Relative Strengths of Acids The relative strength of various acids can be compared using Ka. The larger the value of Ka, the stronger the acid The structure of an acid plays an important role in determining the strength of an acid.

Relative Strengths of Acids For an acid with the general formula, HX, the strength of the acid depends on: the polarity of the H - X bond H X the strength of the H - X bond the stability of the conjugate base, X-

Relative Strengths of Acids Using information about these properties, several trends related to the relative strengths of acids can be identified for binary acids and oxyacids. Binary acid: an acid with the general formula, HX, which is generally composed of two elements (one of which is hydrogen) Oxyacid: an acid that contains oxygen

Relative Strengths of Acids Relative Strengths of Binary Acids: Within a group, the strength of an acid increases moving down the group HCl is stronger than HF Within the same period, the strength increases as the electronegativity of the element X increases (i.e. left to right) HCl is stronger than H2S

Relative Strengths of Acids For binary acids: Periodic Table Increasing acid strength Increasing acid strength

Relative Strengths of Acids Relative Strengths of Oxyacids: For oxyacids with the same central atom, acid strength increases as the number of oxygen atoms attached to the central atom increases. HClO3 is stronger than HClO

Relative Strengths of Acids Relative Strengths of Oxyacids: For oxyacids with the same number of O atoms, acid strength increases as the electronegativity of the central atom increases HClO is stronger than HBrO In general, electronegativity increases toward the top within a group and from the left toward the halogens within a period.

Relative Strengths of Acids For oxyacids with the same # of oxygens: Halogens Periodic Table Increasing acid strength Increasing acid strength

Relative Strengths of Acids Example: Identify the stronger acid in each of the following pairs: HClO2 vs. HIO2 H2SeO3 vs. H2SeO4 H2SeO3 vs. HBrO3 H2O vs. HF

Buffers Before measuring the pH of an aqueous solution using a pH meter, chemists must standardize the pH meter. Adjust the reading of the pH meter to the correct value using standard buffers. Buffer: A solution that resists a change in pH when small amounts of acid or base are added

Buffers Buffers resist changes in pH because they contain both: An acid Neutralizes any OH- ions added A base Neutralizes any H+ ions added At the same time, the acidic and basic components of a buffer must not react with each other. Generally use a weak conjugate acid-base pair

Buffers Examples of buffers: NaC2H3O2 + HC2H3O2 NH4Cl + NH3 Lactic acid + sodium lactate Two important properties of a buffer: pH Buffer capacity

Buffers Buffer capacity the amount of acid or base a buffer can neutralize before the pH begins to change appreciably The buffer capacity depends on the amount of acid and base from which the buffer is prepared.

Henderson- Hasselbalch Equation Buffers The pH of a buffer depends on the pKa of the acid and on the relative concentrations of the acid and base in the buffer. pH = pKa + log10 [base] [acid] pKa = - log Ka Henderson- Hasselbalch Equation

Buffers OR The Henderson-Hasselbalch equation can be used to: determine the pH of a particular buffer determine the ratio of base to acid needed to prepare a buffer with a specific pH. OR

Buffers Example: What is the pH of a buffer containing 0.12 M benzoic acid and 0.20 M sodium benzoate? Ka = 6.3 x 10-5.

Buffers Example: What base to acid ratio is needed to make a pH 4.95 buffer using benzoic acid and sodium benzoate? Ka = 6.3 x 10-5.

Acid-Base Titrations An acid-base titration can be used to determine the concentration of an acid or base solution titration: a technique for determining the concentration of an unknown solution using a standard solution a solution with a known concentration

Acid-Base Titrations The equivalence point in a titration can be determined using either a pH indicator (ex. phenolphthalein) or a pH meter. Equivalence point: the point in the titration where stoichiometrically equivalent amounts of base have been added to the acid (or vice versa) the base added has completely reacted with all available protons (H+)

Region of high buffer capacity Acid-Base Titrations The general shape for a titration curve for a monoprotic acid: pH Equivalence point Region of high buffer capacity pKa Half neutralization point mL base added

Acid-Base Titrations General Shape for a Diprotic Acid Titration Curve:

Acid-Base Titrations General Shape for a Triprotic Acid Titration Curve:

Titration of Polyprotic Acids Example: The titration curve on the next slide was obtained by titrating 0.250 g of a polyprotic acid with 0.100 M NaOH. Answer the following questions. How many equivalence points are there? Is the acid monoprotic, diprotic, or triprotic? What are the values for each pKa? What volume of NaOH was needed to neutralize the acid? What is the molar mass of the unknown acid?

Titration of Polyprotic acids

Lewis Acids and Bases In order for a substance to be a Bronsted-Lowry base (i.e. a proton acceptor), it must have an unshared pair of electrons to bind the proton. Similarly, a Bronsted-Lowry acid (i.e. a H+ ion) always gains a pair of electrons during an acid-base reaction.

Lewis Acids and Bases Lewis noticed this trend and proposed a new definition of acids and bases: Lewis Acid: An electron pair acceptor Lewis Base: An electron pair donor

Lewis Acids and Bases Examples of Lewis Acids: H+ BF3 Fe3+ CO2 Examples of Lewis Bases: OH- CH3NH2

Lewis Acids and Bases Example: Identify the Lewis acid and the Lewis base in the following reactions.