Acids and Bases in all Different
Places
I. Properties of Acids
• A. Molecular substances which ionize when
added to water to form hydronium (H3O+1)
ions  all acids are electrolytes
• B. React with active metals to form H2(g)
– 1. _ Mg(s) + _ HCl(aq) 
– 2. _ Zn(s) + _ HCl(aq) 
– 3. _ Cu(s) + _ HCl(aq) 
I. Properties of Acids
•
C. Acids affect the colors of indicators
– Universal Indicator 
– Phenolphthalein 
•
•
D. Acids neutralize bases
E. Dilute acids taste sour  think citric acid
and ascorbic acid (Vitamin C)
• **SAFETY TIP: Acids release tremendous
amounts of heat when you dilute them
– (esp. H2SO4)  ALWAYS ADD ACID TO WATER
II. Naming Acids
• treated as an ionic compound with H+1 (hydrogen ion) as
cation
• negative ion can be nonmetal (binary acid) or polyatomic
anion (ternary acid)
– A) Binary acids – acids that contain a negative ion ending in “ide”
• 1) Formula  Name
– use prefix: hydro– use root of anion’s name
– use suffix: -ic
•
•
a) HCl
b) HBr
hydrochloric acid
hydrobromic acid
c) HF
hydrofluoric acid
II. Naming Acids
• 2) Name  Formula
– follow above rules in reverse
– be sure to balance charges
• a) hydroiodic acid
•
HI
b) hydrosulfuric acid
H2S
II. Naming Acids
• B) Ternary acids
– DO NOT BEGIN WITH “hydro-“!!!!!!!
– use name of polyatomic ion and switch its ending:
Ion Name Ending
Acid Name ending
-ide
Hydro- ____ -ic acid
-ate
_________ -ic acid
• NOTE: sulfur stays “sulfur-” + ending,
phosphorus stays “phosphor-” + ending
II. Naming Acids
• 1) Formula  Name
– a) H2CO3
• CO3  carbonATE  carbonic acid
– b) H2SO4
•
• SO4  sulfATE  sulfuric acid
2) Name  Formula
–
a) acetic acid
•
–
acetIC  acetATE  HC2H3O2
b) phosphoric acid
•
–
phosphorIC  phosphATE  H3PO4
c) nitric acid
•
nitrIC  nitrATE  HNO3
** some acids are stronger than others:
Acid
HI 
HBr 
HCl 
HNO3
H2SO4
H2SO3 
HSO4-1 
H3PO4 
HF 
HNO2 
HC2H3O2 
Conjugate Base
H+1
H+1
H+1
H+1
H+1
H+1
H+1
H+1
H+1
H+1
H+1
+
+
+
+
+
+
+
+
+
+
+
I-1
Br-1
Cl-1
NO3-1
HSO4-1
HSO3-1
SO4-2
H2PO4-2
F-1
NO-1
C2H3O2-1
Ka
very large
very large
very large
very large
large
1.5 x 10-2
1.2 x 10-2
7.5 x 10-3
6.3 x 10-4
5.6 x 10-4
1.8 x 10-5
Acid
H2CO3 
HSO3-1
H2S 
H2PO4-1 
NH4+1 
HCO3-1 
HPO4-2 
HS-1 
H2O
OH-1 
Conjugate
Base
H+1 + HCO3-1
H+1 + SO3-2
H+1 + HS-1
H+1 + HPO4-2
H+1 + NH3
H+1 + CO3-2
H+1 + PO4-3
H+1 + S-2
H+1 + OH-1
H+1 + O-2
Rank the following acids from weakest to strongest: sulfuric acid,
carbonic acid, hydrochloric acid, hydrofluoric acid, acetic acid
H2CO3 , HC2H3O2 , HF, H2SO4 , HCl
Ka
4.3 x 10-7
1.1 x 10-7
9.5 x 10-8
6.2 x 10-8
5.7 x 10-10
5.6 x 10-11
2.2 x 10-13
1.3 x 10-14
1.0 x 10-14
< 10-36
III. Bases - ionic substance which dissociates to form
hydroxide (OH-1) ions in water
* examples: lye (NaOH) , lime (Ca(OH)2) ,
milk of magnesia (Mg(OH)2)
• Naming Review. Name (or give the formula for) the following
bases:
• 1. NaOH
• sodium hydroxide
• 2. Mg(OH)2
• magnesium hydroxide
• 3. aluminum hydroxide
• Al(OH)3
• 4. ammonium hydroxide
• NH4OH
IV. Properties of Bases - often referred
to as caustic or alkaline substances
• A. Bases are electrolytes - dissociate in water
to form OH-.
• B. Bases affect the colors of indicators.
– Universal Indicator 
PURPLE
– Phenolphthalein  MAGENTA
• C. Bases neutralize acids.
• D. Water solutions are bitter and slippery.
