ACIDS AND BASESTopic 8

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

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

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

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

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

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?

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

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

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 SO4

2- + H3O+

• Sulfuric acid is only considered a strong acid in the first dissociation

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

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-

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

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

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

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

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

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

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

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