1. Teacher
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2. Acids & bases
Basic Chemistry

3. Definition of the pH in H2O
Water molecules dissociate to give hydronium (H3O+) hydroxide (OH-) and ions:
(autoprotolysis or self-ionization)
2H2O H3O+ + OH-
In neutral water at 25 °C, [H3O+] = [OH-] = 10-7 M. The P operator indicates -log10:
pH = -log10([H3O+]) = 7 in neutral water
Addition of acid to water leads to pH < 7, ACIDIC
Addition of base to water leads to pH > 7, BASIC/ ALKALINE

4. Importance and applications
Some examples:
pH of human blood ≈ ; pathology indicator
Acids and bases are very important in industrial processes:
As reactants: Ammonium nitrate production
As catalysts: Formation of carbocations in oil refining
Control of pH and understanding of acid-base chemistry is important in virtually all scientific research structure design & maintenance, safety…..

5. Lewis acids & bases Do not necessarily comprise a different group!
….A different feature is emphasized:
Lewis base: electron-pair donor; Lewis acid: electron-pair acceptor
Example:

6. Typical acids & bases Acids: Hydrogen halides (HF, HCl, HBr, HI)
Oxyacids (H2CO3, H3PO4, HNO3, HNO2…)
(Lewis acids) High-charge (transition) metals (Al3+ , Cr3+, Fe3+)
Bases
Alkali and alkaline earth hydroxides (NaOH, KOH, Ca(OH)2…)
Alkali and alkaline earth oxides (K2O, CaO, MgO…)
Metal hydrides (LiH, NaH, KH….)

7. Acid dissociation constant Ka; pKa = -log(Ka)
The strength of acids and bases is described by their equilibrium constants
General acid HA: HA+ H2O H3O+ + A- Equilibrium constant K: 𝐾= 𝑎 H3O+ 𝑎[A−] 𝑎 HA 𝑎[H2O] ≈ H3O+ A− HA H2O …But the concentration of water is almost constant: 𝐾 𝑎 =K H2O = H3O+ [A−] HA
Acid dissociation constant Ka; pKa = -log(Ka)

8. General base B: B + H2O BH+ + OH- 𝐾 𝑏 = BH+ [OH−] B
The strength of acids and bases is described by their equilibrium constants
General base B: B + H2O BH+ + OH- 𝐾 𝑏 = BH+ [OH−] B
Base dissociation constant Kb; pKb = -log(Kb)

9. Acid-base conjugate pairs
HA/A- and B/BH+ are both acid/base conjugates:
HA B A BH Acid Base Conjugate base Conjugate acid Reactions always proceed in the direction of the weaker acid and base! pKa + pKb = pKw ; The conjugate base of a strong acid is a very weak base (non-reactive); The conjugate base of a weak acid is a weak base

10. Acid-base conjugate pairs

11. While-> HF < HCl < HBr < HI?
Acid strength:
Why? NH3 < H2O < HF ?
Electronegativity: N (3.04) < O (3.44) < F (3.98)
While-> HF < HCl < HBr < HI?
Bond dissociation energies (kJ/mol):
HI (287) < HBr (354) < HCl (419) < HF (543)

12. pH calculation in aqueous solutions
Strong acids (pKa < 0) dissociate completely:
[H3O+] = CHA, pH = -log(CHA)
Weak acids dissociate negligibly, so [HA] ≈ CHA
𝐾 𝑎 = H3O+ [A−] HA
[H3O+] = [A-] = X
X2 = Ka[Ha] ≈ KaCHA
X = [H3O+] = 𝐾 𝑎 𝐶 𝐻𝐴
pH = −log 𝐾 𝑎 𝐶 𝐻𝐴 = 1 2 (𝐾 𝑎 −𝑙𝑜𝑔 (𝐶 𝐻𝐴 ))

13. Summary Reactions proceed to yield the weaker acid and base
Acids are proton donors or electron-pair acceptors
Bases are proton acceptors or electron-pair donors
pKa of acids and pKb of bases describes their strength; the lower pKa (pKb) the stronger the acid (base)
The strength of a conjugate base is easily deduced from the strength of the acid:
pKa + pKb = pKw
Reactions proceed to yield the weaker acid and base
The pH of a strong acid in aqueous solution: pH ≈ -log(CHA)
The pH of a weak acid in aqueous solution: pH ≈ (𝐾 𝑎 −𝑙𝑜𝑔 (𝐶 𝐻𝐴 ))

14. Let’s look at Henry’s law…
The concentration of a gas dissolved in any liquid is proportional to it pressure in the atmosphere/gas in contact with the liquid:
C = KP, with K the constant, C the concentration in the liquid, and P the partial pressure
Henry’s constants for water at 25 °C
Gas
K (mol L-1 atm-1)
CO2­­­­­­­­­­­­
3.4×10-2 N2
7×10-4 He
3.7×10-4
But note: 𝐻 2 𝐶 𝑂 3 (𝑎𝑞) 𝐶 𝑂 2(𝑎𝑞) =1.8 ×10 −3
Finally: for H2CO3, Ka1 = 2.5×10-4; Ka2 =4.7×10-11