equilibrium expression for this reaction is
we can treat [H2O] as a constant = 55.5 mol/L
H2O+H2O H3O++OH-
BaseAcid Conjugateacid of H2O
Conjugatebase of H2O
[H2O]2
[H3O+][HO-]Keq =
Self-Ionization of Water combining these constants gives a new constant called the ion ion
product of water, Kproduct of water, Kww
in pure water, the value of Kw is 1.0 x 10-14
this means that in pure water
[H3O+][OH-]Kw = Keq[H2O]2 =
Kw = 1.0 x 10-14
[H3O+]
[OH-]
= 1.0 x 10-7 mol/L
= 1.0 x 10-7 mol/Lin pure water
Self-Ionization of Water
the product of [H3O+] and [OH-] in any aqueous solution is equal to 1.0 x 10-14 for solutions as well.
for example, if we add 0.010 mole of HCl to 1 liter of pure water, it reacts completely with water to give 0.010 mole of H3O+
in this solution, [H3O+] is 0.010 or 1.0 x 10-2
this means that the concentration of hydroxide ion is
[OH-] = 1.0 x 10-14
1.0 x 10-2= 1.0 x 10-12
pH and pOH
we commonly express these concentrations as pH, where
pH = -log [H3O+] we can now state the definitions of acidic and basic
solutions in terms of pH acidic solution:acidic solution: one whose pH is less than 7.0 basic solution:basic solution: one whose pH is greater than 7.0 neutral solution:neutral solution: one whose pH is equal to 7.0
pH and pOH
just as pH is a convenient way to designate the concentration of H3O+, pOH is a convenient way to designate the concentration of OH-
pOH = -log[OH-] the ion product of water, Kw, is 1.0 x 10-14
taking the logarithm of this equation gives
pH + pOH = 14 thus, if we know the pH of an aqueous solution, we can
easily calculate its pOH
Kw = [H3O+][OH-] = 1.0 x 10-14
pH of Salt Solutions
When some salts dissolve in pure water, there is no change in pH from that of pure water
Many salts, however, are acidic or basic and cause a change the pH when they dissolve
We are concerned in this section with basic salts and acidic salts
pH of Salt Solutions
Basic salt: Basic salt: raises the pH as an example of a basic salt is sodium acetate when this salt dissolves in water, it ionizes; Na+ ions
do not react with water, but CH3COO- ions do
the position of equilibrium lies to the left nevertheless, there are enough OH- ions present in
0.10 M sodium acetate to raise the pH to 8.88
OH-CH3COOHH2OCH3COO- + +Acetic acid
(stronger acid)Acetate ion
(weaker baseHydroxide ion(stronger base)
Water(weaker base)
pH of Salt Solutions
Acidic salt:Acidic salt: lowers the pH an example of an acidic salt is ammonium chloride chloride ion does not react with water, but the
ammonium ion does
although the position of this equilibrium lies to the left, there are enough H3O+ ions present to make the solution acidic
NH4+ + H2O NH3 + H3O+
Ammonia(stronger base)
Ammonium ion(weaker acid)
Hydronium ion(stronger acid
Water(weaker base)
Acid-Base Titrations
Titration:Titration: an analytical procedure in which a solute in a solution of known concentration reacts with a known stoichiometry with a substance whose concentration is to be determined
Acid-Base Titrations
An acid-base titration must meet these requirement1. we must know the equation for the reaction so that we
can determine the stoichiometric ratio of reactants to use in our calculations
2. the reaction must be rapid and complete
3. there must be a clear-cut change in a measurable property at the end pointend point (when the reagents have combined exactly)
4. we must have precise measurements of the amount of each reactant
Acid-Base Titrations
As an example, let us use 0.108 M H2SO4 to determine the concentration of a NaOH solution requirement 1:requirement 1: we know the balanced equation
requirement 2:requirement 2: the reaction between H3O+ and OH- is rapid and complete
requirement 3:requirement 3: we can use either an acid-base indicator or a pH meter to observe the sudden change in pH that occurs at the end point of the titration
