Acids, Bases, and Salts:
Arrhenius Definition:
Acid
A substance, when dissolved in water, produces hydrogen ions (H +)
Base
A substance, when dissolved in water, produces hydroxide ions (OH - )
Types of solutions:
1. Acids dissolve in water to produce acid / acidic solutions
2. Bases dissolve in water to produce basic or “ Alkaline solutions”
Acid and Base Dissociation reactions:
Dissociation is the process by which a molecule “breaks up” or splits into ions
Dissociation reactions describe how an acid or base dissolves in water.
Acid Dissociation reactions
Hydrochloric acid
HCl → H + + Cl –
Nitric acid
HNO3 → H + + NO3 –
Sulfuric acid
H2SO4 → H + + HSO4 –
Base Dissociation reactions
Sodium hydroxide
NaOH → Na + + OH −
Potassium hydroxide
KOH → K + + OH ‾
Barium hydroxide
Ba(OH)2 → Ba+2 + 2OH ‾
Acid Properties:
· Sour Taste
· React with Metals ( Mg, Zn Fe, Al)
· Form Electrolytic Solutions ( conduct electricity)
· Turn Litmus Paper Red
· Corrosive (chemically burn through materials)
Base Properties:
· Bitter Taste
· Feel “Soapy” or slippery
· Don’t react with metals
· Form Electrolytic Solutions
· Turn Litmus Paper Blue
· Caustic
Bronsted – Lowry Definition
Acid donates H+ ions ( proton donor)
Base accepts H+ ions ( proton acceptor)
Hydrogen ion H + = proton
Acid = Proton Donor
Base = Proton Acceptor
Water can be protonated by H+ ion
Water Acts like a Base
H+ + H2O → H3O +
H3O + = “ Hydronium Ion”
HCl + H2O → H3O+ + Cl –
HCl = Acid ( donates H+ )
H2O = Base ( accepts H+ )
Why is a water solution of ammonia basic?
NH3 (ammonia) acts like a base
NH3(g) + H2O(l) → NH4+(aq) + OH – (aq)
NH3 = Base ( H+ acceptor)
H2O = Acid ( H+ donor)
Water can act as an acid or a base
Water is “Amphoteric”
Common acids used in everyday life:
Acid Formula Use
Acetic acid HC2H3O2 Vinegar
Carbonic acid H2CO3 Carbonated
Soft Drinks, Soda
Hydrochloric HCl Stomach acid or gastric juice
Industrial cleaners, (Lysol)
Nitric acid HNO3 Fertilizers, dyes
Phosphoric acid H3PO4 Soft drinks (tart flavor)
Fertilizers, detergents
Sulfuric acid H2SO4 Automobile battery Fluid
Citric acid H3C6H5O7 Orange juice, lemons, citrus fruits
Common Bases used in everyday life
Base Formula Use
Sodium hydroxide NaOH Drain &Oven cleaners;
“Liquid Drano”
Calcium hydroxide Ca(OH)2 Plaster, cement
Potassium hydroxide KOH Liquid soaps
Magnesium hydroxide Mg(OH)2 Antacids: Milk of magnesia, Tums”
Sodium bicarbonate NaHCO3 Baking Soda
Ammonium Hydroxide NH4OH House-hold ammonia
Conjugate Acid – Base Pairs
Conjugate means joined together
An acid loses a H+ ion becomes its conjugate base
HCl → H + + Cl – ( Cl - = conjugate base)
H2O → H + + OH – (OH - = conjugate base)
A Base gains an H + ion becomes its conjugate acid, (protonated base)
NH3 + H + → NH4 + (NH4 + = conjugate acid)
Conjugate Acid – Base pair differ by one H +
NH3 + H2O → NH4+ + OH –
(B) (A) (C.A.) (C.B.)
Acid – Base Classification
Acids and bases classified as strong or weak.
Acid and base classification as strong or weak depends on the degree dissociation into ions in solution.
Acids classified as strong or weak depending on the number of hydrogen ions ( H+) produced in water solution.
Bases classified as strong or weak depending on the number of hydroxide ions (OH-) produced in water solution.
