Acids-Base Equilibria

Acids and Bases:

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Acid: A substance that releases H+ ions in an aqueous solution.
Example: HCl (hydrochloric acid)
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Base: A substance that releases OH- ions in an aqueous solution.
Example: NaOH (sodium hydroxide)
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Bronsted-Lowry (B.L.): Focus on Reactions
Bronsted: Lowry: B.L.
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Acid: A proton (H+) donor "in a reaction".
Base: A proton (H+) acceptor "in a reaction".
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Example: NH3 + H2O = NH4+ + OH- (in equilibrium)
H2O: Acid
NH3: Base
NH4+: Conjugate Acid
OH-: Conjugate Base

Amphoteric: Some substances can act as an acid in one reaction and as a base in another. These substances are called amphoteric. An amphoteric substance can act as an acid in some situations and a base in other situations. It acts as a base when combined with something more acidic than itself. It acts as an acid when combined with something more basic than itself. An example of an amphoteric substance is water. It can either accept or donate a proton.
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Conjugate Acid-base pairs

Conjugate Acid: A substance formed by addition of a proton to a Bronsted-Lowry base.
Conjugate Base: A substance formed by the loss of a proton from a Bronsted-Lowry acid.
Conjugate Acid-Base Pair: An acid and a base, such as H2O and OH-, that differ only in the presence or absence of a proton.
Base Conjugate Acid Examples:
NH3/NH4+ (relation: H+)

Acid Conjugate Base Examples:

HF + H2O = H3O +F- (in equilibrium)
H3O: Hydronium ion. Water with extra H+. Conjugate Acid
HF: Acid
H2O: Base
F-: Conjugate Base
H+ (congruent with) H3O+

An acid and a base always work together to transfer a proton. A substance can function as an acid only if another substance simultaneously behaves as a base. To be a Bronsted-Lowry acid, a molecule or ion must have a hydrogen atom that it can lose as an H+ ion. To be a Bronsted-Lowry base, a molecule or ion must have a nonbonding pair of electrons that it can use to bind the H+ ion.



Strength of Acids and Bases (See Figure 16.4)

Two important facts:

1.) The stronger the acid, the weaker its conjugate base, the stronger the base, the weaker its conjugate acid.

2.) In every acid-base reaction the position of the equilibrium favors the transfer of the proton from the stronger acid to the stronger base to form the weaker acid and the weaker base.

Read Section 16.1,2 Do #’s 16.1, 2 p.709 16.13-16.27 odd pp. 710-711.

The Autoionization of water (water attacks itself – but it’s rare)

Arrhenius -
HOH --> H+ + OH- Kw: [H+][OH-] = 1x10^-14 @ 25C

Bronsted-Lowry -//
H2O + H2O ---> H3O+ + OH- Kw = [H3O+][OH-] = 1x10^-14 @ 25C

2 per billion.


Kw -
K of Water is 1x10^-14 @ 25C
Kw = [H+][OH-] =[H3O+][OH-] = 1x10^-14


The pH Scale
power [H+]
EX: [H+] = 1x10^-3 pH = 3

pH = -log[H+] = -log[H3O+]
log is base 10
exponential scale

pOH = -log [OH-]

*ALWAYS: pH + pOH = 14


If you know the Hydrogen concentration, you can get the Hydroxide.
EX: If pH=4 ---> 1x10^-4 MH+ [OH-] = 1x10^-14/1x10^-4 = 1x10-10M OH-

Pure water
If reaction takes place in PURE water, [H3O+] must equal [OH-], they are bot 7.
pH of 7 is PURE H2O

Acids and Bases

Acids: 0-7
Bases: 7-14

Playing with pH. (YCP WS)

Measuring pH

Sources Used:

- OR

Read Section 16.3,4 Do #’s 16.29-39 odd, 16.40, 42, pp 711-712.

Strong Acids and Bases

List the 7 strong acids. Do it now! Don’t look at your neighbor’s paper.
Hydrochloric Acid, HCl
Hydrobromic Acid, HBr
Hydroiodic Acid, HI
Chloric Acid, HClO3
Perchloric Acid, HClO4
Nitric Acid, HNO3
Sulfuric Acid, H2SO4

What does it mean to be a strong acid?

~ Strong acid is an acid that dissociates completely in an aqueous solution by losing one proton.
~ Strong acids are strong electrolytes.
~ They have a very high equilibrium constant value.

Strong bases What are they? What does it mean to be a strong base?

~ Strong bases are bases that completely dissociate in water.

