8 Acids & Bases Flashcards
different types of acid-bases [2]
bronsted-lowry
lewis
what defines a bronsted-lowry (B-L) acid-base
b-l acid donates proton, b-l base accepts protons
HCl (g) + H₂O (l) ⇌ H₃O⁺ (aq) + Cl⁻ (aq)
-> from the eqn above:
- HCl (g) and H₃O⁺ (aq) are b-l acids as both can donate protons
- H₂O (l) and Cl⁻ (aq) are b-l bases as they can accept protons
- Cl⁻ is said to be the conjugate base of HCl,, H₂O is said is the conjugate base of H₃O⁺
- conjugate base/acid are the species formed from an acid/base losing/gaining a proton
certain substances can act both as a b-l acid and base – described as amphiprotic
what is pH, Kw and how to calculate it
power of hydrogen (pH) = -log₁₀ [H⁺]
dissociation constant of water / equilibrium constant (Kw) = [H⁺][OH⁻] = 1.0 x 10⁻¹⁴
define the following:
strong
weak
concentrated
corrosive
dilute
strong: completely disassociated into ions
weak: slightly disassociated into ions
concentrated: high number of mols of solute per dm³ of solution
dilute: low number of mols of solute per dm³ of solution
corrosive: chemically reactive
what defines a lewis acid-base
lewis acid accepts electrons, lewis bases donates electrons
in this process, a coordinate covalent bond is formed between lewis acid and base, reacting to form product of acid-base reactions
lewis acid is essentially a normal acid or base, and is only identified openly for substances which are not also b-l
- for organic chem
- used to identify reacting species through the use of curly arrows to explain electron movement in mechanisms
- base of arrow always begins from lewis base and points towards lewis acid
pH v pOH v pKw for strong acid/bases
how to use the following variables and their respective eqns to solve determine neutrality of substance
pH = -log₁₀ [H⁺]
pOH = -log₁₀ [OH⁻]
pKw = -log₁₀ [Kw]
Kw = [H⁺] [OH⁻]
pKw = pH + pOH = 14
for weak acid/bases
Ka v Kb v Kw
let weak acid be HA:
dissociation of HA in water:
- HA (aq) ⇌ H⁺ (aq) + A⁻ (aq)
equilibrium expression for above is Ka = [H⁺] [A⁻] / [HA]
- where Ka is acid dissociation constant
if acids are quite weak, equilibrium concentration is assumed to equal initial concentration
dissociation of weak base:
- B (aq) + H₂O (l) ⇌ BH⁺ (aq) + OH⁻ (aq)
equilibrium expression above is Kb = [BH⁺] [OH⁻] / [B]
- where Kb is base dissociation constant
if reverse reaction where BH⁺ acts as acid, then:
Ka = [B] [H⁺] / [BH⁺]
Ka x Kb = Kw = [H⁺] [OH⁻]
pKa = pKb = 14
what are buffer solutions
what it does, how its created,
buffer solutions resists changes in pH when small amounts of acid or alkali are added to it
- i.e. to keep the pH unchanged
acidic buffer created by mixing a weak acid together with the salt of that acid and a strong base
alkali buffers created by mixing a weak base together with the salt of that base and a strong acid
- i.e. created by mixing a weak acid/base together with its conjugate
this works as the weak acid/base only slightly dissociates in solution, but the salt fully dissociates into ions
- on addition of acid/base, H⁺/OH⁻ will be removed by neutralised by ions, forming either undissociated acid (for basic buffer) or water (for acidic buffer)
*eqns need to learn
titration and equivalence points depending on strength of acid/base
strong acid - strong base: long inflexion point – strong acid is 1-3, strong base is 12-14, therefore inflexion point lasts from minimum 3-12
weak acid/base - strong acid/base: shorter inflexion point
- inflexion point depends on strength of acid and base
