6.2 Using Brønsted-Lowry Theory Flashcards
when and what was Svante Arrhenius’ theory?
- acids produced H+ ions in water, bases produced OH+ ions in water
3 limitations of Arrhenius’ theory
(1) there were exceptions
(2) doesn’t account for solvents other than water
(3) doesn’t account for states other than aqueous
when and what was Bronsted-Lowry’s theory
- acids are proton donors, and bases are proton acceptors (concept of conjugate acid-base pairs)
importance of BL theory (3)
(1) shows acidity depends on properties relative to other substances present in reaction, not just on structure
(2) neutralisation didn’t need to involve ionisation to H+ but just proton transfer
(3) hydrolysis of salts to produce pH was separate from acid-base reactions
2 limitations of BL theory
doesn’t explain acidity of acidic oxides (which has no proton transfer)
doesn’t account for acids that don’t have hydrogen
amphiprotic vs amphoteric substance
amphiprotic substances that can both donate and accept protons (weak acid/base) depending on what they react with (doesn’t need to involve hydrogen)
amphoteric substances react as either acids or bases (usually involves hydrogen)
how to recognise BL conjugate acid-base pairs
acid –> base (donated proton)
base –> acid (accepted proton)
weak acid has a … base pair, and vice versa
strong
3 common amphiprotic substances
HPO4 -2
HCO3 -
HSO4 -
two common ways to write equations that identify amphiprotic substances
reaction with HCl and OH-
OR
reactions with H2O
the single arrow is for …
the double arrow is for …
single arrow for strong acid/base
double arrow for weak acid/base
most naturally occurring acids are organic that contains ___ group
COOH (carboxyl)
equilibrium of strong acid equation
right (ionises completely)
equilibrium of weak acid equation
left (majority still molecules)
polyprotic acids…
have more than 2 dissociable protons that undergo stepwise ionisation
multiple hydrogen ions in the acid
degree of ionisation =
[H3O+] / [HA] x 100%
ionisation constant equation
Ka = [H3O+][A-] / [HA]
what is ionisation constant
strength of an acid
a strong acid has Ka…
a weak acid has Ka…
strong, Ka >1
weak, Ka <0.01
relative strength of acids equation
pKa = -log10 (Ka)
lower pKa means…
higher pKa means…
lower pKa = stronger acid
higher pKa = weaker acid
(basically inverted Ka)
which hydrogen ion does formic acid donate HCOOH
from the back. HCOO-
strength vs concentration
strength (strong or weak) is DEGREE OF IONISATION
concentration (concentrated or dilute) is relative to AMT OF WATER PRESENT
electrical conductivity of acids and bases
strong acids and strong bases make good electrolytes
4 strong acids
HCl, HBr, HNO3, H2SO4
the more dilute a weak acid, the …… the %ionisation
greater
strong bases…
dissociate completely
base dissociation constant Kb equation
Kb = [BH+] [OH-] / [B]
(ionic product of water) Kw =
Ka x Kb
accuracy/reliability of indicator vs probe
indicator: inaccurate thus reliable
probe: accurate thus unreliable
is universal indicator an indicator?
no because pH colour ranges vary across indicators
describe litmus colours
blue in neutral/base
red in neutral/acid
which indicator is best for acids and which is for bases, and which is for neutral area
phenolphthalein for basic
methyl orange for acidic
bromothymol blue for neutral
describe phenolphthalein range
colourless until ~8.3 pink acidic
~8.5 fuchsia
describe methyl orange range
~3 orange to yellow ~6/7
describe bromothymol blue
~6.5 yellow
~7 neutral green
~7.5 blue