~ Strong bases include:
Hydroxides of Group I metals
Hydroxides of Group II metals

Weak Acids

~ Substances that increase the concentration of H+ but are not fully ionized (see examples on p 682).
~ Basically anything that is not a strong acid. HX(aq) + H2O(l) ⇌ H3O+ + X-


The higher the Ka the stronger the acid.
The acid dissociation constant is the equilibrium constant of the dissociation reaction of an acid and is denoted by Ka.
Ka = [H3O+] [X-] / [HX]

Calculating pH from Ka and using pH to calculate Ka.s
(See examples from worksheet: "ACIDS AND BASES: WORKSHEET 3")

Polyprotic Acids

~ Polyprotic Acid are able to donate more than one proton per acid molecule
~ It is always easier to remove the first proton from a polyprotic acid than the second. In most cases a satisfactory estimate of the pH can be made by just considering
the Ka1.

More specific names are DIPROTIC ACIDS (two potential protons to donate) and TRIPROTIC ACIDS (three potential protons to donate)

Diprotic Acid Examples: Triprotic Acid Examples:

Sulfuric acid (H2SO4) Phosphoric acid (H3PO4)
Carbonic acid (H2CO3) Citric acid (C6H8O7)
Hydrogen sulfide (H2S)
Chromic acid (H2CrO4)
Oxalic acid (H2C2O4)

Read Sections 16.5 and 16.6 for more information.
For practice questions, do #’s 16.3,4, 5, on page 709, and 16.43-16.67 odd questions on pages 712-713.

Weak Bases (B)

B- + H2O ---> HB + OH-
Kb= [HB][OH-]

B + H2O ---> HB+ + OH-
Kb= [HB+][OH-]

Kb- you'll find acid/conjugate base pairs
Most common weak base: Ammonia (NH3)

NH3 + H2O ---> NH4+ + OH-

Categories of weak bases: (See Table16.4 p 691)

(1) Amines- (smell bad)
Organic compounds that contain Nitrogen- Bitter vegetables


(2) Anions from weak acids
anions- negative charged ions
Ex: (Most common)
NO2- + H2O <--> HNO2 + OH- (raise pH)


YCP again

Relationship between Ka and Kb:

Ka*Kb= Kw

For example:
  • Ka= [H+][A-]/[HA]
  • Kb=[HA][OH-]/[A-]
    • [H+][A-]/[HA] X [HA][OH-]/[A-] = Kx
      • note: notice that [HA], [A-] cancel out, leaving you with [H+] AND [OH-]

Conjugate acid-base pair

  • a conjugate pair refers to acids and bases with common features, such as equal loss or gain of atoms.
    • acid and its conjugate base
Some examples of conjugate acid-base pairs.

Conjugate Base
Conjugate Acid

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Kw = Ka x Kb

Read section 16.7 and 16.8 Do #’s 16.6, 7 p. 710, #’s 16.71-89 odd pp. 713-714.

Acid-Base Properties of Salt Solutions.

Remember if a salt (which is an ionic substance) dissolves, it dissociates.
Salt Dissociates in Water

  • Therefore, The acid-base properties of salt solutions are due to the behavior of the constituent anions and cations.
Question is what does the cation and/or the anion do when it is presented with water?
    • We all know that many substances such as HCl or HBr, are acidic, and other substances such as CaOH and NH3, are basic. However, we have also learned that
ions can also exhibit acidic or basic properties in solutions.
Many ions are able to react with water in a process known as HYDROLYSIS: a chemical reaction in which water molecules are split into H+/H30+ and OH- ions.


  • Neutral (no change in pH) : Group I and II

    • In other words, when Cations from Groups I and II are presented with water, they do not change the pH of the solution.
  • Acidic. (ALL OTHER METAL CATIONS PLUS NH4). These change the pH, and make it acidic.

    • ex: ammonium reacts with water to make hydronium (H30) ions and ammonia.
Can NEVER make solution basic.

Hydrolysis of a hydrated Aluminum ion


an anion in solution can be considered the conjugate base of an acid. Therefore to identify an acid, simply add a proton
to the anion's formula
  • Neutral: any anion derived from strong acids
    • REMEMBER THE SEVEN STRONG ACIDS but do not confuse the anion from the strong acid with the strong acid itself.
    • Strong_Acids.jpg
      • strong acids affect the pH of a solution. On the other hand, anions from strong acids DO NOT affect the pH of the solution.

  • Basic*: these are anions which are from weak acids.
    • these do affect the pH because these anions react to a small extent with water to form a weak acid and hydroxide ions.
    • ex: Fluoride, Nitrite, and Acetate
    • acetate.jpg
    • C2H3O2+ H2O <=> CH3COOH + OH-

*H2PO4 and HSO4- are anions that acts as acids, and are therefore exceptions to this rule.

