Overarching Question: What is the nature of acids, bases, and salts?
Focus Question 1: How does the Arrhenius theory explain acid-base behavior?
A. Svante Arrhenius
1. Conductivity- acids and bases conduct current in aqueous solutions; they act as electrolytes
2. Ionization of acids- acids are “pulled apart” by polar water molecules; water causes them to become
charged; this causes hydronium (H3O+) ions to form
HCl (g) + H2O (l)  H3O+ (aq) + Cl-(aq)
H2SO4 (l) + 2H2O (l)  2H3O+ (aq) + SO42-(aq)
H3PO4 (l) + 3H2O (l)  3H3O+ (aq) + PO43-(aq)
3. Dissociation of bases- bases are already charged; water just separates them
NaOH
Ca(OH)2
Na+ (aq) + OH- (aq)
Ca2+ (aq) + 2OH- (aq)
B. Arrhenius Definition for Acids- acids release hydrogen (H+) or hydronium (H3O+) ions in aqueous solution
C. Arrhenius Definition for Bases- bases release hydroxide (OH-) ions in aqueous solution
FQ 2: What occurs as acids and bases combine?
A. Neutralization
1. Definition- neutralization occurs as acids combine with bases; always produces a salt (ionic
compound) and water
2. ExamplesHCl (aq) + NaOH (aq)  NaCl(aq) + HOH (l)
H2SO4 (l) + 2LiOH (l)  Li2SO4(aq) + 2HOH (l)
3. Water- water (HOH or H2O) is neutral because it contains as many acid ions (H+) as base ions (OH-)
B. Salt
1. Definition- salts are neither acids nor bases, no H+ or OH- ions present
2. Examples- NaCl, KI, MgF2, etc.
FQ 3: How do we show the presence and strength of acids and bases?
A. Ionization and Strength
1. Misconception from formulas- especially for acids, using the formula to determine strength is a
problem; the number of acid/bases ions present falsely indicates actual strength; they indicate potential
strength
Unit 09 UbD C Notes Outline
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B. pH
1. Definition- pH is the actual measure of hydrogen ion concentration; one definition works for both
acids and bases
2. Derivation of pH numbers- taken from negative exponents in the pH formula; nature determined the
number range for pH values
3. pH ranges
a. Acids- 0 ≥ but < 7
b. Bases- 7 > but ≤ 14
c. Neutral- 7 only
4. Calculation of pH
a. Formula- pH = -log[H+] ; the negative sign artificially turns the value into a positive number
b. Logarithms- represent very large or small numbers
FQ 4: How does the Brønsted-Lowry theory of acids and bases compare and contrast to the Arrhenius
theory?
A. Brønsted-Lowry Theory- necessary because not all substances behaving as acids fit the Arrhenius definition;
under this definite acids always exist in combination with bases; form what are called conjugate pairs
1. Acid definition- acids are proton donors; hydrogen ions are only single protons
2. Base definition- bases are proton acceptors
3. Conjugate pairs- differ by one proton or hydrogen ion (H+)
HF acid (donor) – conjugate base F- (acceptor)
F- base (acceptor) – conjugate acid HF (donor)
H2SO4 acid (donor) – conjugate base HSO4- (acceptor)
HSO4- base (acceptor) – conjugate acid H2SO4 (donor)
B. Amphiprotic or amphoteric substances will at times behave as an acid and at others behave as a base. Water
is a good example.
C2H3O2- + H2O  HC2H3O2 + OH[Water donates the H+ as an acid.]
H2CO3 + H2O  H3O + + HCO3[Water accepts the H+ as a base.]
Unit 09 UbD C Notes Outline
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