Identifying Acids and Bases Quiz

Identifying Acids and Bases

• Acids typically contain hydrogen at the beginning of their chemical formula.
• Bases usually contain a hydroxide ion (OH-).
• A hydrogen attached to a nonmetal is generally an acid, while a hydrogen attached to a metal is usually a base.
• Acids are usually positively charged, while bases are negatively charged.

Arrhenius Definition

• An Arrhenius acid releases H+ ions in solution.
• An Arrhenius base releases OH- ions in solution.
• H+ ions in water are bonded to water molecules and exist as hydronium ions (H3O+).

Brønsted-Lowry Definition

• A Brønsted-Lowry acid is a proton (H+) donor.
• A Brønsted-Lowry base is a proton acceptor.
• In a reaction between an acid and a base, the acid forms a conjugate base (by losing a proton), and the base forms a conjugate acid (by gaining a proton).

Writing Conjugate Acids and Bases

• To find the conjugate acid, add an H+ to the molecule and increase the charge by 1.
• To find the conjugate base, remove an H+ from the molecule and decrease the charge by 1.

pH Scale

• The pH scale ranges from 0 to 14, with the possibility of going beyond those limits.
• pH 7 is neutral.
• pH below 7 is acidic.
• pH above 7 is basic.
• pH = -log[H3O+].
• pOH = -log[OH-].
• pH + pOH = 14 at 25°C.

Strong Acids and Bases

• Strong acids ionize completely in solution, forming strong electrolytes.
• Strong bases are soluble ionic compounds that ionize completely in solution, forming strong electrolytes.
• Strong acids form strong electrolytes; weak acids form weak electrolytes.
• Strong bases form strong electrolytes; weak bases form weak electrolytes.

Common Strong Acids

• HCl
• HBr
• HI
• HNO3 (nitric acid)
• H2SO4 (sulfuric acid)
• HClO4 (perchloric acid)

Weak Acids

• HF
• NH4+ (ammonium ion)
• CH3COOH (acetic acid)

Strong Bases

• KOH (potassium hydroxide)
• NaOH (sodium hydroxide)
• Ba(OH)2 (barium hydroxide)

Weak Bases

• NH3 (ammonia)
• F- (fluoride)
• NO2- (nitrite)
• CH3COO- (acetate)
• CN- (cyanide)
• HSO3- (bisulfite)

Strong Bases Other than Hydroxides

• O2- (oxide)
• H- (hydride)

Properties of Acids and Bases

• Acids taste sour.
• Bases taste bitter.
• Acids turn blue litmus paper red.
• Bases turn red litmus paper blue.
• Acids react with active metals to produce hydrogen gas.

Acid Dissociation Constant (Ka)

• Ka = [H3O+][A-] / [HA] (where HA is the acid)
• A higher Ka value indicates a stronger acid.
• pKa = -log(Ka).
• Strong acids have large Ka values and small pKa values.
• Weak acids have small Ka values and large pKa values.

Base Dissociation Constant (Kb)

• Kb = [OH-][BH+]/[B] (where B is the base)
• pKb = -log(Kb)

Autoionization of Water

• Water can act as both an acid and a base.
• Kw = [H3O+][OH-] = 1 x 10^-14 at 25°C.
• As temperature increases, Kw increases.

Amphoteric Substances

• Amphoteric substances can act as both acids and bases.
• Water is an amphoteric substance.
• H2PO4- is also an amphoteric substance.
• The conjugate base of a weak acid is a weak base.

Reactions between Acids and Bases

• Strong acids ionize completely in solution, so their reactions are represented with single arrows.
• Weak acids ionize partially in solution, so their reactions are represented with double arrows indicating a reversible process.
• Strong bases ionize completely in solution, forming strong electrolytes.
• Weak bases ionize partially in solution, forming weak electrolytes.

Mechanisms of Acid-Base Reactions

• The mechanism of reactions between acids and bases involves the transfer of protons (H+) from the acid to the base.
• The lone pairs of electrons on the base can be used to attack the hydrogen atom of the acid, breaking the bond and forming a new bond with the base.

Summary of Definitions for Acids

• Arrhenius Acid: Releases H+ ions in solution.
• Brønsted-Lowry Acid: Proton donor.
• Lewis Acid: Electron pair acceptor.

Summary of Definitions for Bases

• Arrhenius Base: Releases OH- ions in solution.
• Brønsted-Lowry Base: Proton acceptor.
• Lewis Base: Electron pair donor.

pH and pOH

• At 25°C, the product of the concentrations of hydronium ions (H3O+) and hydroxide ions (OH-) is equal to 1 x 10^-14.
• This is represented by the equation: [H3O+] x [OH-] = Kw = 1 x 10^-14
• The pH and pOH of a solution at 25°C are related by the equation: pH + pOH = 14
• pH is the negative logarithm of the hydronium ion concentration: pH = -log[H3O+]
• pOH is the negative logarithm of the hydroxide ion concentration: pOH = -log[OH-]

Ka and Kb

• Ka (acid dissociation constant) measures the strength of an acid.
• Kb (base dissociation constant) measures the strength of a base.
• The product of Ka and Kb for a conjugate acid-base pair is equal to Kw: Ka x Kb = Kw
• pKa and pKb are related to Ka and Kb by the equations: pKa = -log(Ka) and pKb = -log(Kb).
• At 25°C, pKa + pKb = 14.

Acid-Base Definitions

• Arrhenius Acid: An acid releases H+ ions in solution.
• Arrhenius Base: A base releases OH- ions in solution.
• Brønsted-Lowry Acid: A proton (H+) donor.
• Brønsted-Lowry Base: A proton (H+) acceptor.
• Lewis Acid: An electron pair acceptor.
• Lewis Base: An electron pair donor.

Strong vs. Weak Acids and Bases

• Strong acids: Dissociate completely in solution, resulting in a high concentration of H+ ions and a low pH.
• Weak acids: Dissociate partially in solution, resulting in a lower concentration of H+ ions and a higher pH compared to strong acids with the same concentration.
• Strong bases: Dissociate completely in solution, resulting in a high concentration of OH- ions and a high pH.
• Weak bases: Dissociate partially in solution, resulting in a lower concentration of OH- ions and a lower pH compared to strong bases with the same concentration.

Relationship between Ka, pKa, Acid Strength

• A higher Ka value indicates a stronger acid.
• A lower pKa value indicates a stronger acid.
• The stronger the acid, the weaker its conjugate base.

Examples

• Hydrochloric acid (HCl) is a strong acid.
• Acetic acid (CH3COOH) is a weak acid.
• Sodium hydroxide (NaOH) is a strong base.
• Ammonia (NH3) is a weak base.

pH Scale

• The pH scale ranges from 0 to 14.
• pH less than 7 indicates an acidic solution.
• pH greater than 7 indicates a basic solution.
• pH of 7 indicates a neutral solution.

Metal Ions as Lewis Acids

• Metal cations with a high positive charge can act as Lewis acids, accepting electron pairs from water molecules.
• This can lead to the formation of H+ ions, making the solution acidic.
• For example, aluminum chloride (AlCl3) dissolved in water creates an acidic solution.

Test your understanding of acids and bases, their definitions, and their classifications. This quiz covers key concepts, including the Arrhenius and Brønsted-Lowry definitions, as well as how to write conjugate acids and bases. Perfect for chemistry students looking to refine their skills!