acid - baierscience10

Report
Acid/Base Chemistry
Part 3 (5.4-5.5)
Science 10
CT05D05
Resource:
Brown, Ford, Ryan, IB Chem
Topic 05 – Acids/Bases
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5.1
5.2
5.3
5.4
5.5
5.6
Solutions
Definitions of Acids and Bases
Properties of Acids and Bases
Calculating pH, pOH, H+, OHNeutralization equations
Titrations
5.4 Calculating pH, pOH,
OH
+
H ,
 5.4.1 Calculate the concentration of ions (H+ and
OH-) and acidity (pH and pOH) of strong acids and
bases
 5.4.2 Calculate the above of a mixture of strong
acids and bases
5.4 – pH Scale Proposed
 Søren Sørenson
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Denmark
Biochemist
Early 20th century
Proposed the pH
scale
5.4 - Concentrations of
+
[H ]
 The more acidic a solution is, the more H+ ions
are donated into solution
 Actual concentrations of the hydronium ion,
[H+], are often very small
 Therefore, Sørenson proposed a manipulation of
the concentration of H+ in a way that made the
data much more simple to relate
 The pH scale is based on the logarithm of the
concentration of H+
5.4 - Concentration
 We represent concentration by molarity, therefore
the concentration of an acid and a base will give us
the following information:
+
 0.45 M HCl = 0.45 mol [H ]
 From an acid we find the [H+]
 We can calculate pH directly

0.65 M NaOH = 0.65 mol [OH-]
 From a base we find the [OH-]
 We can calculate pOH directly

If we know the concentration of one, we can find
it of the other:
[H+] [OH-] = Kw = 1 x 10-14
5.4 - Calculating pH
 The pH of a solution is defined as the negative
logarithm of the hydrogen ion concentration (in
mol/L).

pH = -log[H3O+]

pH = -log[H+]
 This calculation results in a pH scale, 0-14
 Therefore, the pH range of solutions are as follows:

Acidic Solutions, pH < 7.0 (0.0-6.9)

Neutral Solution, pH = 7.0 (7.0)

Basic Solutions, pH > 7.0 (7.1-14)
5.4 – Problems with scale
 Find the pH of a 12M solution of HCl
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

