Acid and base concepts

1. General Definitions:

Acid: a substance which when added to water produces hydrogen ions [H⁺].

Base: a substance which when added to water produces hydroxide ions [OH^-].

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2. Properties:

Acids:

• react with zinc, magnesium, or aluminum and form hydrogen (H_2(g))
• react with compounds containing CO₃^2- and form carbon dioxide and water
• turn litmus red
• taste sour (lemons contain citric acid, for example) DO NOT TASTE ACIDS IN THE LABORATORY!!

Bases:

• feel soapy or slippery
• turn litmus blue
• they react with most cations to precipitate hydroxides
• taste bitter (ever get soap in your mouth?) DO NOT TASTE BASES IN THE LABORATORY!!

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3. Water dissociation: H₂O_(l)  →  H⁺_(aq) + OH^-_(aq)

equilibrium constant, K_W = [H⁺][OH^-] / [H₂O] Note:  water is not involved in the equilibrium expression because it is a pure liquid, also, the amount of water not dissociated is so large compared to that dissociated that we consider it a constant
Value for K_w = [H⁺][OH^-] = 1.0 x 10^-14
Note: The reverse reaction, H⁺_(aq) + OH^-_(aq)  →   H₂O_(l) is not equal to 1 x 10^-14

[H⁺] for pure water = 1 x 10^-7
[OH^-] for pure water = 1 x 10^-7

Definitions of acidic, basic, and neutral solutions based on [H⁺]

• acidic: if [H⁺] is greater than 1 x 10^-7 M
  
• basic: if [H⁺] is less than1 x 10^-7 M
  
• neutral: if [H⁺] if equal to 1 x 10^-7 M

Example 1: What is the [H⁺] of a sample of lake water with [OH^-] of 4.0 x 10^-9 M? Is the lake acidic, basic, or neutral?
Solution: [H⁺] = 1 x 10^-14 / 4 x 10^-9 = 2.5 x 10^-6 M
Therefore the lake is slightly acidic
Remember: the smaller the negative exponent, the larger the number is.
Therefore:

• acid solutions should have exponents of [H⁺] from 0 to -6.
  
• basic solutions will have exponents of [H⁺] from -8 on.

Example 2: What is the [H⁺] of human saliva if its [OH^-] is 4 x 10^-8 M? Is human saliva acidic, basic, or neutral? Solution: [H⁺] = 1.0 x 10^-14 / 4 x 10^-8 = 2.5 x 10^-7 M
The saliva is pretty neutral.

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4. pH relationship between [H⁺] and pH

pH = -log₁₀[H⁺]

Definition of acidic, basic, and neutral solutions based on pH

acidic: if pH is less than 7
basic: if pH is greater than 7
neutral: if pH is equal to 7

The [H⁺] can be calculated from the pH by taking the antilog of the negative pH
Example 3:  calculate the [OH^-] of a solution of baking soda with a pH of 8.5.
Solution:  First calculate the [H⁺]

if pH is 8.5, then the antilog of -8.5 is 3.2 x 10^-9. Thus the [H⁺] is 3.2 x 10^-9 M

Next calculate the [OH^-]

1.0 x 10^-14 / 3.2 x 10^-9 = 3.1 x 10^-6 M

Example 4:  Calculate the pH of a solution of household ammonia whose [OH^-] is 7.93 x 10^-3 M.
Solution:  This time you first calculate the [H⁺] from the [OH^-]

7.93 x 10^-3 M OH^- = 1.26 x 10^-12 M H⁺

Then find the pH

-log[1.26 x 10^-12] = 11.9 

 

5. Strength of Acids and Bases: Acids

1. Strong Acids:

• completely dissociate in water, forming H⁺ and an anion. example: HN0₃ dissociates completely in water to form H⁺ and N0₃^1-.
  The reaction is
  
  HNO_3(aq)  →   H⁺_(aq) + N0₃^1-_(aq)
  A 0.01 M solution of nitric acid contains 0.01 M of H⁺ and 0.01 M N0₃^- ions and almost no HN0₃ molecules. The pH of the solution would be 2.0.
  
• There are only 6 strong acids: You must learn them. The remainder of the acids therefore are considered weak acids.
  1. HCl
     
  2. H₂SO₄
     
  3. HNO₃
     
  4. HClO₄
     
  5. HBr
     
  6. HI
• Note: when a strong acid dissociates only one H⁺ ion is removed. H₂S0₄ dissociates giving H⁺ and HS0₄^- ions.
  H₂SO₄  →   H⁺ + HSO₄^1-
  A 0.01 M solution of sulfuric acid would contain 0.01 M H⁺ and 0.01 M HSO₄^1- (bisulfate or hydrogen sulfate ion).

