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Section 2.1: Acids and Bases. Solutions. Learning Goals:. Identify the physical and chemical properties of acids and bases. Classify solutions as acidic, basic, or neutral. Compare the Arrhenius and Brønsted -Lowry models of acids and bases. Properties of Acids and Bases.

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Section 2.1: Acids and Bases

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Section 2.1: Acids and Bases

Solutions


Learning Goals:

  • Identify the physical and chemical properties of acids and bases.

  • Classify solutions as acidic, basic, or neutral.

  • Compare the Arrhenius and Brønsted-Lowry models of acids and bases.


Properties of Acids and Bases


Properties of Acids and Bases

  • All water solutions contain hydrogen ions (H+) and hydroxide ions (OH–).

  • An acidic solutioncontains more H+ions than OH-ions.

  • A basic solutioncontains more OH-ions than H+ions.


Properties of Acids and Bases

  • The usual solvent for acids and bases is water.

    • Water produces equal numbers of H+and OH-ions in a process called self-ionization.

      H2O(l) + H2O(l) H3O+(aq) + OH–(aq)


Arrhenius Model

  • States that an acid is a substance that contains hydrogen and ionizes to produce hydrogen ions in aqueous solution

  • States that a base is a substance that contains a hydroxide group and dissociates to produce a hydroxide ion in solution.


Arrhenius Model

  • HCl ionizes to produce H+ ions.

    • HCl(g) → H+(aq) + Cl–(aq)

  • NaOH dissociates to produce OH– ions.

    • NaOH(s) → Na+(aq) + OH–(aq)

  • Some solutions produce hydroxide ions even though they do not contain a hydroxide group


Brønsted-Lowry Model

  • States that an acid is a hydrogen ion donor, and a base is a hydrogen ion acceptor.

  • The Brønsted-Lowry Model is a more inclusive model of acids and bases.


Brønsted-Lowry Model

  • A conjugate acidis the species produced when a base accepts a hydrogen ion.

  • A conjugate baseis the species produced when an acid donates a hydrogen ion.


Brønsted-Lowry Model

  • A conjugate acid-base pair consists of two substances related to each other by donating and accepting a single hydrogen ion.


Brønsted-Lowry Model

  • HF(aq) + H2O(l) ↔ H3O+(aq) + F–(aq)


Brønsted-Lowry Model

  • HF = acid

  • H2O = base

  • H3O+ = conjugate acid

  • F– = conjugate base


Brønsted-Lowry Model

  • NH3(aq) + H2O(l) ↔ NH4+(aq) + OH–(aq)

  • NH3 =

  • H2O =

  • NH4+ =

  • OH–=


Acid Conjugate Base


Section 2.2: pH and pOH

Solutions


Learning Goals:

  • Explain pH and pOH.

  • Relate pH and pOH to the ion product constant for water.

  • Calculate the pH and pOH of aqueous solutions.


Water

  • Pure water contains equal concentrations of H+ and OH– ions.

    H2O  H+ + OH-


Water

  • The ion product constant of water when water self-ionizes is Kw

    Kw= [H+][OH–]


Water

  • With pure water at 25°C, both [H+] and [OH–] are equal to 1.0 × 10–7M.

    Kw = 1.0 × 10–14

1.0 x 10-14= [H+][OH–]


Water

  • In a neutral solution, [H+] = [OH-]

  • In an acidic solution, [H+] > [OH-]

  • In a basic solution, [H+] < [OH-]

[H+][OH–]= 1.0 x 10-14


Practice

  • Calculate the ion concentration for each of the following solutions and state whether they are neutral, acidic, or basic.

    • 1.0 x 10-5 M OH- calculate H+

    • 1.0 x 10-7 M OH-  calculate H+

    • 10.0 M H+  calculate OH-


pH and pOH

  • pH is a measurement of the concentration of hydrogen ions.

    • Low pH = acid

    • High pH = base


pH and pOH

  • Calculating pH:

    • pH = –log [H+]


Calculating pH

  • Enter the [H+]

  • Press the “log” key

  • Press the “+/-” (change of sign) key


Practice

  • Calculate the pH value of the following:

    • A solution in which [H+] = 1.0x10-9

    • A solution in which [OH-] = 1.0x10-6


pH and pOH

  • pOHis a measurement of concentration of hydroxide ions.

