Acids and Bases

• Acid is a substance that donates protons (H⁺) in an aqueous solution
• Base is a substance that accepts protons (H⁺) in an aqueous solution

Strong Acids

• Completely dissociate or ionize and release all H⁺ ions in solution
• Strong acids are good electrical conductors due to the high concentration of ions, where more ions result in better conductivity

Weak Acids

• Partially dissociate or ionize, resulting in fewer H⁺ ions released
• Weak acids are bad electrical conductors because there are fewer ions in solution

Alkalis

• Soluble bases that produce OH⁻ ions in solution

pH

• A measure of the concentration of H⁺ ions
• All alkalis are bases, but not all bases are alkalis
• pH = -log[H⁺]
• [H⁺] = 10^(-pH)
• Kw = [H⁺][OH⁻] = 1.0 x 10⁻¹⁴
• Concentration is Moles/Volume

Concentrated vs Dilute

• Refers to the number of atom molecules in a given volume
• High concentration means high molarity, while dilute means low molarity
• Solutions can have strong dilute acid or weak concentrated acid

Acid Deposition

• Precipitation with any acidic components

Arrhenius Theory (Limited)

• A base is a substance releasing hydroxide ions (OH⁻) in an aqueous solution, neutralizing acids

Bronsted-Lowry Acids and Bases

• Requires a reversible reaction
• Acid is a proton (H⁺) donor in aqueous solutions
• Base is a proton acceptor in aqueous solutions

Bronsted-Lowry Acid

• Donates a proton

Conjugate Base

• What remains after a Bronsted-Lowry acid donates a proton

Bronsted-Lowry Base

• Accepts a proton

Conjugate Acid

• The new substance formed when a Bronsted-Lowry base accepts a proton

• Metal hydroxides do not act as Bronsted-Lowry bases because they cannot accept protons.

• Hydroxide ions in a solution can act as a Bronsted-Lowry base by accepting a proton

Conjugate Strength

• Strong Bronsted-Lowry acids result in weak conjugate bases

• Strong Bronsted-Lowry bases result in weak conjugate acids

• Weak Bronsted-Lowry acids result in strong conjugate bases

• Weak Bronsted-Lowry bases result in strong conjugate acids

• All Bronsted-Lowry acids contain a proton (H⁺)

  • Charge can be 0, +, or -

• All Bronsted-Lowry bases contain a lone electron pair, or a dative covalent/coordinate bond

• Charge can only be 0 or +

Amphiprotic Species

• A molecule/ion can either donate or accept a proton (H⁺)

• Must have a net negative charge to accept H⁺

• Amphiprotic molecules must also be amphoteric

• Metallic oxides (except aluminum oxide Al₂O₃) have Hydrogen

• Oxides transition from basic to amphoteric to acidic across a period, with increasing acidic/covalent character

Amphoteric Reactions

• Amphoteric Oxide + Acid/Alkali produces Salt + Water

Determining Acid Strength

• Use a pH meter, probe, or strip

• Check electrical conductivity with a conductivity meter

  • Strong acids/bases dissociate fully, more ions result in better electrical conductivity

• Measure rate reactions

  • React strong/weak acids with metal, metal carbonates, and metal hydrogen carbonates to produce gas, collect gas with syringe for Rate of Reaction

• Use Titration

• Metal hydroxides are strong bases (e.g., NaOH)

• Weaker acids/bases needs additional energy to ionize, its less exothermic.

pH Scale

• Logarithmic, every integer pH value increase/decrease changes the H+ concentration by tenfold
  • Solution pH 3 has ten times more H+ ions than a solution at pH 4
  • pH 3 → [H+] x 10⁻³ moldm⁻³, pH 5 → [H+] x 10⁻⁵ moldm⁻³
• To dilute 100cm³ of a pH 6 solution to pH 7 requires adding 900cm³ of water

Measuring pH

• Use UI solution/paper
• Color readings are subjective
• Use a pH meter

Indicators

• Universal indicator is red in acid, blue in alkali

• Methyl orange is red in acid, yellow in alkali

• Red/blue litmus paper is red in acid, blue in alkali, and no change otherwise

• Phenolphthalein is colorless in acid, pink in alkali

• pH = -log[H+]

• [H+] = 10^(-pH)

• [Strong Monoprotic Acid/Base] = [H+]

• [Diprotic Acid/Base] = 2[H+]

• [Triprotic Acid/Base] = 3[H+]

Ion Product of Water

• Kw = [H+][OH⁻]
• At 298 Kelvin: Kw = [H+][OH⁻] = 1.0 x 10⁻¹⁴
• In pure water, [H+] = [OH⁻]

Self-Ionization of Water

• Water undergoes self-ionization:

• H₂O ⇔ H+ + OH-

• H₂O + H₂O ⇔ H₃O+ + OH-

• The concentrations of H+ and OH- are small because the equilibrium lies on the left [𝐻+][𝑂𝐻−]

