chapter 6 acids, bases, salts, solubility acids, bases, salts, solubility, and stuff like that!
TRANSCRIPT
Chapter 6
Acids, Bases, Acids, Bases, Salts, SolubilitySalts, Solubility,
And stuff like that!And stuff like that!
Definitions
Solubility:Solubility:
those compounds with low solubility are said to be insolubleinsoluble,
those compounds with higher solubility are said to be solublesoluble
More Definitions
saturated solution:saturated solution:
unsaturated solution:unsaturated solution:
supersaturated solution:supersaturated solution:
Soluble or Insoluble
Explain why some substances are soluble and other substances are not soluble by giving one example of each. Used balanced equations in you discussion.
You may use the solubility rules – Use your intelligence and understanding if the internet to find them! – These are observation based – no explanation needed at this time.
Solvation
What happens when substances dissolve? What forces are involved? Use water as a solvent for specific examples.
Ionic?
Covalent?
Water as a Solvent How water dissolves ionic compounds
water is a
ions
Water as a Solvent
How water dissolves molecular compounds nonpolar covalent molecules
polar covalent molecules dissolve because
Each individual molecule is
Electrolytes
Video Link-electrolytes and non-electrolytes
Video Link – Weak and strong electrolytes
Electrolytes cations migrate to the negative electrode (the
cathode) anions migrate to the positive electrode (the anode) the movement of ions constitutes an electric current electrolyte:electrolyte:
nonelectrolytenonelectrolyte
strong electrolyte:strong electrolyte:
weak electrolyte:weak electrolyte:
Arrhenius Acids and Bases
In 1884, Svante Arrhenius proposed these definitions acid:acid: a substance that produces H3O+ ions
aqueous solution base:base: a substance that produces OH- ions in
aqueous solution
Arrhenius Acids and Bases
when HCl, for example, dissolves in water, its reacts with water to give hydronium ion and chloride ion
we use curved arrows to show the change in position of electron pairs during this reaction
HCl(aq)+H2O(l) H3O+(aq) + Cl-(aq)
H O
:
+ H Cl:
: : H O H
:
+H
+Cl -
::: ::
H
Arrhenius Acids and Bases
With bases, the situation is slightly different many bases are metal hydroxides such as KOH,
NaOH, Mg(OH)2, and Ca(OH)2
these compounds are ionic solids and when they dissolve in water, their ions merely separate
other bases are not hydroxides; these bases produce OH- by reacting with water molecules
NaOH(s) H2O Na+(aq) +OH-(aq)
NH3(aq) + H2O(l) NH4+(aq) + OH-(aq)
Arrhenius Acids and Bases
we use curved arrows to show the transfer of a proton from water to ammonia
HO H
::+ H N H
H
H+ + O
:::
-H N
H
H: H
Acid and Base Strength Strong acid:Strong acid: one that reacts completely or almost
completely with water to form H3O+ ions Strong base:Strong base: one that reacts completely or almost
completely with water to form OH- ions here are the six most common strong acids and the four
most common strong bases
HClHBrHIHNO3
H2SO4
HClO4
LiOHNaOHKOH
Ba(OH)2
Hydrochloric acidHydrobromic acidHydroiodic acidNitric acidSulfuric acidPerchloric acid
Lithium hydroxideSodium hydroxidePotassium hydroxideBarium hydroxide
Formula Name Formula Name
Acid and Base Strength Weak acid:Weak acid: a substance that dissociates only
partially in water to produce H3O+ ions
acetic acid, for example, is a weak acid; in water, only 4 out every 1000 molecules are converted to acetate ions
Weak base:Weak base: a substance that dissociates only partially in water to produce OH- ions ammonia, for example, is a weak base
CH3COOH(aq) + H2O(l) CH3COO-(aq) + H3O+(aq)
Acetic acid Acetate ion
NH3(aq) + H2O(l) NH4+(aq) + OH-(aq)
Brønsted-Lowry Acids & Bases Acid:Acid: a proton donor Base:Base: a proton acceptor Acid-base reaction:Acid-base reaction: a proton transfer reaction Conjugate acid-base pair:Conjugate acid-base pair: any pair of molecules or ions
that can be interconverted by transfer of a proton
HCl(aq) + H2O(l) H3O+(aq)+Cl-(aq)
WaterHydrogenchloride
Hydroniumion
Chlorideion
(base)(acid) (conjugateacid of water)
(conjugatebase of HCl)
conjugate acid-base pair
