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Acid/Base DefinitionsAcid/Base Definitions Types of Acids/basesTypes of Acids/bases
Polyprotic AcidsPolyprotic Acids The Ion Product for WaterThe Ion Product for Water
The pH and Other “p” ScalesThe pH and Other “p” Scales Aqueous Solutions of Acids and BasesAqueous Solutions of Acids and Bases
HydrolysisHydrolysis The Common Ion EffectThe Common Ion Effect
Buffer SolutionsBuffer Solutions Indicators and TitrationsIndicators and Titrations
Chapter 15. Acids & Bases
Types of Reactions a) Precipitation Reactions. Ionic compounds or salts b) Acid/base Reactions. Acids and Bases c) Redox Reactions. Oxidizing & Reducing agents
What are Acids &Bases?
Definition?
a) Arrhenius
b) Bronsted-Lowry
c) Lewis
Arrhenius definitions AcidAcid Anything that produces hydrogen
ions in a water solution.» HCl (aq) H+ + Cl-
BaseBase Anything that produces hydroxide ions in a water solution.
» NaOH (aq) Na+ + OH-
Arrhenius definitions are limited to aqueous solutions.
Acid base reactions: HCl(aq) + NaOH(aq) NaCl(aq) + H2O(l)
Brønsted-Lowry definitions
Expands the Arrhenius definitions
AcidAcid Proton donor
BaseBase Proton acceptor This definition explains how substances
like ammonia can act as bases.
Eg. HCl(g) + NH3(g) ------> NH4Cl(s)
HCl (acid), NH3 (base).
NH3(g) + H2O(l) NH4+ + OH-
Lewis Definition Lewis was successful in including acid and
bases without proton or hydroxyl ions.
Lewis Acid: A substance that accepts an
electron pair.
Lewis base: A substance that donates an
electron pair.
E.g. BF3(g) + :NH3(g) F3B:NH3(s)
Types of Acids and Bases• Binary acids • Oxyacid • Organic acids • Acidic oxides• Basic oxides • Amine• Polyprotic acids
Binary Acids Compounds containing acidic protons
bonded to a more electronegative atom.
e.g. HF, HCl, HBr, HI, H2S The acidity of the haloacid (HX; X = Cl, Br, I, F) Series increase in the following order: HF < HCl < HBr < HI
Oxyacids Compounds containing acidic - OH groups
in the molecule. Acidity of H2SO4 is greater than H2SO3
because of the extra O (oxygens) The order of acidity of oxyacids from the a
halogen (Cl, Br, or I) shows a similar trend.
HClO4 > HClO3 > HClO2 >HClO perchloric chloric chlorus hyphochlorus
Acidic Oxides
These are usually oxides of non-metallic elements such as P, S and N.
E.g. NO2, SO2, SO3, CO2 They produce oxyacids when
dissolved in water
Basic Oxides
Oxides oxides of metallic elements such as Na, K, Ca. They produce hydroxyl bases when dissolved in water.
e.g. CaO + H2O --> Ca(OH)2
Protic Acids
Monoprotic Acids: The form protic refers to acidity or protons. Monoprotic acids have only one acidic proton. e.g. HCl.
Polyprotic Acids: They have more than one acidic proton.
e.g. H2SO4 - diprotic acid
H3PO4 - triprotic acid.
Amines
Class of organic bases derived from ammonia NH3 by replacing hydrogen by organic groups. They are defined as bases similar to NH3 by Bronsted or Lewis acid/base definitions.
What acid base concepts (Arrhenius/Bronsted/Lewis) would best describe the following reactions:
a) HCl(aq) + NaOH(aq) ---> NaCl(aq) + H2O(l)
b)HCl(g) + NH3(g) ---> NH4Cl(s)
c)BF3(g) + NH3(g) ---> F3B:NH3(s)
d)Zn(OH)2(s) + 2OH-(aq) ---> [Zn(OH)4]2- (aq)
Common acids and bases AcidsAcids Formula Molarity* nitric HNO3 16
hydrochloric HCl 12 sulfuric H2SO4 18
acetic HC2H3O2 18
BasesBases ammonia NH3(aq) 15 sodium hydroxide NaOH solid
*undiluted.
Acids and bases
AcidicAcidic BasicBasic
– Citrus fruits Baking soda– Aspirin Detergents– Coca Cola Ammonia
cleaners– Vinegar Tums and Rolaids– Vitamin C Soap
Dissociation Equilibrium,
HCl(aq) + H2O(l) H3+O(aq) + Cl-(aq)
H2SO4(aq) + H2O(l) H3+O(aq) + HSO4
-(aq)
H2O(l) + H2O(l) H3+O(aq) + OH-(aq)
This dissociation is called autoionization of water.
