Acids & Bases

1. Acids & Bases

2. Acid & Base Properties Acid: • Sour taste (but never do this in lab!) • Is an electrolyte • Changes the color of an indicator • Reacts with active metals producing hydrogen gas (H2) • Reacts with bases to produce an ionic salt and water

3. Base • Bitter taste (ahh, shouldn’t have to mention since it is bitter, but don’t actually taste in the lab!) • Base solutions have a slippery feel. Probably not a good idea to go around touching chemicals though. • Changes the color of an indicator • Is an electrolyte • Reacts with acids to produce an ionic salt and water.

4. Acid Base Definitions Arrhenius Definitions: • Acid = A chemical that donates a H+ ion into solution. • Base = A chemical that donates a OH− ion into solution. • Examples: HCl is an acid because it will donate the hydrogen as an ion into solution. KOH is a base because it will donate the hydroxide ion into solution. • This is the least correct definition, but is the easiest to understand, and therefore the 1 we will use most

5. Brönstad—Lowry Definitions: • Acid = A chemical that donates a proton into solution. A proton is the same as H+ ion! • Base = A chemical that removes a proton from solution. Which explains OH−, as it will accept a proton to become H2O. • This definition is better than Arrhenius because it explains why many other chemicals other than hydroxides behave as bases. Still not the best though!!

6. Lewis Definitions: • Acid = A chemical that is an electron pair acceptor. • Base = A chemical that is an electron pair donor. • This is the best, most comprehensive explanation of acids and bases we have. It explains all chemicals that behave as acids & bases. • This is the most difficult to understand however, so we will not use it to describe the chemicals we will work with in class. The chemicals we will work with in class will be easily explained using the Arrhenius definition.

7. Strengths of Acids & Bases • Acids & Bases are electrolytes • Electrolytes can be strong or weak depending on how well they dissolve in water. • Chemicals that completely dissolve into ions and don’t recombine easily are strong electrolytes. Acids & Bases that completely dissolve are strong acids & strong bases. • Chemicals that partially dissolve into ions and tend to recombine easily are weak electrolytes. Acids & Bases that partially dissolve are weak acids & weak bases.

8. Ex. of strong acids: • HClHCl H+ + Cl− • H2SO4H2SO4  H+ + HSO4− • HNO3HNO3  H+ + NO3− • HClO4HClO4  H+ + ClO4− • HBrHBr  H+ + Br− • Ex. of weak acids: • HC2H3O2HC2H3O2⇄ H+ + C2H3O2− • H3PO4H3PO4⇄ H+ + H2PO4− • H2O H2O⇄ H+ + OH− • H2CO3H2CO3⇄ H+ + HCO3−

9. Ex. Of Strong Bases: • NaOHNaOH Na+ + OH− • KOH KOH K+ + OH− • Ca(OH)2 Ca(OH)2 Ca+2 + 2 OH− • RbOHRbOH Rb+ + OH− • Ex. Of Weak Bases: • Al(OH)3 Al(OH)3⇄ Al+3 + 3 OH− • NH4OH NH4OH⇄ NH4+ + OH− • H2O H2O⇄ H+ + OH− • F− F− + H2O ⇄ HF + OH− *Note all weak acids & bases are reversible reactions!

10. Acid Terms • Monoprotic Acid: Acid with only one H to donate into solution. Ex. HCl • Diprotic Acid: Acid with 2 H to donate into solution. Ex. H2SO4 • Triprotic Acid: Acid with 3 H to donate into solution. Ex. H3PO4 • Polyprotic Acid: Acid with 2 or more H to donate into solution. • Amphiprotic: Chemical that can behave as both an acid & a base. • Oxyacid: Acid containing a polyatomic ion with O • Protic term comes from H+ ions being protons.