conjugate acid-base pair
Brønsted-Lowry Acids & Bases
Brønsted-Lowry definitions do not require water as a reactant
NH4+CH3COOH CH3COO-
NH3
(base) (conjugate baseacetic acid)
(conjugate acidof ammonia)
conjugate acid-base pair
+ +Acetic acid Ammonia
(acid)
conjugate acid-base pair
Acetate ion
Ammoniumion
Brønsted-Lowry Acids & Bases
we can use curved arrows to show the transfer of a proton from acetic acid to ammonia
CH3-C-OO
H N HH
H CH3-C-O -
OH N H
H
H+ +
Acetic acid(proton donor)
Acetate ion
+:
:: ::
:: :: :
Ammonia(proton acceptor)
Ammoniumion
C2H5OH C2H5O-H2O OH-HPO4
2- PO43-
HCO3- CO3
2-
C6H5OH C6H5O-HCN CN-
NH3NH4+
H2PO4- HPO4
2-
H2S HS-H2CO3 HCO3
-CH3COOH CH3COO-H3PO4 H2PO4
-HSO4
- SO42-
H2OH3O+HNO3 NO3
-H2SO4 HSO4
-HCl Cl-HI I-Hydroiodic acid
Hydrochloric acidSulfuric acid
Dihydrogen phosphateAcetateBicarbonate
Hydrogen phosphateAmmonia
Phenoxide
Carbonate
PhosphateHydroxideEthoxide
Hydrogen sulfide
Nitric acidHydronium ion
Hydrogen sulfate ion
Name of acid Name of ion
Phosphoric acidAcetic acidCarbonic acid
Dihydrogen phosphateAmmonium ion
Phenol
Bicarbonate ion
Hydrogen phosphate ionWaterEthanol
Hydrogen sulfide
AcidConjugate Base
IodideChlorideHydrogen sulfateNitrateWater
Sulfate
StrongAcids
Weak Acids
Weak Bases
StrongBases
Hydrocyanic acid Cyanide
Brønsted-Lowry Acids & Bases
Note the following about the conjugate acid-base pairs in the table
1. an acid can be positively charged, neutral, or negatively charged; examples of each type are H3O
+, H2CO3, and H2PO4-
2. a base can be negatively charged or neutral; examples are OH-, Cl-, and NH3
3. acids are classified a monoprotic, diprotic, or triprotic depending on the number of protons each may give up; examples are HCl, H2CO3, and H3PO4
Brønsted-Lowry Acids & Bases carbonic acid, for example can give up one proton to become
bicarbonate ion, and then the second proton to become carbonate ion
4. several molecules and ions appear in both the acid and conjugate base columns; that is, each can function as either an acid or a base
Carbonic acid
Bicarbonateion
Bicarbonateion
Carbonateion
H2CO3 H2O
HCO3- H2O
HCO3-
CO32-
H3O+
H3O+
+ +
+ +
Brønsted-Lowry Acids & Bases the HCO3
- ion, for example, can give up a proton to become CO3
2-, or it can accept a proton to become H2CO3
a substance that can act as either an acid or a base is said to be amphiproticamphiprotic
the most important amphiprotic substance in Table 8.2 is H2O; it can accept a proton to become H3O+, or lose a proton to become OH-
5. a substance cannot be a Brønsted-Lowry acid unless it contains a hydrogen atom, but not all hydrogen atoms in most compounds can be given up acetic acid, for example, gives up only one proton
Brønsted-Lowry Acids & Bases
6. there is an inverse relationship between the strength of an acid and the strength of its conjugate base the stronger the acid, the weaker its conjugate base HI, for example, is the strongest acid in Table 8.2, and its
conjugate base, I-, is the weakest base in the table CH3COOH (acetic acid) is a stronger acid that H2CO3
(carbonic acid); conversely, CH3COO- (acetate ion) is a weaker base that HCO3
- (bicarbonate ion)
6.4. DISSOCIATION OF ACIDS AND BASES IN WATER
Table 6.1. Dissociation of Acids% dissociated
Formula Name Common uses in 1 M solution Strength
H2SO4 Sulfuric Industrial chemical 100 Strong
HNO3 Nitric Industrial chemical 100 Strong
H3PO4 Phosphoric Fertilizer, food 8 Moderately additive weak
H3C6 H5O7 Citric Fruit drinks 3 WeakCH3CO2H Acetic Foods, industry 0.4 WeakHClO Hypochlorous Disinfectant 0.02 WeakHCN Hydrocyanic Very poisonous 0.002 Very weak
industrial chemical electroplating waste
H3BO3 Boric acid Antiseptic, ceramics 0.002 Very weak
Acid-Base Equilibria we know that HCl is a strong acid, which means that the
position of this equilibrium lies very far to the right
in contrast, acetic acid is a weak acid, and the position of its equilibrium lies very far to the left
but what if the base is not water? How can we determine which are the major species present?
HCl + H2O H3O++Cl-
H3O+CH3COO-H2OCH3COOH + +
Acetic acid Acetate ion
CH3COOH NH3 CH3COO-NH4
++ +
Acetic acid Acetate ion
?