Strong acid
· High degree of acid dissociation into ions (~100%)
· High concentration of H+ ions in solution
· Strong acids are corrosive
Weak acid
· Slight degree of acid dissociation into ions
· Low concentration of H+ ions in solution
Strong Base
· High degree of base dissociation into ions
· High concentration of OH ‾ ions in solution
Weak Base
· Slight degree of base dissociation into ions
· Low concentration of OH‾ ions in solution
Common strong acids and weak acids
Strong acids
1. Nitric acid
2. Sulfuric acid
3. Hydrochloric acid
Weak acids
1. Citric acid (orange juice)
2. Carbonic acid (soda)
3. Acetic acid (vinegar)
Strong Bases
1. Sodium hydroxide NaOH
2. Potassium hydroxide KOH
Weak Bases
1. Magnesium hydroxide Mg(OH)2
2. Ammonia NH3
Note:
Ammonia is a base that does not contain an (OH‾) ion in its structure; however, ammonia reacts with water to produces hydroxide ions (OH‾).
NH3 + H2O → NH4+ + OH ‾
Key point to review,
The “strong” or “weak” classification of an acid or base refers to the degree of dissociation, not the concentration of the acid or base solution.
A strong acid does not mean that the solution is concentrated, nor does a weak acid mean that the solution is dilute.
Given two 0.1 M solutions of Nitric acid and Acetic acid,
If nitric acid is spilled on clothing, it will burn through the fabric,
However, if the acetic acid solution is spilled on clothing it will not be corrosive.
Why?
Even though both acid solutions are at the same concentration, Nitric acid is a stronger acid than acetic acid. Therefore, nitric acid will have a higher concentration of hydrogen ions than acetic acid, which makes the nitric acid corrosive.
Acetic acid on the other hand is a weak acid. This classification means that solution of acetic acid contains more acid molecules that H+ ions.
Gastric juice is a dilute (5% by weight) solution of the strong acid HCl.
However, a 35 % by weight solution of HCl is a concentrated solution of a strong acid
Neutralization reactions
Acid + Base = Salt + Water
HCl + NaOH ------NaCl + H2O
Hydrochloric acid + Sodium hydroxide = Sodium chloride + Water
(Table salt)
Antacid tablet “ Tums “
2HCl + Mg(OH)2 ------MgCl2 + 2 H2O
Magnesium hydroxide partially neutralizes excess stomach acid
Self – ionization of water
2 H2O (l) → H3O + (aq) + OH – (aq)
In pure water at 25 °C, both H3O + and OH – are at equal concentrations of x 10 – 7 Molar
[ ] = Molarity ( M )
[ H3O + ] [OH - ] = 1.00 x 10 – 14 M2
Acidic Solution: [ H3O + ] > 10 – 7 M or [ OH – ] < 10 – 7 M
Neutral Solution [H3O +] = [OH – ] = 10 – 7 M
Basic Solution : [ H3O +] < 10 – 7 M or [ OH - ] > 10 – 7 M
pH scale
The pH scale measures the acidic, basic or neutral character of a solution.
The pH scale ranges from 0 to 14.
0 ------7 ------14
Acidic Neutral Basic
· PH = 7 indicates a neutral solution
· PH < 7 indicates an acidic solution
· PH > 7 indicates a basic solution
Low pH represents a more acidic solution
High pH represents a more basic solution
Lowering the pH of a solution corresponds to an increase in the hydrogen ion concentration; therefore the solution is more acidic.
Measuring pH
Use an Acid – Base Indicator
Indicators are weak acids or bases that change color when they gain or lose an H+ ion.
Examples:
1. Litmus paper
2. Phenolphthalein
Litmus turns Red in Acidic solutions
Litmus turns Blue in Basic solutions
Phenolphthalein is Colorless or cloudy in Acids and Pink in Bases
Mathematically define pH using common logarithms
A common logarithm is the exponent which base 10 is raised to obtain that number.
1,000 = 10 x 10 x 10 = 10 3
log (1,000) = 3 ; or log (10 3) = 3
1 / 1000 = 0.001 = 10 – 3
log (0.001) = − 3
[ H+] is very small, therefore use the
pH scale
Low pH = higher acidity or higher proton concentration
For weak acids,
pH ≠ [ Acid ]
pH measures the [ H + ions] or [H3O+ ions]
pH is the negative common log of the hydronium ion concentration.