pH = -log (12M) = -1.07
Usually the pH of 12 Molar HCl represented as
0.00 and not -1.07?
Any solution with a pH or pOH calculation that
results in a negative number you are welcome to
round to a pH or pOH of zero (0).
 For our purposes in this class, the scale
will be from 0-14!
5.4 - The pH scale is logarithmic
 What does the logarithm scale mean?
 The logarithm is base 10, so a change in one value
has 10x the effect
 Since pH 7 is neutral
 pH 5 is 10 x more acidic than pH 6
 pH 4 is 100 x more acidic than pH 6
 pH 3 is 1000 x more acidic than pH 6
 pH 9 is 10 x more alkaline than pH 8
 pH 10 is 100 x more alkaline than pH 8
 pH 11 is 1000 x more alkaline than pH 8
5.4 - What is pOH?
 pOH is the opposite of pH, and a measure of
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
alkalinity (how basic)
If pH goes down, pOH goes up.
The pOH scale is (opposite pH):
 Basic Solutions:
pOH < 7.0 (0.0-6.9)
 Neutral Solutions: pOH = 7.0 (7.0)
 Acidic Solutions: pOH > 7.0 (7.1-14)
pH + pOH = 14
14 – pH = pOH
Calculated:
pOH = -log [OH-]
5.4 – Common Solutions
Practice #1
 What is the pOH of a 0.005 M Mg(OH)2 ?
 What is the pH ?
-log [OH-] = pOH
-log [0.005] = 2.3 = pOH
14 – pOH = pH
14 – 2.3 = 11.7 = pH
Practice #2
 If 20g of NH3 are dissolved in 2.5 L of distilled
water, what would the pH of the solution be?
20g NH3
1 mol NH3
17g NH3
= 1.17 mol NH3
-log [OH-] = pOH
1.17 mol NH3
2.5 L solution
= 0.47 M NH3
-log [0.47] = pOH = 0.32
14 – pOH = pH
14 – 0.32 = 13.68
5.4 - Strong Acids and Bases
Strong Acids
 HClO4 (perchloric acid)•
 HI (hydroiodic acid)
•
 HBr (hydrobromic acid) •
 HCl (hydrochloric acid) •
 H2SO4 (sulfuric acid) •
 HNO3 (nitric acid)
•
•
We will
calculate
primarily
Mainly
the acids
of halides!
•
with strong monoprotic
acids and bases (BOLD)
Strong Bases
LiOH (lithium hydrox.)
NaOH (sodium hydrox.)
KOH (potassium hydrox.)
RbOH (rubidium hydrox.)
CsOH (cesium hydrox.)
Ca(OH)2 (calcium hydrox.)
Sr(OH)2 (strontium hydrox.)
Ba(OH)2 (barium hydrox.)
Mainly the bases of alkali
metals and some alkaline
earths!
Applicable Solution Def’s
 Strong acids and bases will completely ionize in
water and we can therefore use a “” yields
symbol in the equation. (for weak acids and bases,
an equilibrium arrow would be used ⇌)
 Ionization: Process by which a neutral compound
is split into charged particles by action when
dissolved in liquid water
 Equilibrium: When reactants and products are in
a constant ratio. The forward and reverse reactions
occur at the same rate when a system is in
equilibrium.
Ionization Reactions
 Completion: For those that include a strong
acid or strong base, the reaction will run to
completion and can be shown as such with a
generic ‘yields’ symbol ()
 Equilibrium: For those that include a weak
acid or base or do not go to completion, the
reaction can be represented by an equilibrium
symbol ()
5.5 Neutralization equations
 5.5.1 Balance simple acid base equations
 5.5.2 Conjugate Acid/Base pairs
5.5 - Neutralization
 When equal concentrations of H+ and OH- are
added to one another, a neutral solution results
 NaOH + HCl  NaCl + H2O
 Base + Acid  Salt + Water
 In equal amounts this is always the case, whether
the acid is strong or weak, as long as
concentrations are taken into account
 Try the following, assume complete neutralization:
 Mg(OH)2 + HCl 
 KOH + H2SO4 
5.4 - Aqueous Solutions of A&B
• What happens when you put an acid or a base into
water?
• Each have the property of being electrolytes so will
therefore dissociate
• Water itself can act as an acid or a base
+
+
 H + H2O  H3O
(or H+ + OH-  H2O)
 H3O+  H+ + H2O
(or H2O H+ + OH- )
5.4 – Conjugate Acid/Base
 When acids and bases go through the process of
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donating or accepting protons, they then switch
roles as they can easily reverse the reaction
Acids become conjugate bases
Bases become conjugate acids
Strong become weak
Weak become strong
Arrhenius acid is a substance that produces H+ (H3O+) in water
Arrhenius base is a substance that produces OH- in water
A Brønsted acid is a proton donor
A Brønsted base is a proton acceptor
base
acid
conj. acid
conj. base
A Brønsted acid must contain at least one
ionizable proton!
5.5 - Conjugate Acids and Bases
• Conjugate pairs are two substances that differ by
one H+ (they gain or lose one PROTON)
HCl (acid)  Cl- (conjugate base)
NH3 (base)  NH4+ (conjugate acid)
• When an acid loses a proton it becomes its
conjugate base
H2O  OH-
or
HCl  H+ + Cl-
• When a base gains a proton it becomes its
conjugate acid
H2O  H3O+ or
NH3 + H+  NH4+
5.5 - Conjugate pairs
Conjugate base of the acid H2SO4
HSO4Conjugate acid of the base HSH2S
Conjugate acid of H2O
H3O+
Conjugate base of H2O
OH-
5.5 - Bronsted-Lowry Acids and
Bases
NH4+ + OH-  NH3 + H2O
Acid
Base
C. Base C. Acid
C2H3O2- + H3O+  HC2H3O2 + H2O
Base
Acid
C. Acid C. Base

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