2. Weak acids:

• a weak acid only partially dissociates in water to give H⁺ and the anion
  for example, HF dissociates in water to give H⁺ and F^-. It is a weak acid. with a dissociation equation that is
  HF_(aq)  ↔   H⁺_(aq) + F^-_(aq)
• Note the use of the double arrow with the weak acid. That is because an equilibrium exists between the dissociated ions and the undissociated molecule. In the case of a strong acid dissociating, only one arrow (  →  ) is required since the reaction goes virtually to completion.
• An equilibrium expression can be written for this system:
  K_a = [ H⁺][F^-] / [HF]
• Which are the weak acids?   Anything that dissociates in water to produce H⁺ and is not one of the 6 strong acids.
  1. Molecules containing an ionizable proton. (If the formula starts with H then it is a prime candidate for being an acid.) Also: organic acids have at least one carboxyl group, -COOH, with the H being ionizable.
  2. Anions that contain an ionizable proton. ( HSO₄^1-  →   H⁺ + SO₄^2- )
  3. Cations:  (transition metal cations and heavy metal cations with high charge)
     also NH₄⁺ dissociates into NH₃ + H⁺

Bases

1. Strong Bases:

• They dissociate 100% into the cation and OH^- (hydroxide ion).
  example:  NaOH_(aq)  →   Na⁺_(aq) + OH^-_(aq)
  a. 0.010 M NaOH solution will contain 0.010 M OH^- ions (as well as 0.010 M Na⁺ ions) and have a pH of 12.
  
• Which are the strong bases?
  The hydroxides of Groups I and II.
• Note: the hydroxides of Group II metals produce 2 mol of OH^- ions for every mole of base that dissociates. These hydroxides are not very soluble, but what amount that does dissolve completely dissociates into ions. exampIe: Ba(OH)_2(aq)  →   Ba²⁺_(aq) + 2OH^-_(aq)
  a. 0.000100 M Ba(OH)2 solution will be 0.000200 M in OH^- ions (as well as 0.00100 M in Ba²⁺ ions) and will have a pH of 10.3.

2. Weak Bases:
What compounds are considered to be weak bases?

1. Most weak bases are anions of weak acids.
2. Weak bases do not furnish OH^- ions by dissociation. They react with water to furnish the OH^- ions.
   Note that like weak acids, this reaction is shown to be at equilibrium, unlike the dissociation of a strong base which is shown to go to completion.
3. When a weak base reacts with water the OH^- comes from the water and the remaining H⁺ attaches itsef to the weak base, giving a weak acid as one of the products. You may think of it as a two-step reaction similar to the hydrolysis of water by cations to give acid solutions. examples:
   
   NH_3(aq) + H₂O_(aq)  →   NH₄⁺_(aq) + OH^-(aq) methylamine: CH₃NH_2(aq) + H₂0_(l)  →   CH₃NH₃⁺_(aq) + OH^-_(aq)
   acetate ion: C₂H₃O₂^-_(aq) + H₂O(aq)  →  HC₂H₃0_2(aq) + OH^-_(aq)
   General reaction: weak base_(aq) + H₂O_(aq)  →   weak acid_(aq) + OH^-_(aq)
   
   Since the reaction does not go to completion relatively few OH- ions are formed.

Acid-Base Properties of Salt Solutions:

definition of a salt:

• an ionic compound made of a cation and an anion, other than hydroxide.
• the product besides water of a neutralization reaction

determining acidity or basicity of a salt solution:

1. split the salt into cation and anion
   
2. add OH^- to the cation
   
   a. if you obtain a strong base. the cation is neutral
   b. if you get a weak base, the cation is acidic
3. Add H⁺ to the anion
   
   a. if you obtain a strong acid, the anion is neutral
   b. if you obtain a weak acid. the anion is basic
4. Salt solutions are neutral if both ions are neutral
   
5. Salt solutions are acidic if one ion is neutral and the other is acidic
   
6. Salt solutions are basic is one of the ions is basic and the other is neutral.
   
7. The acidity or basicity of a salt made of one acidic ion and one basic ion cannot be determined without further information. 6. Acid-Base Reactions:
   

   
   
   
   

   Strong acid + strong base:  HCl + NaOH  →   NaCl + H₂O

   net ionic reaction:  H⁺ + OH^-  →   H₂O

   Strong acid + weak base: example:  write the net ionic equation for the reaction between hydrochloric acid, HCl, and aqueous ammonia, NH₃. What is the pH of the resulting solution?
   

   Strong base + weak acid: example:  write the net ionic equation for the reaction between citric acid (H₃C₆H₅0₇) and sodium hydroxide. What is the pH of the resulting solution?

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   7. Titrations
   

   1. Nomenclature:  these are terms that are used when talking about titrating one substance with another. You need to learn these definitions well enough to explain them to someone else.
   

   • titration
     
   • titrant
     
   • indicator
     
   • equivalence point
     
   • end point
     
   • titration cuve
     

   2. Strong acid-strong base titration
   

   example: titration curve
   pH at equivalence point
   species present
   appropriate indicators

   3. Strong acid-weak base titration
   

   example titration curve
   pH at end point
   species present
   appropriate indicators
   8. Three models of acids:
   

   l. Arrhenius Model Basis for the model--action in water
   
   

   • acid definition: produces H in water solution
   • base definition: produces OH^1- in water solution

   2. Bronsted-Lowry Model
   Basis for the model-- proton transfer
   
   

   • acid definition: donates a proton ( H )
   • base definition: accepts a proton
   • conjugate acid definition: the acid becomes the conjugate base after it donates the proton because it can now accept it back.
   • conjugate base definition: the base becomes the conjugate acid after it accepts the proton because it can now donate it back.

   3. Lewis Model
   Basis for model--electron pair transfer
   
   

   • acid definition: accepts a pair of electrons
   • base definition: donates a pair of electrons