In Summary:

  • An anion that is the conjugate base of a strong acid does not affect the pH of the solution. (ACT AS SPECTATOR IONS)
  • A cation that is the conjugate acid of a strong base does not affect the pH of the solution. (ACT AS SPECTATOR IONS)
  • An anion that is the conjugate base of a WEAK ACID will cause pH to INCREASE.
  • A cation that is the conjugate base of a WEAK BASE will cause pH to DECREASE.


The Combined Effect

  • When a solution contains both anions and cations that affect pH, the ion with the larger equilibrium constant.
    • If Ka>Kb ; acidic solution
    • If Kb>Ka ; basic solution
    • (Use the pink reference sheet)

  • NH4F - two effects, the NH4+ is acidic and the F- is basic; in this case the solution is acidic.
    • Kb F- = 1*10^-14 / 6.8 * 10^-4 = 1.5*10^-11
      • F- + H2O <--> HF + OH- Kb = 1.5*10^-11
    • Ka NH4+ = 1*10^-14 / 1*10^-5 = 1*10^-9
      • NH4F + H2O <--> H30+ + NH3 Ka = 1*10^-9
    • Acidic because the NH4+ value of Ka is higher than the F- value of Kb.

Acid-base Behavior and Chemical Structure

Three factors affect acid strength:
  • the polarity of the H-X bond
    • If X is highly electronegative then the acid strength increases
    • H2O vs. HF
      • Fluorine is greater because of stronger electronegativity

  • the strength of the H-X bond
    • Bond strength: stronger bond means a weaker acid
    • HF - weak acid because water cannot pull the ions apart
    • Has to do with the size of F-

  • the stability of the conjugate base X-
    • Conjugate base: the more stable the conjugate base, the stronger the acid.

Binary Acids

Strongest Bond

Weakest Bond
HF <<
HCl <
HBr <
Weakest Acid

Strongest Acid

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*Use bond strength for binary acids


  • Have Oxygen in them

  • - H - O - Cl +
-------> electron density

  • - H - O - Na +
<------ electron density

  • Example - which is strongest?
    • H - O - Cl
    • H - O - Br
    • H - O - I
    • HOCl because it's the most electronegative
    • HOI is not the strongest because the Hydrogen and Iodine are not directly attached

  • HOCl vs. HClO2 vs. HClO3 vs. HClO4
    • Get stronger as you add Oxygen
    • HOCl <--> H+ + OCl-
      • Electron density is confined; less stable
    • HClO4 --> H+ + ClO4-
      • Electron density is distributed amongst 5 atoms; more stable = strong acid
  • Resonance
    • NO2 vs. NO3
      • More resonance in NO3 means it's more stable and thus a stronger acid
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*Use electronegativity for Oxyacids
*Ability to disperse charge and electronegativity affect it

Carboxylic acids

R = "other stuff"

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  • Acetic Acid / Ethanoic Acid
    • HC2H3O2 <--> C2H3O2- + H+
    • C2H3O2- has resonance, but still is a weak acid
  • Chloroethanoic Acid
    • HC2H2ClO2 <--> C2H2ClO2- + H+
  • Chloroethanoic Acid is stronger due to the electronegativity of Chlorine

*Affected most by the electronegativity



Read Section 16.9,10 Do #’s 16.9, 10, 11 p. 710. 16.91-97 odd p. 714.

Lewis Acids and Bases

*More general than BL Acid-Base because it focuses on electron pairs and not on protons (H+)

Acid: An electron pair acceptor.

Base: An electron pair donor.

NH3 + H2O -> NH4+ + OH-

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-Ammonia is Lewis Base (donates unpaired elections to H+)
-Water is Lewis Acid

Classic example: Ammonia plus boron trifluoride.

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-Ammonia is the Lewis Base and Boron Trifluoride is the Lewis Acid

Other Examples: Cations like Al3+ and Fe3+

Al3+ + 6H2O -> Al(H20)6^3+ -> (Al(H2O)5)OH)^+2

The Amphoteric Behavior of Amino Acids. (See page 703)

-the building blocks of protein, 20 types of Amino Acids
-composed of an Amino group (NH2), Hydrogen, carboxyl group (COOH) around a Carbon and an R group that varies
-Putting an Amino acid in an acidic solution causes protonation, while a base causes deprotanation
Protonated Amino Acid
Protonated Amino Acid

-At a certain pH the Zwitterion causes the amino acid to be charged on each end

-Two Amino Acids can form a peptide bond using the (-OH) from the carboxyl group and an (H+) from the Amino group making a water molecule and a peptide (dehydration synthesis)
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