The document discusses the key properties and reactions of acids and bases. It defines acids as substances with a pH less than 7 that produce hydrogen ions in water. Bases are defined as having a pH greater than 7 and producing hydroxide ions in water. The document also discusses acid-base indicators, the pH scale, and neutralization reactions between acids and bases that produce salts and water.

1. Chapter 19
“Acids, Bases,
and Salts”

2. Properties of Acids
 They taste sour (don’t try this at home).
 They can conduct electricity.
–Can be strong or weak electrolytes in
aqueous solution
 React with metals to form H2 gas.
 Change the color of indicators
(for example: blue litmus turns to red).
 React with bases (metallic hydroxides)
to form water and a salt.

3. Properties of Acids
 They have a pH of less than 7 (more
on this concept of pH in a later lesson)
 They react with carbonates and
bicarbonates to produce a salt, water,
and carbon dioxide gas
 How do you know if a chemical is an
acid?
– It usually starts with Hydrogen.
– HCl, H2SO4, HNO3, etc. (but not water!)

4. Weak vs. Strong Acids
 Weak Acids do not ionize completely:
Acetic, Boric, Nitrous, Phosphoric,
Sulfurous
 Strong Acids ionize completely:
Hydrochloric, Nitric; Sulfuric,
Hydriodic

5. Acids
have a
pH
less
than 7

6. Acids React with Active Metals
Acids react with active metals to
form salts and hydrogen gas:
HCl(aq) + Mg(s) → MgCl2(aq) + H2(g)
This is a single-replacement reaction

7. Effects of Acid Rain on Marble
(marble is calcium carbonate)
George Washington:
BEFORE acid rain
George Washington:
AFTER acid rain

8. Acids Neutralize Bases
HCl + NaOH → NaCl + H2O
-Neutralization reactions
ALWAYS produce a salt (which is
an ionic compound) and water.
-Of course, it takes the right
proportion of acid and base to
produce a neutral salt

9. Properties of Bases (metallic hydroxides)
 React with acids to form water
and a salt.
 Taste bitter.
 Feel slippery (don’t try this either).
 Can be strong or weak
electrolytes in aqueous solution
 Change the color of indicators
(red litmus turns blue).

10. Weak vs. Strong Bases
 Weak Bases: ammonia; potassium
carbonate, sodium carbonate
 Strong Bases: sodium hydroxide;
sodium phosphate; barium hydroxide;
calcium hydroxide

11. Bases
have a
pH
greater
than 7

12. Acids
 Not all compounds that have
hydrogen are acids. Water?
 Also, not all the hydrogen in an
acid may be released as ions
–only those that have very polar
bonds are ionizable - this is
when the hydrogen is joined to
a very electronegative element

13. Organic Acids (those with carbon)
Organic acids all contain the carboxyl group,
(-COOH), sometimes several of them.
CH3COOH – of the 4 hydrogen, only 1 ionizable
The carboxyl group is a poor proton donor, so
ALL organic acids are weak acids.
(due to being bonded to the highly electronegative Oxygen)

14. Why Ammonia is a Base
 Ammonia can be explained as a
base by using Brønsted-Lowry:
NH3(aq) + H2O(l) ↔ NH4
1+
(aq) + OH1-
(aq)
Ammonia is the hydrogen ion
acceptor (base), and water is the
hydrogen ion donor (acid).
This causes the OH1- concentration
to be greater than in pure water,
and the ammonia solution is basic

15. Acids and bases come in pairs
 A “conjugate base” is the remainder of
the original acid, after it donates it’s
hydrogen ion
 A “conjugate acid” is the particle
formed when the original base gains a
hydrogen ion
 Thus, a conjugate acid-base pair is related by
the loss or gain of a single hydrogen ion.
 Chemical Indicators? They are weak
acids or bases that have a different
color from their original acid and base

16. Acids and bases come in pairs
 General equation is:
HA(aq) + H2O(l) ↔ H3O+
(aq) + A-
(aq)
 Acid + Base ↔ Conjugate acid + Conjugate base
 NH3 + H2O ↔ NH4
1+ + OH1-
base acid c.a. c.b.
 HCl + H2O ↔ H3O1+ + Cl1-
acid base c.a. c.b.
 Amphoteric – a substance that can act as
both an acid and base- as water shows

17. Ion Product Constant
 H2O ↔ H1+ + OH1-
 Kw is constant in every aqueous solution:
[H+] x [OH-] = 1 x 10-14 M2
 If [H+] > 10-7 then [OH-] < 10-7
 If [H+] < 10-7 then [OH-] > 10-7
 If we know one, other can be determined
 If [H+] > 10-7 , it is acidic and [OH-] < 10-7
 If [H+] < 10-7 , it is basic and [OH-] > 10-7
–Basic solutions also called “alkaline”

18. The pH concept – from 0 to 14
 pH = pouvoir hydrogene (Fr.)
“hydrogen power”
 definition: pH = -log[H+]
 in neutral pH = -log(1 x 10-7) = 7
 in acidic solution [H+] > 10-7
 pH < -log(10-7)
– pH < 7 (from 0 to 7 is the acid range)
– in base, pH > 7 (7 to 14 is base range)

20. Measuring pH
 Why measure pH?
Everyday solutions
we use - everything
from swimming pools,
soil conditions for
plants, medical
diagnosis, soaps and
shampoos, etc.
 Sometimes we can use
indicators, other times
we might need a pH
meter

21. How to measure pH with wide-range paper
1. Moisten the pH
indicator paper strip
with a few drops of
solution, by using a
stirring rod.
2.Compare the color
to the chart on the vial
– then read the pH
value.

22. Some of the
many pH
Indicators
and their
pH range

23. Acid-Base Indicators
 Although useful, there are limitations
to indicators:
–usually given for a certain
temperature (25 oC), thus may
change at different temperatures
–what if the solution already has a
color, like paint?
– the ability of the human eye to
distinguish colors is limited

24. Acid-Base Indicators
 A pH meter may give more definitive
results
–some are large, others portable
–works by measuring the voltage
between two electrodes; typically
accurate to within 0.01 pH unit of
the true pH
–Instruments need to be calibrated

25. Acid-Base Reactions
Acid + Base  Water +
Salt
 Properties related to every day:
–antacids depend on neutralization
–farmers adjust the soil pH
–formation of cave stalactites
–human body kidney stones from
insoluble salts

26. Acid-Base Reactions
 Neutralization Reaction - a reaction in
which an acid and a base react in an
aqueous solution to produce a salt
and water:
HCl(aq) + NaOH(aq)  NaCl(aq) + H2O(l)
H2SO4(aq) + 2KOH(aq)  K2SO4(aq) + 2 H2O(l)