HC2H3O2(aq) + H2O(l) H3+O(aq) + C2H3O2
-(aq)
NH3 (aq) + H2O(l) NH4+ + OH-(aq)
Brønsted-Lowry definitions Conjugate acid-base pairsConjugate acid-base pairs.. Acids and bases that are related by
loss or gain of H+ as H3O+ and H2O. Examples.Examples. Acid Base
H3O+ H2O
HC2H3O2 C2H3O2-
NH4+ NH3
H2SO4 HSO4-
HSO4- SO42-
Bronsted acid/conjugate base and base/conjugate acid pairs inacid/base equilibria
HCl(aq) + H2O(l) H3+O(aq) + Cl-(aq)
HCl(aq): acid H2O(l): base
H3+O(aq): conjugate acid
Cl-(aq): conjugate base H2O/ H3
+O: base/conjugate acid pair HCl/Cl-: acid/conjugate base pair
Select acid, base, acid/conjugate base pair,base/conjugate acid pair
H2SO4(aq) + H2O(l) H 3+O(aq) + HSO4
-(aq) acid base conjugate acid conjugate base base/conjugate acid pair acid/conjugate base pair
Equilibrium, Constant, Ka & Kb
Ka: Acid dissociation constant for a equilibrium reaction.
Kb: Base dissociation constant for a equilibrium reaction.
Acid: HA + H2O H3+O + A-
Base: BOH + H2O B+ + OH-
[H3+O][ A-] [B+ ][OH-]
Ka = --------------- ; Kb = ----------------- [HA] [BOH]
What is Ka
HCl(aq) + H2O(l) <===> H3+O(aq) + Cl-(aq)
E.g. Ka
HCl(aq) + H2O(l) H3+O(aq) + Cl-(aq)
[H3
+O][Cl-] Ka= ----------------- [HCl]
[H+][Cl-] Ka= ----------------- [HCl]
What is Ka1 and Ka2?
H2SO4(aq) + H2O(l) H3+O(aq) + HSO4
-(aq)
HSO4-(aq) + H2O(l) H3
+O(aq) + SO42-(aq)
What is Kb
NH3 (aq) + H2O(l) NH4+ + OH-(aq)
E.g. H2SO4(aq) + H2O(l) H3
+O(aq) + HSO4-(aq)
HSO4-(aq) + H2O(l) H3
+O(aq) + SO42-(aq)
[H3+O][HSO4
-]
H2SO4 ; Ka1 = -------------------
[H2SO4]
[H3+O][SO4
2-]
H2SO4 ; Ka2 = -------------------
[HSO4-]
E.g.
HC2H3O2(aq) + H2O(l) H3+O(aq) + C2H3O2
-(aq)
[H+][C2H3O2-]
H C2H3O2; Ka= ------------------
[H C2H3O2]
NH3 (aq) + H2O(l) NH4+ + OH-(aq)
[NH4+][OH-]
NH3; Kb= --------------
[ NH3]
Which is weaker?
a. HNO2 ; Ka= 4.0 x 10-4.
b. HOCl2 ; Ka= 1.2 x 10-2.
c. HOCl ; Ka= 3.5 x 10-8.
d. HCN ; Ka= 4.9 x 10-10.
WEAKER/STRONGER Acids and Bases & Ka and Kb values
A larger value of Ka or Kb indicates an equilibrium favoring product side.
Acidity and basicity increase with increasing Ka or Kb.
pKa = - log Ka and pKb = - log Kb
Acidity and basicity decrease with increasing pKa or pKb.
Autoionization of water AutoionizationAutoionization When water molecules react with one
another to form ions.
H2O(l) + H2O(l) H3O+(aq) + OH-
(aq)
– (10-7M) (10-7M)
Kw = [ H3O+ ] [ OH- ]
= 1.0 x 10-14 at 25oC
Note:Note: [H2O] is constant and is
included in Kw.
ion productof water
ion productof water
What is Kw?
H2O(l) + H2O(l) H3+O(aq) + OH-(aq)
This dissociation is called autoionization of water.