• E. Emulsify fats and oils  this is why they
are useful in soap
V. Salt – any ionic compound that
does not contain hydroxide (OH-1)
•
* all are good electrolytes
formed by a neutralization reaction
Acid + Base  Salt + Water
• 1) _____ HCl(aq) + _____ NaOH(aq) 
• 2) _____ H2SO4(aq) + ____ KOH(aq) 
• 3) ____ HBr(aq) + _____ Ca(OH)2(aq) 
• 4) _____ HC2H3O2(aq) + _____ NaOH(aq) 
Acid, Base, Salt, or Neither:
• 1. NaCl
•
Salt
• 1. KBr
•
Salt
• 6. KOH
• Base
2. KCl 3. KOH
4. SO2 5. NH4C2H3O2
salt
base
neither
salt
2. H2SO4 3. HgCl2 4. Al(OH)3 5. HCl
acid
salt
base
acid
7. CaO 8. K3PO4 9. CO2 10. NH4OH
salt
salt
neither
base
VI. pH – a mathematical way of
measuring how acidic a solution is
H+
Acidity
7
neutral
less than 7
acidic
greater than 7
basic
10 M
1
2
battery
lemon
juice
3
vinegar
4
10-10M
10-7M
10-4M
-1
acid
pH
5
6
milk
7
8
9
sea
water
10-13M
10 11 12 13
milk
of
magnesia
lye
• It’s a logarithmic scale; that means each step
is worth 10
– lemon juice is 10 times more acidic than vinegar
– battery acid is 10 times more acidic than lemon
juice
– How many times more acidic is battery acid than
vinegar?
H+
1
2
battery
acid
lemon
juice
3
vinegar
4
10-10M
10-7M
10-4M
10-1M
5
6
milk
7
8
9
sea
water
10-13M
10 11 12 13
milk
of
magnesia
lye
H+
10 M
1
2
3
5
7
6
milk
vinegar
battery
acid
4
10-10M
10-7M
10-4M
-1
8
9
sea
water
10-13M
10 11 12 13
milk
of
magnesia
lemon
juice
Color scale for Universal Indicator:
pH:
Red
Orange
Green
3
5
7
Blue
Purple
9
11
lye
• Which of the solutions above is the most acidic?
• Battery acid
• 2) Which of the solutions above is the most
basic?
• Lye
• 3) Look at the solutions that your teacher is
testing with universal indicator.
– Label each as acidic, basic, or neutral
– Estimate the pH based on the color
– Rank the substances from most acidic to least acidic
Substance
acidic, basic, neutral
1
2
3
4
RANK:
pH
VII. Buffer - a solution which is able to
resist major changes in pH
• example: HC2H3O2(aq) H+1(aq) + C2H3O2-1(aq)
• common-ion effect - by adding a salt with the negative
ion (NaC2H3O2, KC2H3O2), we increase the
concentration of that ion, therefore:
• add H+1:
• the acid will react with the acetate ion to produce
molecular acetic acid, thus “neutralizing” it and
keeping the pH the same
• add OH-1:
• the base will react with the molecular acetic acid to
produce acetate ions, thus “neutralizing” it and
keeping the pH the same
BLOODY BUFFERS!!
• biological example: carbonic acid/bicarbonate
in blood  Hold your Breath!!!
• There is a balance between the ions which
acts as a buffer, keeping the pH of the blood
right around 7.4. The hemoglobin molecule
in red blood cells can only withstand pH
extremes of 7.2-7.6
VIII. Acid-Base Indicators - chemicals
specifically designed to show specific
colors in acids and different colors in bases
Indicator
methyl violet
methyl yellow
bromophenol blue
methyl orange
methyl red
litmus
bromothymol blue
phenol red
phenolphthalein
thymolphthalein
alizarin yellow
pH Range below pH color above pH color
0.0 – 1.6
yellow
blue
2.9 – 4.0
red
yellow
3.0 – 4.6
yellow
blue
3.2 – 4.4
red
yellow
4.8 – 6.0
red
yellow
5.5 – 8.0
red
blue
6.0 – 7.6
yellow
blue
6.6 – 8.0
yellow
red
8.2 – 10.6
colorless
red
9.4 – 10.6
colorless
blue
10.0 – 12.0
yellow
red
IX. Acid-Base Neutralization
H+1 + OH-1  H2O
• if you have 35 molecules of acid, 35 molecules
of base will neutralize it
• equivalence point - when an equivalent
amount of OH-1 ions has been added to H+1
ions  it’s “neutralized”
X. Acid-Base Titration - lab procedure
used to determine the concentration
of an unknown acid or base solution.
• standard solution – solution whose
concentration is known
• unknown solution – solution whose
concentration you are trying to determine
• MaVa = MbVb
Titration Problems
• 1) If you begin a titration with 20.0 mL of
unknown HCl and titrate it to the equivalence
point using 35.6 mL of 0.600 M standard NaOH,
what is the concentration of HCl?
• Ma(20.0 mL) = (0.600M)(35.6 mL) Ma = 1.07 M
• 2) If you titrate 65.0 mL of an unknown NH3
solution to the equivalence point with 31.2 mL of
a 1.50 M HCl solution, what is the concentration
of the ammonia?
• (1.50M)(31.2mL) = Mb(65.0mL)
Mb = 0.720 M
Titration Problems
• 1)
Ma = ??? Va = 50.0 mL
Mb = 1.50 M
• Ma(50.0mL) = (1.50M)(71.3mL)
Vb = 71.3 mL
Ma = 2.14 M
• 2) What is the concentration of an unknown
NaOH solution if you titrate 100.0 mL of it to
the equivalence point with 43.5 mL of 6.0 M
HCl?
• (6.0M)(43.5mL) = Mb(100.0mL)
Mb = 2.6 M
Titration Problems
• 3) What is the concentration of a vinegar
(HC2H3O2) solution if you titrate exactly 20
drops of it to the equivalence point with 26
drops of 0.600M NaOH?
• Ma(2θdr) = (0.600M)(26dr)
Ma = 0.78 M