11. Conjugate Acid Base Pairs • The formation of an acid or a base involves the formation of a conjugate acid or conjugate base. • The equation for the formation of an acid or base will contain the following; an acid, a base, a conjugate acid, and a conjugate base. • Ex. Formation of Hydrochloric acid HCl: • HCl + H2O  H3O+ + Cl− • HCl is the acid, H2O is the base • H3O+ is the conjugate acid, Cl− is the conjugate base

12. With a strong acid or strong base, the acid / base combo is on the left side of the arrow, & the conjugate acid / conjugate base combo is on the right side of the arrow. • With weak acids or weak bases, it could be more tricky, but we treat them the same. A & B on the left, Conj. A & Conj. B on the right. • Ex. Ammonia NH3 • NH3 + H2O ⇄ NH4+ + OH− • NH3 is the Base, H2O is the acid • NH4+ is the conjugate acid, OH− is the conjugate base

13. One last example: HSO4− • HSO4− + H2O ⇄ H3O+ + SO4−2 • Acid: • Base: • Conjugate Acid: • Conjugate Base: • You can treat the reactions as reversible, to determine which is the acid and which is the base. • See which one loses the H, that’s the acid. • See which one gains the H, that’s the base.

14. Neutralization Reactions • One of the properties of acids & bases was that when they react together, they will form an ionic salt and water. • Ex. NaOH + HClNaCl + H2O • Ex. H2SO4 + 2 KOH  K2SO4 + 2 H2O • Ex. NH4OH + HBr H2O + NH4Br • Ex. 3 HNO3 + Al(OH)3 3 H2O + Al(NO3)3 They are double displacement reactions where one product is always water.

15. pH • Water has the ability to ionize itself. • H2O + H2O ⇄ H3O+ + OH− • So pure water will ionize itself, but has an equilibrium relationship for the forward & reverse reactions. The equilibrium is called Kw or water dissociation equilibrium constant. • At room temperature, its value is approximately 1 × 10−14 • Kw = [H3O+] [OH−] [ ] symbols mean concentration

16. Very few water molecules will be broken apart at any given moment. So the concentrations of H3O+ & OH− are very small. In fact they are the square root of 1 × 10−14. Which is 1 × 10−7 for both H3O+ & OH−. • So adding an acid to the water will disturb the equilibrium and result in an increase in H3O+ concentration. • Adding a base to the water will also disturb the equilibrium and result in a decrease in H3O+ concentration.

17. So the concentration of H+ ions in solution can have a large range of values. To make measuring the amount of acidic H+ ions are present in solution, Chemists use a mathematical trick of logarithms. • pH is the negative logarithm of the H+ ion concentration in solution. • Since the concentration of H+ ions in pure water is 1 × 10−7, then if we take the negative log of this number, we get a value of 7. As you know, this is the value of a neutral solution, which is what pure water would be.

18. You can also calculate a pOH as well. • pH + pOH = 14 • Remember [H3O+] = 1 × 10−7 & [OH−] = 1 × 10−7 so…pH = 7 & pOH = 7 in pure water. Example Problems: • What is the pH of a solution with [H3O+] = 6.82 × 10−9 ? Ans. = 8.166 * Sig. Figs. are tricky with logs!!! • What is the pH of a solution with [OH−] = 3.41 ×10−4 ? Ans. = 10.533

19. There are 12 pH/pOH/[H+]/[OH−] calculation combinations that can be made. • pH to pOH & pOH to pH • [H+] to pH & pH to [H+] • [OH−] to pOH & pOH to [OH−] • [OH−] to pH & pH to [OH−] • [H+] to pOH & pOH to [H+] • [H+] to [OH−] & [OH−] to [H+] Just remember that pH + pOH = 14, pX = −log [ X ]. To get [ ] from a p value, is 10^(−p number) Also, [H+] × [OH−] = 1 × 10−14 can be useful too!

20. Acid Base Titrations & Calculations • When an acid & base react together, they produce an aqueous ionic salt and water. This is called a neutralization reaction. • The reaction is a double displacement reaction, where ions switch places. • Write the reaction and balance it. • A titration is a specialized lab procedure where a solution of unknown concentration is reacted with a solution of known concentration. • With the balanced equation, stoichiometry can be used to find the unknown concentration.

21. EX. 50.0 mL of an unknown solution Ca(OH)2 is titrated with a 0.500 M hydrochloric acid solution. The neutralized end point of the reaction requires 12.7 mL of the HCl solution. What is the concentration of the calcium hydroxide solution? Write & balance the reaction, then calculate! 2 HCl(aq) + Ca(OH)2(aq)  CaCl2(aq) + 2 H2O(l) The stoichiometry results with 0.0635 M Ca(OH)2.