Ch. 15–Acid/Base Notes
Properties of Acids and Bases: You need to be familiar with the FIVE properties of acids on p. 454 and bases on p. 457-458.
Also be familiar with the three major theories of Acids/Bases:
Arrhenius: p. 459, Bronsted-Lowry p. 464-467 Lewis p. 467-468 (summary table p. 468)
ACIDS GIVE UP A PROTON,
BASES RECEIVE A PROTON. A PROTON
IS ALSO AN H+ ion.
When an acid or a base is put into water, it becomes aqueous (aq). Example: HCl in water becomes H+(aq) and Cl-(aq) ions. The water then takes the H+(aq) ion and makes H3O+(aq) The HYDRONIUM ion.
Amphoteric: (p. 471) Water is amphoteric. This means that water can act as either an acid or a base. For example, if you pour HCl into water, the water would accept a proton to become a hyrdonium ion (H3O+ (aq)). If you pour ammonia into water, the water will give away a proton to become a hydroxide ion (OH- (aq) ).
How to balance Acid/Base RXNs: Recall that there are single displacement, double displacement, synthesis, decomposition and combustion rxns. Well, in this chapter, you will be introduced to acid/base rxns. In an acid/base rxn, an acid gives up its proton (Hydrogen ion) and a base receives the proton. Here are the three main ways you will see an acid/base rxn:
1. Acid + Base ---à Salt + Water (this occurs when a strong acid mixes w/ strong base)
2. Acid + Water ----à Conjugate Base + H3O+ (aq)
3. Base + Water -----à Conjugate Acid + OH- (aq)
Look at p. 469-472 for a listing of some acid/base neutralization rxns.
Here are some sample acid/base rxns. Notice that the balancing is VERY EASY!
HCl(g) + NaOH(aq) -à NaCl(aq) + H2O(l)
HCl(g) + H2O(l) ---à Cl-(aq) + H3O+ (aq)
NH3(g) + H2O(l) ß----à NH4+ (aq) + OH- (aq) (this is household “ammonia”)
What are conjugate acid-base pairs (p. 469-471 and Table 15-6): The conjugate base of CH3COOH acid is CH3COO-. The conjugate base of HCl acid is Cl-. Strong acids have weak conjugate bases. Weak acids have strong conjugate bases. For example: CH3COO- really wants a proton more than CH3COOH wants to give away its proton. Similarly, HCl wants to give away its proton much more than Cl- wants to receive one.
Strong vs. weak
acids/bases (these cause neutralization reactions – Acid + Base makes a Salt
and a Water – p. 474-475): Two
different acids with the same Molarity (concentration) may have different
strengths. 1M of HCl acid is much
stronger than 1M of CH3COOH (acetic acid). Strength depends upon the number of H+ ions produced
per mole of acid. A highly concentrated acid may be a weak
acid. And vice-versa. Concentration has nothing to do with the
equilibrium of the acid.
Strong acids are: HCl (Hydrochloric – Murietic), HBr, HI, HNO3 (nitric) and H2SO4(sulfuric).
Weak acids are: CH3COOH (Acetic – vinegar), H2CO3 (carbonic) H2O (water)
H3PO4 (phosphoric), HNO2 (nitrous), HOCl (Hypochlorous –in fish tanks) NH4+ (ammonium)
Strong bases are: NaOH (sodium hydroxide – Lye – Drano), KOH (potassium hydroxide),
NaOCl (sodium hypochlorite – bleach)
Weak bases are: NH3(g) (Ammonia. Notice ammonia is a gas. Most acids and bases are aqueous. Ammonia you buy in the store is in equilibrium between (aq) and (g) NH3(g) ß-> NH3(aq)
QUIZ HINTS: NO homework for Chapter 15. Just read this to help for the quiz!
1. MEMORIZE STRONG/WEAK ACIDS AND BASES
GIVEN ABOVE.
(tables 15-3/15-4 p. 460-461 have more,
but I’m only going to test you on those above)
2. Memorize the chemical formula for the names
of the acids given above in the Strong/Weak Acid/Base list. You should also know any “street” name for
those acids or bases.
