Summary (CAIE) Cambridge A Level Chemistry (9701) - Acids and Bases

Cambridge International AS and A Level Chemistry (9701)
UI CHEM 005
Study Notes & 15’s Acids and Bases Problems with Solutions




Acids and Bases

, The Acid – Base Theory

Comparison Arrhenius Bronsted-Lowry Lewis
Svante August Arrhenius Johan Nicolaus Bronsted &
Inventor & Years Thomas Martin Lowry Gilbert N. Lewis (1923)
(1886)
(1923)
Substances that when
Acid Definition The species that releases Electron pair acceptor
dissolved in water release
protons (H+) = proton donor
H+ ions
Substances that when The species that accept
Base Definition dissolved in water can protons (H+) = proton Electron pair donor
release OH- ions acceptor
Cannot recognizing the key Cannot explain an acid-base
Limitation role of the solvent in the reaction that does not
ionization of a solute involve proton transfer




The Arrhenius Acid – Base Theory


ACID BASE

 Substances that in water release H+ ions  Substances that in water release OH- ions
M(OH)x(aq) Mx+(aq) + xOH-(aq)
HxZ(aq) xH+(aq) + Zx-(aq)
 Base valence: The number of OH- ions that can
 Acid valence: The number of H+ ions that can be produced by 1 molecule of base.
be produced by 1 molecule of acid.
 The remaining base ions: Positive ions are
 The remaining acid ions: Negative ions are formed from bases after releasing OH- ions.
formed from acids after releasing H+ ions.
 Example:
 Example:
NaOH(aq) Na+(aq) + OH-(aq)
HCl(aq) H+(aq) + Cl-(aq)
Ca(OH)2(aq) Ca2+(aq) + 2OH-(aq)
H2SO4(aq) 2H+(aq) + SO42-(aq)


The Arrhenius theory does have limitations. One of the most glaring is in its treatment of the weak base
ammonia, NH3. The Arrhenius theory suggest that all bases contain OH-. Where is the OH- in NH3? To get
around this difficulty, chemists began to think of aqueous solutions of NH 3 as containing the compound
ammonium hydroxide, NH4OH, which as a weak base is partially ionized into NH4+ and OH- ions:

NH3(g) + H2O(l) NH4OH(aq)
NH4OH(aq)  NH4+(g) + OH-(aq)

, The Bronsted - Lowry Acid – Base Theory

Acid = proton (H+) donor Acid  H+ + conjugate base
Base = proton (H+) acceptor Base + H+  conjugate acid

Example:
Determine the conjugate acid-base pairs in the following reactions.



Acid HNO3 Acid H2O

Conjugate Base NO3- Conjugate Base OH-

Base H2O Base CO32-

Conjugate Acid H3O+ Conjugate Acid HCO3-



Substances that can act as both acids and bases are called amphiprotic




Hydrochloric acid is an acid because it donates a proton to water. This means that water is
acting as a Bronsted-Lowry base. The water is accepting a proton.




Water can also act as an acid. When ammonia reacts with water, it accepts a proton from
the water and becomes an NH4+ ion.

, CH3COOH(aq) + H2O(l)  CH3COO-(aq) + H3O+(aq)
Acid Base Base Acid


In reaction above, CH3COOH acts as an acid. It gives up a proton, H+, which is taken up by
H2O. Thus, H2O acts as a base. In the reverse reaction, the hydronium ion, H3O+, acts as an
acid and CH3COO- acts as a base.

When CH3COOH loses a proton, it is converted into CH3COO-. Notice that the formulas of
these two species differ by a single proton, H+. Species that differ by a single proton (H+)
constitute a conjugate acid-base pair. Within this pair, the species with the added H+ is
the acid, and the species without the H+ is the base. Thus, for the reaction above, we can
identify two conjugate acid-base pairs.