    • Low pOH = base

    • High pOH = acid


pH and pOH


pH and pOH

  • Calculating pOH

    • pOH= –log [OH–]


Calculating pOH

  • Enter the [OH-]

  • Press the “log” key

  • Press the “+/-” (change of sign) key


Practice

  • Calculate the pH and pOH values of the following:

    • 1.0 x 10-3 M OH-

    • 1.0 M H+


pH and pOH

pH + pOH = 14


Calculating [H+] from pH

  • Enter the pH

  • Press the “+/-” (change of sign) key

  • Take the inverse log by pressing “10x” button (inv log)


Practice

  • The pH of a human blood sample was measured to be 7.41. What is the [H+] in this blood?


Practice

  • Calculate [H+] and [OH-] in each of the following solutions:

    • Milk, pH = 6.50

    • Ammonia, pH = 11.90


Section 2.3: Neutralization

Solutions


Learning Goals:

  • Write chemical equations for neutralization reactions.

  • Explain how neutralization reactions are used in acid-base titrations.

  • Compare the properties of buffered and unbuffered solutions.


Acid-Base Reactions

  • A neutralization reactionis a reaction in which an acid and a base in an aqueous solution react to produce a salt and water.

  • A salt is an ionic compound made up of a cation from a base and an anion from an acid.

  • Neutralization is a double-replacement reaction.


Acid-Base Reactions


Practice

  • What volume of a 1.00M HCl solution is needed to neutralize 25.0 mL of a 0.350 M NaOH solution?


What volume of a 1.00M HCl solution is needed to neutralize 25.0 mL of a 0.350 M NaOH solution?

1.) Balanced equation:

HCl(aq) + NaOH(aq)  H2O(l) + NaCl(aq)

H+(aq) + OH-(aq)  H2O(l)


What volume of a 1.00M HCl solution is needed to neutralize 25.0 mL of a 0.350 M NaOH solution?

2.) Calculate moles of reactants:


What volume of a 1.00M HCl solution is needed to neutralize 25.0 mL of a 0.350 M NaOH solution?

3.) Determine limiting reactant:


What volume of a 1.00M HCl solution is needed to neutralize 25.0 mL of a 0.350 M NaOH solution?

4.) Calculate moles of H+ required:


What volume of a 1.00M HCl solution is needed to neutralize 25.0 mL of a 0.350 M NaOH solution?

5.) Calculate volume of HCl required:


Practice

  • What volume of a 0.101 M HNO3 is required to neutralize 1.21 L of 0.102 M KOH?


Acid-Base Reactions

  • Titration is a method for determining the concentration of a solution by reacting a known volume of that solution with a solution of known concentration.


Acid-Base Reactions

  • In a titration procedure, a measured volume of an acid or base of unknown concentration is placed in a beaker, and initial pH recorded.

  • A buret is filled with the titrating solution of known concentration, called a titrant.


Acid-Base Reactions

  • Measured volumes of the standard solution are added slowly and mixed into the solution in the beaker, and the pH is read and recorded after each addition. The process continues until the reaction reaches the equivalence point, which is the point at which moles of H+ ion from the acid equals moles of OH– ion from the base.


Acid-Base Reactions

  • An abrupt change in pH occurs at the equivalence point.


Acid-Base Reactions

  • Chemical dyes whose color are affected by acidic and basic solutions are called acid-base indicators.


Acid-Base Reactions

  • An end point is the point at which an indicator used in a titration changes color.

  • An indicator will change color at the equivalence point.


Buffered Solutions

  • The pH of blood must be kept in within a narrow range.

  • Buffersare solutions that resist changes in pH when limited amounts of acid or base are added.


Buffered Solutions

  • Ions and molecules in a buffer solution resist changes in pH by reacting with any hydrogen ions or hydroxide ions added to the buffered solution.

    HF(aq) H+(aq) + F–(aq)


Buffered Solutions

  • Additional H+ ions react with F– ions to form undissociated HF molecules but the pH changes little.

  • The amount of acid or base that a buffer solution can absorb without a significant change in pH is called the buffer capacity.


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