3. Know how to balance an acid/base
reaction. How to find a conjugate base
or acid.
4. Know the 5 properties of acids and bases
5. Know the basic ideas behind the three
descriptions of acids/bases (Arrhenius, Bronsted-Lowry, Lewis)
Chapter 16 Acid/Base Notes
Ka or Kb or Keq or Kw… What are these things? (p. 569 from chapter 18)
K is a constant which is
different depending upon the substance in question. For example, HCl has a different value for K than does CH3COOH.
The subscript just tells you if it is a
K for acid (a) or base (b) or
equilibrium (eq) or water (w)
Fine, but still, what is K?
K = [products]/[reactants]. [] = concentration in molarity
(moles/liter)
Ka values for weak acids are all less than 1. This means there is more reactants and less products - it favors the reactant side. In other words, weak acids hardly dissociate in water. You will find a listing of Ka values on your homework paper (questions 25-30)
You will not find a table
of Ka values for strong acids (such as HCl) because their rxn with
water has much more product and very little reactant is left. Strong acids have a Ka greater than 1 and
thus favors the product. You can assume
there is no equilibrium in a strong acid/water rxn for our purposes. You can assume that the reactants all turn
into products.
Equilibrium constants (Kw) for water (p. 571 chapter 18): Water will break down spontaneously into H3O+ and OH- ions. But the equilibrium of this reaction is much less than 1 and favors the reactants.
H2O(l) + H2O (l) <-> H3O+ (aq) + OH- (aq).
If you were to write a Keq for this rxn, remember to put the concentration of the products above those of the reactants. And, since the reactants are both liquids, they do not appear in the Keq expression.
Therefore, Kw
= [H3O+] [OH-]. But, this Keq
is called Kw (w is for water)
Ka is for acids other than water.
The Kw= 1.0 x 10-14. So, that means the [H3O+] is 1.0 x 10-7M and the [OH-] is also 1.0 x 10-7M. Since the mole ratio between [H3O+] and [OH-] is 1:1, they are the same value. Remember that concentration is measured in units of Molarity (moles/liters)
Notice that Kw is VERY SMALL. This means the rxn favors the reactants!
Logarithms and pH scale: The concentration of H+ (or H3O+ (aq) ) ions in solution changes dramatically as you add an acid to a solution. The pH of an acid is the logrithmic representation of the number of H+ (or H3O+ (aq) ) ions in solution. In order to keep track of it, we must use logarithms. For every 1 change in the pH scale, this represents a 10 fold increase in H+ (or H3O+ (aq) ) ions in solution.
The concentration of H+ ions in solution is usually less than 1M. This means that the concentration is going to be in scientific notation with a negative value (example: 2.3 x 10-8M)
The pH = -log[H+] or pH = -log[H3O+]
Learning about Logs.
The log of 2.3 x 10-8M is –7.64. This means that if you took 10 and raised it to the –7.64 power, you would get 2.3 x 10-8M as your answer. To do this on your calculator, enter the number 2.3 x 10-8 and hit the log key. It should give you –7.64. +7.64 will be the pH of the concentration of 2.3 x 10-8M H+ ions in solution. It is positive because pH is equal to the NEGATIVE value of the Log of [H+] ions.
If you
know the pH to be 7.64, then you simply plug –7.64 into your calculator, and
hit the INVERSE LOG key and you should get 2.3 x 10-8M concentration
as your answer. Notice that the pH was
POSITIVE 7.64, so you must plug a –7.64 into your calculator before hitting the
INVERSE LOG key. So, as the concentration of [H+]
ions increases, the pH will decrease.
Reading the pH scale (p. 485-486): Something which is acidic has a pH of less than 7 (or, more [H+] ions than 1.0 x 10-7M). Something which is basic has a pH greater than 7 (or, less [H+] ions than 1.0 x 10-7M).
So, a [H+] of 1.0 x 10-8M would be basic with a pH of 8.
So, a [H+] of 1.0 x 10-6M would be acidic with a pH of 6.
To say it another way: A low [H+] ions is basic. A high [H+] ions is acidic.