Ammonia Ammonium ion(conjugate baseof CH3COOH
(conjugate acidof NH3
(acid) (base)
Acid-Base Equilibria
To predict the position of an acid-base equilibrium such as this, we do the following identify the two acids in the equilibrium; one on the left and one
on the right using the information in Table 10.1, determine which is the
stronger acid and which is the weaker acid also determine which is the stronger base and which is the
weaker base; remember that the stronger acid gives the weaker conjugate base, and the weaker acid gives the stronger conjugate base
the stronger acid reacts with the stronger base to give the weaker acid and weaker base; equilibrium lies on the side of the weaker acid and weaker base
Acid-Base Equilibria identify the two acids and bases, and their relative
strengths
the position of this equilibrium lies to the right
CH3COOH NH3 CH3COO-NH4
++ +
Acetic acid(stronger acid)
Acetate ion(weaker base)
Ammonia(stronger base)
Ammonium ion(weaker acid)
?
CH3COOH NH3 CH3COO- NH4++ +
Acetic acid(stronger acid)
Acetate ion(weaker base)
Ammonia(stronger base)
Ammonium ion(weaker acid)
Acid-Base Equilibria
Example:Example: predict the position of equilibrium in this acid-base reaction
H2CO3 OH- HCO3- H2O+ +
?
Acid-Base Equilibria
Example:Example: predict the position of equilibrium in this acid-base reaction
Solution:Solution: the position of this equilibrium lies to the right
H2CO3 OH- HCO3- H2O+ +
?
H2CO3 OH- HCO3- H2O+ +
Strongeracid
Strongerbase
Weakerbase
Weakeracid
Acid Ionization Constants
when a weak acid, HA, dissolves in water
the equilibrium constant, Keq, for this ionization is
because water is the solvent and its concentration changes very little when we add HA to it, we treat [H2O] as a constant equal to 1000 g/L or 55.5 mol/L
we combine the two constants to give a new constant, which we call an acid ionization constant, Ka
HA H2O A- H3O++ +
[HA][H2O]
[A-][H3O+]Keq =
[HA]
[A-][H3O+]Ka = Keq[H2O] =
Acid Ionization Constants Ka for acetic acid, for example is 1.8 x 10-5
because the acid ionization constants for weak acids are numbers with negative exponents, we commonly express acid strengths as pKa where
the value of pKa for acetic acid is 4.75
values of Ka and pKa for some weak acids are given in Table 10.2
as you study the entries in this table, note the inverse relationship between values of Ka and pKa
the weaker the acid, the smaller its Ka, but the larger its pKa
pKa = -log Ka
H3PO4
HCOOH
CH3CH(OH)COOH
CH3COOH
H2CO3
H2PO4-
H3BO3
NH4+
C6H5OH
HPO42-
HCO3-
HCN
Phosphoric acid
Formic acid
Lactic acid
Acetic acid
Carbonic acid
Dihydrogen phosphate ion
Name
7.21
pKa
9.14
9.25
9.89
12.66
10.25
Boric acid
Ammonium ion
Phenol
Hydrogen phosphate ion
Bicarbonate ion
Acid
7.5 x 10-3
1.8 x 10-4
8.4 x 10-4
1.8 x 10-5
4.3 x 10-7
6.2 x 10-8
Ka
7.3 x 10-10
5.6 x 10-10
1.3 x 10-10
2.2 x 10-13
5.6 x 10-11
2.12
3.75
3.08
4.75
6.37
Hydrocyanic acid 4.9 x 10-10 9.31
Properties of Acids & Bases
Neutralization acids and bases react with each other in a process
called neutralization. Reaction of acids with metals
strong acids react with certain metals (called active metals) to produce a salt and hydrogen gas, H2
Mg(s) + 2HCl(aq) MgCl2(aq) + H2(g)Magnesium Hydrochloric
acidMagnesium
chlorideHydrogen
Properties of Acids & Bases
Reaction with metal hydroxides reaction of an acid with a metal hydroxide gives a salt plus water
the reaction is more accurately written as
omitting spectator ions gives this net ionic equation
HCl(aq)Hydrochloric
acid
+ KOH(aq)
Water
+KCl(aq)Potassiumchloride
Potassiumhydroxide
H2O
H3O+ Cl- K+ OH- 2H2O Cl- K++ + + + +
H3O+ OH- 2H2O+
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 and pOH pH of some common materials
pH
Battery acidGastric juiceLemon juiceVinegarTomato juiceCarbonated beveragesBlack coffee
UrineRain (unpolluted)
Milk
