Solutions
Unit
Honors
Chemistry
Naming Acids Review:
A. Binary – H +one anion Prefix “hydro”+ anion name +“ic”acid
Ex) HCl
Ex) H3P
hydrochloric acid
hydrophosphoric acid
B. Tertiary – H + polyatomic anion
(oxo)
Ex) H2SO4
Ex) H2SO3
no Prefix “hydro”
end “ate” = “ic” acid
end “ite” = “ous” acid
sulfuric acid
sulfurous acid
Properties of Acids and Bases:
Acid
Base
(alkali)
Reactions
Electrical
with Metals Conductivity
Taste
Touch
sour
looks like
water,
burns,
stings
Yesproduces
H2 gas
electrolyte
in solution
bitter
looks like
water,
feels
slippery
No
Reaction
electrolyte
in solution
Indicators: Turn 1 color in an acid and
another color in a base.
A. Litmus Paper: Blue and Red
An aciD turns blue litmus paper reD
A Base turns red litmus paper Blue.
B. Phenolphthalein: colorless in an acid and
pink in a base
C. pH paper: range of colors from acidic to
basic
D. pH meter: measures the concentration
of H+ in solution
Reactions
• Neutralization: A reaction between an
acid and base. When an acid and base
neutralize, water and a salt (ionic solid)
form.
Acid + Base → Salt + Water
Ex) HCl + NaOH → NaCl + HOH
Arrhenius Definition (1884):
A. An acid dissociates in water to produce more
hydrogen ions, H+.
HCl  H+1 + Cl-1
B. A base dissociates in water to produce more
hydroxide ions, OH-.
NaOH  Na+1 + OH-1
C. Problems with Definition:
• Restricts acids and bases to water solutions.
• Oversimplifies what happens when acids
dissolve in water.
• Does not include certain compounds that have
characteristic properties of acids & bases.
Ex) NH3 (ammonia) doesn’t fit
Bronsted-Lowry Definition (1923):
A. An acid is a substance that can donate hydrogen ions.
Ex) HCl → H+ + Cl–
–
–
Hydrogen ion is the equivalent of a proton.
Acids are often called proton donors.
Monoprotic (HCl), diprotic (H2SO4) , triprotic (H3PO4)
B. A base is a substance that can accept hydrogen ions.
Ex) NH3 + H+ → NH4+
–
Bases are often called proton acceptors.
C. Advantages of Bronsted-Lowry Definition
•Acids and bases are defined independently of how
they behave in water.
•Focuses solely on hydrogen ions.
Hydronium Ion:
Hydronium Ion – H3O+ This is a complex ion
that forms in water.
H+1 + H2O  H3O+1
To more accurately portray the Bronsted-Lowry,
the hydronium ion is used instead of the
hydrogen ion.
STRONG Acid/Base versus WEAK
Acid/Base
Strength refers to the % of molecules that form IONS.
A strong acid or base will completely ionize (>95%
as ions). This is represented by a single ()
arrow.
HNO3 + H2O  H3O+ + NO3A weak acid or base will partially ionize (<5% as
ions). This is represented by a double (↔) arrow.
HOCl + H2O
↔
H3O+ + ClO-
HF < HCl < HBr < HI
increasing strength
7 Strong Acids
HNO3
H2SO4
HClO4
HCl
HI
HClO3
HBr
8 Strong Bases
LiOH
NaOH
RbOH
CsOH
Sr(OH)2
Ba(OH)2
KOH
Ca(OH)2
Strength vs. Concentration
• Strength refers to the percent of
molecules that form ions
• Concentration refers to the amount of
solute dissolved in a solvent. Usually
expressed in molarity.
Ionization of Acids & Bases
• H2SO4  2 H+ + SO4-2
– Sulfuric acid
• H3PO3 
3 H+ + PO3-3
– Phosphorous acid
• Ca(OH)2  Ca+2 + 2 OH-1
– Calcium hydroxide
Conjugate Acid-Base Pairs: A pair of
compounds that differ by only one
hydrogen ion
A. Acid donates a proton to become a
conjugate base.
B. Base accepts proton to become a conjugate
acid.
•
•
A strong acid will have a weak conjugate
base.
A strong base will have a weak conjugate
acid.
Acid (A), Base (B),
Conjugate Acid (CA), Conjugate Base (CB)
NH3 +
H2O
↔ NH4+ +
OH-
HCl +
H2O
↔ Cl-
H3O+
B
A
A
B
CA
CB
+
CB
CA
• Base and Conjugate Acid are a Conjugate
Pair.
• Acid and Conjugate Base are a Conjugate
Pair.