YOU MAY USE H+ and H3O+ interchangeably for our purposes
Rules to live by: A strong acid (such as hydrochloric
acid) will completely dissociate in water.
Therefore, its concentration is equal to the concentration of [H3O+]. So, if you know the concentration of
hydrochloric acid, simply take the negative log of it to find the pH
A weak acid (such as acetic acid) does not completely dissociate and therefore the concentration of [H3O+] is less than the concentration of acetic acid. You need to first solve for the concentration of [H3O+] by doing the following and then find the negative log.
Titrations (p. 493-503): When you add acid to base, or vice versa, you can change the pH to neutral (pH of 7) if you add the same number of moles of acid to base. When you reach this point, it is called the equivalence point. (You will learn more about pH later, but realize that neutral pH is 7. Your body is pH 7.35 and so is sea water!)
Look at the sample problems for titration on p. 502.
Acid/Base Indicator dyes: Based on the amount of H+ ions in solutions, an indicator dye can have the majority of its molecules attached to a H+ ion, or a minority of its molecules attached by an H+ ion. Different indicators change colors depending upon the amount of H+ ions present. Look at table 16-6 on p. 495 see which indicators change at the different levels of pH. A good test question is which indicators are used for measuring the presence of a base and which are for acid.
What do you need to know about indicators
for a quiz? Which indicators are used
for acids and which for bases.
Buffers (p. 570 chapter 18): Since water is right at pH 7, it doesn’t take much acid or base to push it to the acidic or basic side of the scale. A buffer can be added to water to help it resist the addition of an acid or a base.
A buffer is a mixture of a weak acid and its strong conjugate base. For example, in your body, you have a buffer in your blood which is a mixture of the weak acid H2CO3 (carbonic acid) and its strong conjugate base HCO3- (hydrogen carbonate ion). The hydrogen carbonate ion will absorb any acid your body might build up (such as lactic acid produced by your muscles when you run) and turn it into carbonic acid and water. Since carbonic acid is a weak acid and hardly dissociates in water, it will “buffer” your blood pH so it will remain at about 7.35.
If your body receives some base, then the carbonic acid will be tempted to release a proton to form the hydrogen carbonate ion and water.
Base in body: H2CO3 (aq) + OH- (aq) <-> H2O (l) + HCO3- (aq). With a Ka >1 (products favored)
Acid in body: HCO3- (aq). + H3O+ (aq) <->H2CO3 (aq) + H2O (l) With a Ka > 1 (products favored)
Look at the Ka for Carbonic acid and the Ka for Hydrogen carbonate ion in water (p. 663 table 19-6). You will see that these Kas are much less than 1 and therefore they stay on the reactant side. It is not until you add an acid or a base that they then go to product.
Cool applications:
1. CO2 is slightly basic. Therefore, if you have a lot of acid in your body, you breathe very hard to get rid of it by pumping out a large volume of CO2. Running makes you breathe hard because you need O2 to power your ATP rxn and you need to get rid of the lactic acid build up by releasing CO2.
2. Don’t mix household ammonia with bleach. Remember that household ammonia is really ammonium (NH4+). They make household ammonia this way, by bubbling in ammonia gas into water. It is in equilibrium and eventually will turn more into ammonia gas (and thus escape) leaving you with a less potent solution
NH3(g) + H2O(l) ß----à NH4+ (aq) + OH- (aq)
So, if you mix ammonia with bleach, here is what you get (remember that ammonia is really ammonium which is acidic. Ammonia gas is basic. Bleach is very basic.)
NH4+ (aq) + NaOCl (aq) + OH- (aq)----à NH3(g) + Na+ (aq) + OH-(aq)+ HOCl (aq)
You produce ammonia gas which is poisonous. It is true that the ammonia you have in your house is in equilibrium and is also producing quantities of ammonia gas, but it is not as much production as that which is provided by the addition of bleach which makes the ammonium ion give up a proton to become ammonia gas.
What if you add acid to bleach?
HCl + NaOCl ---à HOCl + NaCl. No big deal.
What if you add acid to ammonia?
HCl + NH4+ + OH- --à H2O + NH4Cl Remember that an acid + base gives you salt and water.