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Acids and Bases

Acids and Bases. Topic 8. 8.1 Reactions of acids and bases. Acids with metals Produces a salt and hydrogen gas Mg + 2HCl  MgCl 2 + H 2 Acids with carbonates and hydrogencarbonates Produces salt + carbon dioxide + water Na 2 CO 3 + H 2 SO 4  Na 2 SO 4 + H 2 O + CO 2. Con’t.

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Acids and Bases

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  1. Acids and Bases Topic 8

  2. 8.1 Reactions of acids and bases • Acids with metals • Produces a salt and hydrogen gas • Mg + 2HCl  MgCl2 + H2 • Acids with carbonates and hydrogencarbonates • Produces salt + carbon dioxide + water • Na2CO3 + H2SO4 Na2SO4 + H2O + CO2

  3. Con’t • Acids with bases and alkalis • Bases are metal oxides • Produce a salt and water • CuO + H2SO4 CuSO4 + H2O • Alkalis are bases that dissolve in water • Produce a salt and water • NaOH + HNO3 NaNO3 + H2O

  4. 8.2 Definitions of acids and bases • BrØnsted-Lowry definitions • Acid proton (H+) donor • Base/ alkal proton (H+) acceptor • Conjugate base the base formed when an acid reacts and donates a proton to become a base • Conjugate acid the acid formed when a base reacts and accepts a proton to become an acid • Referred to as conjugate acid- base pair

  5. Conjugate acid-base pairs • CH3COOH + H2O  CH3COO- + H3O+ • Which is the beginning acid? Base? • Which is the acid’s conjugate? • The base’s? • Water is called amphoteric. What is that? Can act as an acid or base

  6. Another way to phrase it… • In the forward reaction the CH3COOH acts as the acid and the H2O acts as the base • In the reverse reaction the CH3COO- acts as the base and the H3O+ acts as the acid

  7. Lewis theory of acids and bases • Acid electron pair acceptor • Base electron pair donor • Must understand the Lewis structure of the compound to know which substance will accept the electrons • Ex. NH3 + H+ NH4+ • Which substance gained electrons? Which donated?

  8. Con’t • A dative covalent bond is always formed in a Lewis acid-base reaction • What is a dative covalent bond? Both electrons come from the same atom • For a substance to act as a Lewis base, it must have a lone pair of electrons • For a substance to act as a Lewis acid, it must have space to accept a pair of electrons

  9. 8.3 Strong and weak acids and bases • When acid reacts with water it dissociates or ionizes • Can use the Bronsted-Lowry theory to understand this • Strong acids completely dissociate in aqueous solution • Which direction does the equilibrium dominantly lie? To the right (products) • HA  H+ + A- • Uses a non-reversible arrow

  10. Strong acids • HCl is considered a monoprotic acid it dissociates to form one proton per molecule • H2SO4 is considered diprotic dissociates to form two protons per molecule • H2SO4 + H2O  HSO4- + H3O+ • HSO4- + H2O  SO42- + H3O+ • Sulfuric acid is only considered a strong acid in the first dissociation

  11. Weak acids • Only partially dissociate in aqueous solution • The equilibrium arrow is used for these equations • HA  H+ + A- • Ex. Carbonated water is acidic due to dissolved CO2, which acts as a weak acid

  12. Bases • Strong bases ionize completely in aqueous solution • Ex. NaOH Na+ + OH- • The group 1 hydroxides are strong bases; along with Ba(OH)2 • Weak bases ionize partially in aqueous solution • Equilibrium arrows are used in these equations • Ex: NH3+ H2O NH4+ + OH-

  13. Distinguishing experimentally between strong and weak acids and bases • Solutions of strong acids conduct electricity better than weak acids • Why? There is a large concentration of ions to carry the electrical current • Can also be called strong electrolytes or weak electrolytes • The same concept is true for strong and weak bases

  14. Con’t • Strong acids have a lower pH than weak acids • What does pH measure? The concentration of H+ ions in solution • Lower pH = more H+ ions • Would the pH for strong bases be higher or lower? Higher • Why? There are very few H+ ions in the solution of strong bases

  15. Con’t • Strong acids react more violently with metals or carbonates • The higher concentration of free H+ ions cause a more rapid reaction with metal to form H2(g) • There is a similar effect when a carbonate is added

  16. Con’t • strength vs. concentration • Concentration refers to the number of moles of acid in a certain volume (i.e. mol dm-3) • Strength refers to what? How much the acid dissociates • Ex. Ethanoic acid is considered a weak acid. No matter how concentrated the acid solution is, it will still not fully dissociate. • Similar for bases

  17. 8.4 pH • Definition: pH is the negative logarithm to base 10 for the hydrogen ion concentration in aqueous solution pH= -log10 [H+(aq)] • The pH scale is used to indicate how acidic or alkaline a solution is • The scale is from 1 to 14 • One being the most acidic • Fourteen is the most alkaline • Seven is neutral

  18. pH • Since pH is on a log scale, a 1 unit change in pH means there is a tenfold change in H+ ion concentration • Calculating [H+] from pH [H+]= 10-pH • This is the inverse of the previous equation

  19. Calculating pH of a strong acid • It can be assumed that a strong acid fully dissociates and the [H+] is equal to [acid] • Ex: calculate the pH of a 0.00150 mol dm-3 solution of HCl. • pH=-log10[H+]= -log[0.00150]= 2.82 • Just plug in the [acid] for hydronium ions

  20. pH is not a measure of acid strength • This is the measure of what? [H+] ions • It is possible for a dilute solution of a strong acid to have a higher pH than a concentrated solution of a weak acid • pH can be used to compare the strength of acids,ONLY IF THE CONCENTRATIONS OF THE ACIDS ARE EQUAL

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