AciDonates & Bases accept
1. H2O
B
+
H 2O
↔
A
B
B
4. OH− + H3O+
B
A
HSO4−
OH−
CB
+
H 2O
CB
3. HSO4− + H2O ↔
A
+
CA
2. H2SO4 + OH− ↔
A
H3O+
SO4−2
CA
+
H3O+
CB
↔
CA
H 2O + H 2 O
CA
CB
The Self-ionization of Water & pH
1. Water is amphoteric, it acts as both an acid and a base in the
same reaction.
Ex)
H2O(l) + H2O(l) ↔ H3O+(aq) + OH-(aq)
Keq = equilibrium constant = [H3O+] [OH-]
Because reactants and products are at equilibrium, liquid water is
not included in the equilibrium expression
@ 25C, [H3O+] = 1 x 10-7 M and [OH-] = 1 x 10-7 M
Kw = ion product constant or equilibrium constant for water
Kw = [H3O+] [OH-] = 1 x 10-14 M2
1.0 x 10-14 M2 = [1.0 x 10-7 M] [1.0x10-7 M]
1.0 x 10-14 =
[H3O+]
[OH-]
Acids: [H3O+] > 1 x 10-7 M
Bases: [OH-] > 1 x 10-7 M
Using Kw in calculations: If the concentration of H3O+ in
the blood is 4.0 x 10-8 M, what is the concentration of
OH ions in the blood? Is blood acidic, basic or neutral?
Kw = [H3O+] [OH-]
1.0 x 10-14 M2 = [4.0 x 10-8 M] [OH-]
2.5 x 10-7 M = [OH-]
slightly basic
The pH scale (1909): the power of Hydrogen
A. Measure of H3O+ in solution.
B. pH = -log[H3O+]
C. Range of pH: 0-14
pH < 7: acid
pH > 7: base
pH = 7: neutral
D. pOH = -log[OH-]
E. pH + pOH = 14
H+
OH-
14
1
pH
[H3O+]
[OH-]
14
1x10-14
1x100
13
1x10-13
1x10-1
12
1x10-12
1x10-2
11
1x10-11
1x10-3
10
1x10-10
1x10-4
9
1x10-9
1x10-5
8
1x10-8
1x10-6
7
1x10-7
1x10-7
6
1x10-6
1x10-8
5
1x10-5
1x10-9
4
1x10-4
1x10-10
3
1x10-3
1x10-11
2
1x10-2
1x10-12
1
1x10-1
1x10-13
Significant Digits Rule
• The number of digits AFTER THE
DECIMAL POINT in your answer
should be equal to the number of
significant digits in your original
number
• Ex -log[8.7x10-4M]
– Calc Answer = 3.0604807474
– Sig Fig pH = 3.06
Concentration
• Percent concentration by mass (mass %
– (solute/solution) x 100% = % Concentration
• Molarity (M)
– Moles of solute/Liters of solution = mol/L
• Molality (m)
– Moles of solute/mass of solvent = mol/kg
• ppm and ppb
– Used for very dilute solutions
• Dilution – a process in which more solvent is added
to a solution
– How is this solution different?
• Volume, color, molarity
– How is it the same?
• Same mass of solute, same moles of solute
– In Dilution ONLY – M1V1 = M2V2
Dissolution Process
electrolyte
• Ionic Compounds
NaCl(s)  Na+1(aq) + Cl-1(aq)
nonelectrolyte
– For dissolution to occur, must overcome solute attractions
and solvent attractions.
– Dissociation Reaction: the separation of IONS when an
ionic compound dissolves (ions already present)
– Try calcium chloride
Dissolving
NaCl in water
hexahydrated for Na+1;
Solvation: process of solvent molecules
most cations have 4-9 H2O molecules
surrounding solute
6 is most common
Hydration: solvation with water
V. Solution Stoichiometry
A. Many reactants are introduced to a reaction chamber
as a solution.
B. The most common solution concentration is molarity.
molarity =
mol/liter
C. Examples
1. Excess lead(II) carbonate reacts with 27.5 mL of 3.00M nitric
acid. Calculate the mass of lead(II) nitrate formed
PbCO3 + 2HNO3  Pb(NO3)2 + H2CO3
2. Calculate the volume, in mL, of a 0.324 molar solution of
sulfuric acid required to react completely with 2.792 g of
sodium carbonate according to the equation below.
H2SO4 + Na2CO3  Na2SO4 + CO2 + H2O
Energy Changes
• Heat of solution = Hsoln
• Endothermic
– Solute particles separating in solid
– Solvent particles moving apart to allow solute to enter
liquid
– Energy absorbed
• Exothermic
– Solute particles separating in solid
– Solvent particles attracted to solvating solute particles
– Energy released