requirement 4:requirement 4: we use volumetric glassware
2NaOH(aq)+H2SO4(aq) Na2SO4(aq) + 2H2O(l)(concentration
known)(concentrationnot known)
Acid-Base Titrations experimental measurements
doing the calculations
Trial I
Volume of 0.108 M H2SO4
Volumeof NaOH
25.0 mL 33.48 mLTrial II 25.0 mL 33.46 mL
Trial III 25.0 mL 33.50 mL
average = 33.48 mL
2 mol NaOH1 mol H2SO4
= 0.161 mol NaOHL NaOH
= 0.161 M
mol NaOHL NaOH = 0.108 mol H2SO4
1 L H2SO4x x0.0250 L H2SO4
0.03348 L NaOH
pH Buffers
pH buffer:pH buffer: a solution that resists change in pH when limited amounts of acid or base are added to it a pH buffer as an acid or base “shock absorber” a pH buffer is common called simply a buffer the most common buffers consist of approximately equal
molar amounts of a weak acid and a salt of the conjugate base of the weak acid
for example, if we dissolve 1.0 mole of acetic acid and 1.0 mole of its conjugate base (in the form of sodium acetate) in water, we have an acetate buffer
pH Buffers
How an acetate buffer resists changes in pH if we add a strong acid, such as HCl, added H3O+ ions
react with acetate ions and are removed from solution
if we add a strong base, such as NaOH, added OH- ions react with acetic acid and are removed from solution
CH3COO- H3O+ CH3COOH H2O+ +
CH3COOH OH- CH3COO- H2O+ +
CH3COOH H2O CH3COO- H3O++ +
Added asCH3COOH
Added asCH3COO-Na+
pH Buffers
The effect of a buffer can be quite dramatic consider a phosphate buffer prepared by
dissolving 0.10 mole of NaH2PO4 (a weak acid) and 0.10 mole of Na2HPO4 (the salt of its conjugate base) in enough water to make 1 liter of solution
waterpH
0.10 M phosphate buffer7.07.21
2.0 12.07.12 7.30
pH afteraddition of
0.010 mole HCl
pH afteraddition of
0.010 mole NaOH
pH Buffers
Buffer pHBuffer pH if we mix equal molar amounts of a weak acid and
a salt of its conjugate base, the pH of the solution will be equal to the pKa of the weak acid
if we want a buffer of pH 9.14, for example, we can mix equal molar amounts of boric acid (H3BO3), pKa 9.14, and sodium dihydrogen borate (NaH2BO3), the salt of its conjugate base
pH Buffers
Buffer capacity depends both its pH and its concentration
pH The closer the pH of the buffer is to the pKaof the weak acid, the greater the buffer capacity
Concentration The greater the concentration of the weak acid and its conjugate base, the greater the buffer capacity
Henderson-Hasselbalch Eg. Henderson-Hasselbalch equation:Henderson-Hasselbalch equation: a mathematical
relationship between pH, pKa of the weak acid, HA concentrations HA, and its conjugate base, A-
It is derived in the following way
taking the logarithm of this equation gives
HA H2O A- H3O++ +
[HA]
[A-][H3O+]Ka =
[HA]log [H3O+] + log
[A-]log Ka =
Henderson-Hasselbalch Eg. multiplying through by -1 gives
-log Ka is by definition pKa, and -log [H3O+] is by definition pH; making these substitutions gives
rearranging terms gives
[HA]-log [H3O+] - log
[A-]-log Ka =
[HA]
[A-]+ logpH = pKa Henderson-Hasselbalch Equation
[HA][A-]
pKa = pH - log
Henderson-Hasselbalch Eg.
Example:Example: what is the pH of a phosphate buffer solution containing 1.0 mole of NaH2PO4 and 0.50 mole of Na2HPO4 dissolved in enough water to make 1.0 liter of solution
Henderson-Hasselbalch Eg. Example:Example: what is the pH of a phosphate buffer
solution containing 1.0 mole of NaH2PO4 and 0.50 mole of Na2HPO4 in enough water to make one liter of solution
SolutionSolution the equilibrium we are dealing with and its pKa are
substituting these values in the H-H equation gives
H2PO4- H2O HPO4
2- H3O++ + pKa = 7.21
1.0 mol/L 0.50 mol/L
= 7.21 - 0.30 = 6.91
+ logpH = 7.21 0.501.0
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