pH = − log10 [ H3O + ]
10 – pH = [ H3O + ]
Note: 1og scale is a 10-fold change in [H+]
If pH drops from 4 to 3, then [H+] increases from 10 – 4 M to 10 –3 M
Lower pH = increase [H+] = more acidic solution
Summary
pH = − log [ H+]
pOH = − log [ OH - ]
pH + POH = 14
For a strong acid
HX → H+ + X – (100% ionization)
1 mole acid yields 1 mole of protons
Therefore, [ strong acid ] = [ H + ions]
pH = - log [ strong acid ]
For a strong base,
MOH → M+ + OH – (100% dissociation)
1 mole of strong base yields 1 mole of OH - ions
[strong base ] = [OH - ]
pOH = - log [ strong base]
Concentrations of Acids & Bases
Chemical Equivalents (Equiv)
Solute amount having equivalent combining capacity
Acid equivalents for neutralization rxn with KOH
HCl → H + + Cl –
(1 mol HCl → 1 mol H + )
H2SO4 → 2 H + + SO4 – 2
( 1 mol H2SO4 → 2 mol H + )
HCl ≈ H2SO4
1mol ≈ ½ mol
36.5 g ≈ 49.0 g ( ½ 98 g)
In proton transfer reactions,
Use Equivalent mass of acid:
Acid mass in grams that provides one mole of H + ions, (protons)
M eq (acid) = Acid Molar Mass
mole H + ions
units: g / mol H +
Determine mass for 1 equivalent of HCl
Meq = 36 .5 g HCl x 1 mol HCl
Mol HCl 1 mol H +
= 36.5 g HCl
mol H +
Mass for one equivalent of H2SO4,
Meq = 98.1 g H2SO4 x 1mol H2SO4
1 mol H2SO4 2 mol H +
= 49.0 g H2SO4
mol H +
Calculate Equivalent masses for:
HNO3 , H2SO3
Meq (HNO3) = 63.0 g HNO3
mol H +
Meq (H2SO3) = 41.0 g H2SO3
mol H +
Equivalent Mass of Base
Base mass in grams that reacts with 1 mole H + or provides 1 mole of OH –
Meq (Base) = Base Molar mass
mol OH – ions
KOH → K + + OH –
Meq KOH = 56.1g KOH x 1mol KOH
1 mol KOH 1 mol OH –
= 56.1 g KOH / mol OH –
Ca(OH)2 → Ca+2 + 2OH –
74.1 g Ca(OH)2 x 1 mol Ca(OH)2
1 mol Ca(OH)2 2 mol OH –
Meq Ca(OH)2 = 37 g Ca(OH)2 / mol OH –
Normality – Solution Concentration Unit
# Solute equivalents per solution Liter ( Equiv. = mol H + or OH – ions)
Normality (N) = equiv solute / L solution
One – Normal solution = 1 N
0.25 N = 0.25 solute equiv. / L solution
For monoprotic acid HCl,
36.5 g HCl = 1 mol H + = 1 N
Normality of 1 mol of diprotic acid H2SO4 in 1 L solution
2 equivalents x 1 mol H2SO4 = 2 equiv
1 mol H2SO4 1 L 1 L
= 2 N H2SO4
Acid – Base Titration
Controlled, nearly complete, neutralization reaction to determine the acid or base concentration in a solution
Materials:
Standard solution
Contains an acid or base of a known concentration
Indicator – substance that changes color at a certain pH
(phenolphthalein) color change over a pH range ( 8 – 10)
Buret – finely calibrated pipette for exact volume measurements
Equivalence point
Point when enough standard solution is added to neutralize all acid or base in the unknown solution
End Point (Observed)
Point at which indicator changes color
End Point ≈ Equivalence Point
At the end point,
Total moles H + Total moles H +
donated by Acid = accepted by Base
Total moles H + = Total moles OH –
(from acid ) (from base)
Normality and Titrations
Equal volumes of different solutions of same normality are chemically equivalent
Equiv. (acid) = Na Va = ( equiv. x L)
L
At the end point,
Acid Equivalence = Base Equivalence
NaVa = Nb Vb
Na = acid Normality
Va = acid volume
Nb = base Normality
Vb = base volume
Acid Rain
· Rain that contains acid
· Environmental pollutant (contaminates lakes, lowers pH of lake water)
· Harmful to plants and fish
Acid rain originates from the burning of fossil fuels such as coal, oil, methane
Industry burns coal to produce energy,
“dirty” Coal contains sulfur
Sulfur + Oxygen = Sulfur dioxide gas is released into the atmosphere
S + O2 → SO2
SO2 + H2O → H2SO3
( Moist air ) Sulfurous acid
Acid rain
Normal Rain pH = 5.6
Acid Rain pH = 4.0 – 4.5