Autoionization of water: Kw = [H3
+O][OH-]
Kw is called ionic product of water
Kw = 1 x 10-14
Why is water important for acid/base equilibria? Water is the medium/solvent for acids
and bases. Acids and bases alter the dissociation
equilibrium of water based on Le Chaterlier’s principle
H2O(l) + H2O(l) H3+O(aq) + OH-(aq)
Comparing Kw and Ka & Kb
Any compound with a Ka value greater than Kw of water will be a an acid in water.
Any compound with a Kb value greater than Kw of water will be a base in water.
pH and other “p” scales
We need to measure and use acids and bases over a very large concentration range.
pH and pOH are systems to keep track of these very large ranges.– pH = -log[H3O+]
– pOH = -log[OH-]– pH + pOH = 14
pH scale
A logarithmic scale used to keep track of the large changes in [H+].
14 7 0
10-14 M 10-7 M 1 M Very Neutral VeryBasic Acidic
When you add an acid, the pH gets smaller.
When you add a base, the pH gets larger.
pH of somecommon materials
Substance pH
1 M HCl 0.0Gastric juices 1.0 - 3.0Lemon juice 2.2 - 2.4Classic Coke 2.5Coffee 5.0Pure Water 7.0Blood 7.35 - 7.45Milk of Magnesia 10.5Household ammonia 12.0
1M NaOH 14.0
Substance pH
1 M HCl 0.0Gastric juices 1.0 - 3.0Lemon juice 2.2 - 2.4Classic Coke 2.5Coffee 5.0Pure Water 7.0Blood 7.35 - 7.45Milk of Magnesia 10.5Household ammonia 12.0
1M NaOH 14.0
What is pH? Kw = [H3
+O][OH-] = 1 x 10-14
[H3+O][OH-] = 10-7 x 10-7
Extreme cases: Basic medium [H3
+O][OH-] = 10-14 x 100 Acidic medium [H3
+O][OH-] = 100 x 10-14 pH value is -log[H+] spans only 0-14 in water.
pH, pKw and pOH The relation of pH, Kw and pOH
Kw = [H+][OH-]
log Kw = log [H+] + log [OH-]
-log Kw= -log [H+] -log [OH-] ; previous equation multiplied by -1 pKw = pH + pOH; pKw = 14
since Kw =1 x 10-14 14 = pH + pOH pH = 14 - pOH pOH = 14 - pH
Acid and Base Strength Strong acidsStrong acids Ionize completely in water.
HCl, HBr, HI, HClO3, HNO3, HClO4, H2SO4.
Weak acids Weak acids Partially ionize in water. Most acids are weak.
Strong basesStrong bases Ionize completely in water. Strong bases are metal
hydroxides - NaOH, KOH
Weak basesWeak bases Partially ionize in water.
pH and pOH calculations of acid and base solutions a) Strong acids/bases
dissociation is complete for strong acid such as HNO3 or base NaOH
– [H+] is calculated from molarity (M) of the solution
b) weak acids/bases
needs Ka , Kb or percent(%)dissociation
pH of Strong Acid/bases
HNO3(aq) + H2O(l) H3+O(aq) + NO3-(aq)
Therefore, the moles of H+ ions in the solution is equal to moles of HNO3 at the beginning.
[HNO3] = [H+] = 0.2 mole/L pH = -log [H+] = -log(0.2) pH = 0.699
pH of 0.5 M H2SO4 Solution
H2SO4(aq) + H2O(l) H3+O(aq) + HSO4
-(aq)
HSO4-(aq) + H2O(l) H3
+O(aq) + SO42-(aq)
[H3+O][HSO4
-]
H2SO4 ; Ka1 = -------------------
[H2SO4]
[H3+O][SO4
2-]
H2SO4 ; Ka2 = ------------------- ; Ka2 ignored
[HSO4-]
pH of 0.5 M H2SO4 Solution
H2SO4(aq) + H2O(l) H3+O(aq) + HSO4
-(aq)
the moles of H+ ions in the solution is equal to moles of H2SO4 at the beginning.
[H2SO4] = [H+] = 0.5 mole/L
pH = -log [H+]
pH = -log(0.5)
pH = 0.30
1.5 x 10-2 M NaOH.1.5 x 10-2 M NaOH.
NaOH is also a strong base dissociates completely in water.
[NaOH] = [HO- ] = 1.5 x 10-2 mole/L
pOH = -log[HO-]= -log(1.5 x 10-2)
pOH = 1.82
As defined and derived previously: pKw= pH + pOH; pKw= 14
pH = pKw + pOH
pH = 14 - pOH
pH = 14 - 1.82 ; pH = 12.18