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Acid/Base

Acid/Base. Chapter 14 . Arrhenius Acids . Arrhenius Acid – a chemical compound that increases the concentration of hydrogen ions in aqueous solution. Strong acids – one that ionizes completely in aqueous solution. Ex: HCl (g) + HOH (l)  H + (aq) + Cl - (aq)

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Acid/Base

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  1. Acid/Base Chapter 14

  2. Arrhenius Acids • Arrhenius Acid – a chemical compound that increases the concentration of hydrogen ions in aqueous solution. • Strong acids – one that ionizes completely in aqueous solution. • Ex: HCl(g) + HOH(l) H+(aq) + Cl-(aq) • Weak acids – weak electrolytes • Ex: HCN(g) + HOH(l)<-> H3O+(aq) + CN-(aq)

  3. Arrhenius Bases • Arrhenius Base – a substance that increases the concentration of hydroxide ions in aqueous solution. • Strong bases – one that dissociates completely in aqueous solution. • Ex: KOH(s) + HOH(l) K+(aq) + OH-(aq) • Weak bases – relatively low concentration of hydroxide ions. • Ex: NH3(g) + HOH(l) <-> NH+4(aq) + OH-(aq)

  4. Bronsted-Lowry Acids • Bronsted-Lowry acid – a molecule or ion that is a proton donor. • Ex: HCl + NH3 NH+4 + Cl- • HCl – proton donor • NH3 – proton acceptor • Bronsted-Lowry base – a molecule or ion that is a proton acceptor.

  5. Bronsted-Lowry acid-base reaction • Protons are transferred from one reactant (the acid) to another (the base) • HCl(aq) + NH3(aq) NH4Cl(g)

  6. Monoprotic and Polyprotic Acids • Monoprotic acid – an acid that can donate only one proton per molecule. • Ex: HCl, HNO3 … • Ex: HCl(g) + H20 (l) H30+(aq) + Cl-(aq) • Polyprotic acid – an acid that can donate more than one proton per molecule. • Ex: H2SO4 , H3PO4 … • E1: H2SO4(l) + H2O(l) H3O+(aq) + HSO4-(aq) • E2: HSO4-(aq) + H2O(l) <-> H3O+(aq) + SO42-(aq)

  7. Conjugate Acids and Bases • Conjugate base – the species that remains after a bronsted-lowry acid has given up a proton is the conjugate base of that acid. • Conjugate acid – the species that is formed when a bronsted-lowry base gains a proton is the conjugate acid of that base. • Ex: HCl + HOH  H3O+ + Cl- Acid Base C.Acid C.Base

  8. Amphoteric Compounds • Any species that can react as either an acid or a base . • Ex: water • Sulfuric acid and water -> hydronium ion and hydrogen sulfate ion * water acts as a base H2SO4 + HOH -> H3O+ + HSO4- Ammonia and water  ammonium ion and hydroxide ion * water acts as a acid NH3 + HOH <-> NH4+ + OH-

  9. Neutralization Reactions • Reaction of hydronium ions and hydroxide ions to form water molecules. • A salt is an ionic compound composed of a cation from a base and an anion from an acid. • HCl(aq) + NaOH  NaCl(aq) + H2O(l)

  10. Lewis Acids and Bases • Lewis acid – an atom, ion, or molecule that accepts an electron pair to form a covalent bond. • Lewis base – an atom, ion, or molecule that donates an electron pair to form a covalent bond. • Ex: Boron trifluoride and Ammonia Lewis acid Lewis base

  11. Lewis acid-base reaction • The formation of one or more covalent bonds between an electron-pair donor and an electron-pair acceptor. • NH3(g) + BF3(g) H3N:BF3(g)

  12. Acid-Base Systems

  13. Ammonia • Arrhenius base because OH- ions are created when in solution. • NH3(g) + HOH(l) NH4+(aq) + OH-(aq) • Bronsted-lowry base because is accepts a proton in an acid-base reaction. • NH3(g) + HCl(aq) NH4+(aq) + Cl-(aq) • Lewis base in all reactions in which it donates its lone pair to form a covalent bond. • NH3(g) + BF3(g) H3N:BF3(g)

  14. Self-ionization of Water • Two water molecules produce a hydronium ion and a hydroxide ion by transfer of a proton. • Ex: HOH + HOH <-> H3O+ + OH- • Conductivity measurements show that concentrations of H3O+ and OH- in pure water are each only 1.0 x 10-7 mol/L of water at 250C. • Ionization constant of water • Kw = [H3O+][OH-] • Kw = (1.0 x 10-7M)(1.0 x 10-7M) • Kw = 1.0 x 10-14M2

  15. What you need to know for solving pH problems • Kw = [H3O+] [OH-] • pH = - log [H3O+] • pOH = - log [OH-] • pH + pOH = 14 • [H3O+] = antilog (- pH) • [OH-] = antilog (-pOH)

  16. Titration • Controlled addition and measurement of the amount of a solution of known concentration required to react completely with a measured amount of a solution of unknown concentration. • Provides a sensitive means of determing the chemically equivalent volumes of acidic and basic solutions.

  17. Equivalence Point • The point at which two solutions used in a titration are present in chemically equivalent amounts. • Acid-base indicators: compounds whose colors are sensitive to pH • Indicators that undergo transition (pH range over which an indicator changes color) at about pH 7 are used to determine the equivalence point of strong-acid/strong-base titrations because the neutralization of strong acids with strong bases produce a salt solution with a pH of 7 (bromthymol blue) • Indicators that change color at pH lower than 7 are useful in determining the equivalence point of strong-acid/weak-base titrations. The equivalence point of a strong-acid/weak-base titration is acidic because the salt formed is itself a weak acid. Thus the salt solution has a pH lower than 7. (methyl orange) • Indicators that change color at pH higher than 7 are useful in determining the equivalence point of weak-acid/strong-base titrations. These reactions produce salt solutions whose pH is greater than 7. This occurs because the salt formed is a weak base. (phenolphthalein) • What type of indicator is used to determine the equivalence point of weak-acid/weak-base titration? • None at all. Dependant on the relative strengths of the reactants.

  18. Standard Solution • The solution that contains the precisely known concentration of a solute.

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