• K = {𝐻₂𝑂]

• Kw depends on temperature, at 298 Kelvin, Kw = [H+][OH-] = 1.0 x 10⁻¹⁴

• H+ + OH- ⇔ H₂O (neutralization)

• Forward reaction is exothermic (bond making)

• Reverse reaction is endothermic (bond breaking)

• Water at higher temperatures has lower pH

• Because pH = -log[H+]

• At higher temperatures, the reaction favors the endothermic side, thus, resulting in more H⁺ and OH⁻

• Higher H⁺ value results in a lower pH value

Determining pH From H⁺ and OH⁻ ions

• [H+] = [OH-] → neutral
• [H+] > [OH-] → acidic
• [H+] < [OH-] → basic

Neutralization Reactions

• General form: HCl + NaOH → NaCl + H₂O

• Sulfuric acid becomes sulfate

• Hydrochloric acid becomes chloride

• Phosphoric acid becomes phosphate

• Ethanoic acid becomes ethanoate

• Nitric acid becomes nitrate

• Acid + Metal Oxide produces Salt + Water

• Acid + Metal Hydroxide produces Salt + Water

• Acid + Alkali produces Salt + Water

• Alkalis release OH ions

• Acid + Metal Carbonate produces Salt + Carbon Dioxide + Water

• Acid + Metal Hydrogen Carbonate produces Salt + Carbon Dioxide + Water

Soluble Carbonates & Hydrogen Carbonates

• NH₄CO₃
• Na₂CO₃
• NaHCO₃
• KCO₃
• KHCO₃
• Ca(HCO₃)₂

Soluble Hydroxides

• NH₄OH
• LiOH
• NaOH
• KOH

Other Acid-Base Reactions

• Ammonia (NH₃) is a gas dissolving in water to make an alkali (ammonium hydroxide)
• Ammonia solution (NH₄OH) is a soluble base reacting with acids to form water and salt
  • HNO₃ + NH₄OH → NH₄NO₃ + H₂O

Amines

• Amines are bases by accepting a proton (H⁺)
• Amines are weak bases compared to inorganic bases like NaOH
• Primary, secondary, and tertiary amines influence reaction (number of Hydrogens attached to Nitrogen)
• Amine + Acid produces Ammonium Salt

Standard Enthalpy Change of Neutralization

• Enthalpy change when 1 mole of water forms from neutralization under standard conditions (298K, 100kPa)
• ΔHn = -57.3 kJmol⁻¹
• Exothermic
• Enthalpy change of neutralization for any strong acid and strong base reaction are almost identical
  • Always -57.3 kJmol⁻¹ for strong acids/bases
• Enthalpy change of neutralization of weak acids/bases is less than strong ones
• Ionization of weak acids/bases are mildly endothermic
• Energy is used to ionize weak acids/bases, energy releases during the formation of water
• Weaker acids/base's enthalpy of neutralization has a smaller net difference

Titration

• Process of finding the unknown concentration of a solution

  • Titrant describes a solution with a known concentration
  • Titrand / Analyte describes a solution with an unknown concentration

• Equivalence point indicates the point of neutralization

• End point indicates the point of indicator color change

• Polyprotic acids has more than one proton (H+)

• Back titration is same as typical titration, but it proceeds in reverse direction

• First, add excessive amount of known reagent to unknown, and allow reaction

• Known amount of another reagent is added to react with the excess reagent

• Remaining excess reagent is titrated with a known solution.

• Basic titration follows the same procedure as acid titration, but a base serves as unknown titrant

pH Curves

• Below Neutralization Point (Acidic) Moles = concentration x volume

• Above Neutralization Point (Basic)

• During heat calculation with thermometric method assume

• Acid solution density matches water density

• Acid-base solution specific heat capacity matches water

• No evaporation occurs

• pH measures Hydrogen ion concentration of solution

Strong Acids

• Fully dissociates/ionizes
• HCl-
• Sulfuric acid

Strong Bases

• Fully dissociates/ionizes
• Sodium hydroxide
• Potassium hydroxide
• Lithium hydroxide
• Acids react with reactive metals and produce hydrogen
• Acid deposition encompasses all acidic component precipitation
• Rainwater mixed with dissolved CO₂ is weakly acidic (pH 5.6)
• Gaseous non-metal oxides results in acid rain

Nitrogen Oxide

• N₂ + O₂ → 2NO
• Nitrogen (II) oxide reacts with oxygen to form nitrogen dioxide:
  • 2NO + O₂ → 2NO₂
  • 2NO₂ + H₂O → HNO₂ + HNO₃
• Nitric (III) acid reacts with more oxygen to produce nitric (IV) acid:
  • 2HNO₃ + O₂ → 2HNO₃
• Nitrogen dioxide is a red-brown gas