- Strong vs. Weak Acids
- Indicators
Mr. Shields
Regents Chemistry
U15 L04
1
Acid strength
We all know there must be a difference between
Different acids.
For instance Vinegar is acetic acid and we all
Put vinegar on our food from time to time.
People used to chew willow bark to help reduce fever.
In doing so they were swallowing salicylic acid (an Organic
Acid).
But would anyone try putting sulfuric acid
(car battery acid) instead of vinegar on their food
and eat it???
I HOPE NOT !!!
2
Strong vs. Weak acids/bases
Obviously some acids are stronger than others.
But what makes one acid strong and the other
Weak?
A Strong acid is one that is mostly dissociated
(separated into ions) when dissolved in Water.
The same can be said for bases. Strong bases are
Also highly dissociated in water.
3
NOTE:
-Do not confuse strong acid vs. weak
acids & bases with concentration.
- 6M Acetic acid (CH3COOH) is just as
concentrated as 6M H2SO4 or 6M HCl
But it is a much weaker acid
4
Strong & Weak Acids
Example of a Strong Acid:
HCl  H+ + ClLots of H+ & Cl-. Very little undissociated HCl in Sol’n.
Example of a Strong Base:
KOH  K+ + OHLots of K+ & OH-. Very little undissociated KOH in Sol’n.
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Strong vs. Weak acids/bases
So if strong acids and bases are highly dissociated
In solution what defines a weak acid or base?
Weak acids are NOT highly dissociated into their
respective ions in solution
The same can be said for weak bases. They are
Also NOT highly dissociated in solution.
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Weak Acids and bases
• Weak acids/bases don’t ionize very much in solution. Instead
they tend to reach an equilbrium that is mostly the
undissociated acid or base:
C2H4O2  H+ + C2H3O2- (4% dissociated)
“Vinegar”
Lots of C2H4O2. Very little H+, C2H3O2-.
(CH3COOH)
NH4OH  NH4+ + OH-
(CH3COO-)
(in equilibrium)
Lot’s of NH4OH and very little NH4+, OH-
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The greater the dissociation of the acid (or base)
The Stronger the acid (or base)
Strength of acid/base
Degree of
dissociation
1%
5%
40% 90%
100%
Metals react more vigorously with strong acids
than weak acids
Mg + HCl  MgCl2 + H2
Strong acid/
Fast reaction
Mg + CH3COOH  CH3COOMg + H2 Weak Acid/ 8
Slow reaction
H2
Mg &
Strong
acid
Like H2SO4
Mg &
Weaker
acid
Like
Acetic Acid
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Strong Acids
•
•
•
•
HCl
HBr
H2SO4
HNO3
Sulfate
Nitrate
Weak Acids
•H2C2H3O2 (acetic)
•
•
•
•
Strong Bases
NaOH
KOH
Ca(OH)2
Mg(OH)2
Weak Bases
•NH3
•HCN
•HF
•HNO2 (Nitrite) or H2SO3 (Sulfite)
•H2CO3
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Indicators
Remember that pH is a measure of how much H+
is in solution. Let’s look at how we can measure
pH.
There are two ways we might do this:
1) Measure pH directly using electronic
instruments
2) Use INDICATORS
Indicators are substances that change color
Based on the pH of the solution.
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Indicators
Look at Table M in your Reference Table Booklet. The table
lists the color changes of specific indicators at specific pH’s
It also lists
The useful pH
Range of each
Indicator.
Using these
Indicators we
Can determine
The pH of an
unknown.
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Color range of a few indicators
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Solving Problems w/ Indicators
Problem: An Unknown sol’n turns methyl orange
Yellow & turns bromocresol green…green. What
is the pH of the solution?
1) Check Methyl orange in table M:
2) Methy orange
if pH > 4.4 it’s yellow
3) Check Bromocresol green
<3.8 yellow
>5.4 blue
4) pH must be >4.4 but < 5.4
and a green color must be the midpoint of 3.8 and
14 5.4
or about 4.6
Solving Problems w/ Indicators
Problem: An Unknown sol’n causes
Bromothymol blue to turn blue and phenolphthalein
To turn colorless. What is the approximate pH of
The solution?
1) Check Bromothymol blue pH range in table M:
2) Bromthymol Blue
if pH is >7.6 it’s blue
3) Check Phenolphthalein
<8.2 colorless
4) pH is between 7.6 and 8.2
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