Acids & Bases

Acids:

·      taste sour

·      can dissolve some metals

·      usually produce H+ ions in water.

Bases:

·      taste bitter

·      feel “slippery” (like soap)

·      usually produce OH− ions in water.

 “Acids dissolve plants; bases dissolve people.”

neutralization: a reaction between an acid and a base, producing a salt and water.  For example:

HCl + NaOH → NaCl + H2O

acid + base   → salt + water

dissociation: the process of an acid or base splitting into ions.  For example:

HCl  H+ + Cl−
NaOH  Na+ + OH−      

protic: “having protons”.  Remember that H+ is a proton.  Acid-base chemists often use the word “proton” to mean an H+ ion.  Acids can be “monoprotic” (can dissociate to give one H+), “polyprotic” (more than one H+), “diprotic” (2 H+), etc. 

Three separate formal definitions of acids & bases:

1.   Arrhenius:  acids release H+ ions (actually form H3O+ ions) in water and bases release OH− ions in water.

2. Brønsted-Lowry:  acids donate H+ ions and bases accept H+ ions.  (The Brønsted-Lowry definition includes all Arrhenius acids & bases).

3. Lewis:  bases donate electrons and acids accept electrons.  (The Lewis definition includes all Brønsted-Lowry acids & bases.)

 


conjugates: the acid & base forms of a compound.  The acid form has an extra H+ that can dissociate.  The base form is the same thing without the H+.

conjugate base: the base formed by removing H+ from the acid.  For example, the conjugate base of HCl is Cl−.

conjugate acid: the acid formed by adding H+ to the base.  For example, the conjugate acid of NH3 is NH4+.

amphoteric: a substance that “can go either way”—i.e., it has both a conjugate acid and a conjugate base.  For example, the HSO4− ion is amphoteric:

H2SO4  HSO4−  SO42−

indicator: a substance whose conjugate acid and base forms have different colors.  E.g., the acid form of phenolphthalein is clear, but its conjugate base is pink.

pH: a measure of the strength of an acidic or basic solution.  Equal to −log [H+].  (The “p” in pH is a mathematical function that is literally equal to “−log”.)

      For example, if [H+] = 0.001 M, then pH = −log (0.001) = 3.

      Lower pH = more H+ ions = more acidic.  Higher pH = fewer H+ ions = more basic.

      In aqueous solutions:

·      pH range is from 1-14

·      pH 7 = neutral (neither acid nor base)

·      pH < 7 = acidic; pH > 7 = basic

pOH: similar to pH, but −log [OH−].

In aqueous solutions, pH + pOH = 14.

acid dissociation constant (Ka): is the equilibrium constant for the dissociation of an acid.  For the “generic” acid HA:

      The higher the Ka value, the stronger the acid.

pKa: −log Ka  (analogous to pH).  The pKa of an acid equals the pH at which the acid is exactly half dissociated.

base dissociation constant (Kb) is the equilibrium constant for the dissociation of a base.  For the “generic” base B:

water dissociation constant (Kw) is the equilibrium constant for the dissociation of H2O into H+ and OH−.  Kw = 1 × 10−14.

strong acid: an acid with a pKa lower than that of H3O+ (1.0).  Strong acids include HCl, HBr, HI, H2SO4 and HNO3.

      The conjugate base of a strong acid is a weak base.

      Strong acids dissociate completely into H+ and their (weak) conjugate bases.  The dissociated H+ converts H2O molecules to H3O+ ions.

strong base: a base with a pKa higher than that of OH− (14).  All hydroxides are strong bases.

      The conjugate acid of a strong base is a weak acid.

      Strong bases pull H+ off of H2O molecules to form their (weak) conjugate acids plus OH− ions.

weak acid: an acid with a pKa higher than that of H3O+.

weak base: a base with a pKa lower than that of OH−.

Weak acids & bases only partially dissociate in water.

buffer: an acid or base that prevents the pH of a solution from changing drastically until it neutralizes the ions of the buffer.  The effective pH range of a buffer is near its pKa.

      Weak acids & bases act as buffers in aqueous solutions.  (For strong acids & bases, the water acts as the buffer.)