SalivaPure waterBloodBilePancreatic fluidSeawaterSoap
Milk of magnesiaHousehold ammonia
Lye (1.0 M NaOH)
0.51.0-3.02.2-2.42.4-3.44.0-4.44.0-5.05.0-5.1
5.5-7.56.2
6.3-6.6
6.5-7.57.0
7.35-7.456.8-7.07.8-8.08.0-9.08.0-10.0
10.511.7
14.0
Material pHMaterial
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
Blood Buffers
The average pH of human blood is 7.4 any change larger than 0.10 pH unit in either direction
can cause illness To maintain this pH, the body uses three buffer
systems carbonate buffer:carbonate buffer: H2CO3 and its conjugate base,
HCO3-
phosphate buffer:phosphate buffer: H2PO4- and its conjugate base,
HPO42-
proteins:proteins: discussed in Chapter 21
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
PREPARATION OF ACIDSCombination of H with nonmetal: H2 + Cl2 2HCl
Nonmetal with water: Cl2 + H2O HCl + HClO
Nonmetal oxide plus water: SO3 + H2O H2SO4
Evolution of volatile acid: 2NaCl(s) + H2SO4(l) 2HCl(g) + Na2SO4(s)
• HCl gas collected in water gives hydrochloric acidOrganic acids, such as acetic acid, have the carboxylic acid group:
Carboxylic acidgroup, -CO2 HC
OOH
PREPARATION OF BASESActive metal plus water
• 2K + H2O 2K+ + 2OH- + H2(g)
Metal oxide plus water• CaO + H2O Ca(OH)2
Substances that generate OH- in water
NH3 + H2O NH4+ + OH-
Salt anions that react with water to produce OH-
• From Na2CO3: 2Na+ + CO32- + H2O 2Na+ + HCO3
-+ OH-
• This reaction is a hydrolysis reaction
Organic bases, particularly amines
• (CH3)3N + H2O (CH3)3NH+ + OH-
PREPARATION OF SALTSReaction of acid with base• 2NaOH + H2SO4 2H2O + Na2SO4 (sodium sulfate)
Reaction of metal and nonmetal• Ca + F2 CaF2 (calcium fluoride)
Metal reacting with acid• Mg + H2SO4 H2 + MgSO4 (magnesium sulfate)
Active metal reacting with base• 2Al + 6NaOH 3H2(g) + Na3AlO3 (sodium aluminate)
Addition of a base to a salt to form another salt and an insoluble base• 2KOH + MgSO4 Mg(OH)2(s) + K2SO4(aq)
Evolution of a volatile acid leaving a salt• 2NaCl(s) + H2SO4(l) 2HCl(g) + Na2SO4(s)
Displacement of a metal from a salt, such as in cementation• Fe(s) + CdSO4(aq) Cd(s) + FeSO4(aq)
Specialized processes, such as the Solvay synthesis of NaHCO3
NaCl + NH3 + CO2 + H2O NaHCO3 + NH4Cl
6.10. ACID SALTS AND BASIC SALTSAcid salts are salts that contain H and can act as acids• NaHSO4 + NaOH Na2SO4 + H2O
• Sodium bicarbonate: NaHCO3
• Sodium dihydrogen phosphate, NaH2PO4, used to prepare buffers
• Disodium hydrogen phosphate, Na2HPO4, buffers
• Potassium hydrogen tartrate, KH4C4H4O6, acid in baking powder
Basic salts contain OH and can react with H+ ion
Example: Calcium hydroxyapatite source of phosphorus• Ca5OH(PO4)3
Many rock-forming minerals are basic salts
6.11. WATER OF HYDRATIONWater molecules bound to other compounds, typically salts• Example: Sodium carbonate decahydrate, Na2CO3•10H2O
6.12. NAMES OF ACIDS, BASES, AND SALTSAcidsH and a nonmetal: Hydro-ic acid• Hydrochloric acid, HClOxygen-containing acids• H2SO3, sulfurous acid
Table 6.7. Names of Oxyacids of ChlorineFormula Name Anion nameHClO4 perchloric acid perchlorate
HClO3 chloric acid chlorate
HClO2 chlorous acid chlorite
HClO hypochlorous acid hypochlorite
Bases For ionic bases containing OH, name of cation followed by hydroxide• NaOH, sodium hydroxide Ca(OH)2, calcium hydroxide
Names of SaltsName of cation followed by name of anionSee Table 6.8 for some important ions and their namesExamples:• Na2SO4, sodium sulfate
• KH2PO4, potassium dihydrogen sulfate
• Ca(ClO)2, calcium hypochlorite
Formulas of SaltsSum of charge on cations times their subscripts plus sum of charge on anions times their subscripts must equal zero• Example: Iron(III) sulfate• Formula before adding subscripts: Fe(SO4)
• Cation charge: +3 • Anion charge: -2• 2 Fe3+ cations gives a total cation charge of 2 x 3 = 6• 3 SO4
2- gives an anion charge of 3 x (-2) = -6
• Therefore, the formula is Fe2(SO4)3