∵ because
∴ therefore
Acids and Bases part 1 Equilibrium positions of Acid-ba
Dissociation: process in which molecules separate into smaller parts like atoms and ions
See: https://www.youtube.com/watch?v=tXBI2e_77gs "What is Dissociation in Chemistry" by Vincent
Sapone
Arrhenius definition: 1. Weaker acid and base for
Acids are substances that dissociate in water to produce H30+ ions for the compounds to be
Bases are substances that dissociate in water to produce -OH ions (base has to contain hydroxide ion) 2. Weaker acid larger pKa; w
Does not explain ammonia or many other organic acids and bases 3. Weaker acids and bases a
Strong acids and bases dissociate more than weak so reactions lean towards the products (towards the pKa is important for identifying s
right) equation is favoured and if the r
Strength is measured by concentration of hydronium ions (H30+) ∴ pH = -log10[H3o+] between forward and reverse re
• Use arrows to indicate wh
Bronsted-Lowry definition: • If we don't know pKa the
Bronsted acids as compounds that can donate a proton (H+ ion)
Solvent effects on acidity and bas
Bronsted bases are compounds that can accept a proton (H+ ion) (must have Water: am
electron pair ready to accept proton ie be an anion) Acid: H3O+; Base: -OH
Bronsted-Lowry acids compounds that have an H attached to an electronegative atom Any acid stronger than water (pK
When a base accepts a proton it becomes a conjugate acid capable of donating a proton to give H3O+. When we add HCl t
therefore acid's pKa is actually -1
Acid strength
Acid-dissociation constant (Ka): relative strength of the acid
We do not need to know
Expresses the extent of ionization in water; dissociation constants, just how
to compare strengths of acids
and bases through principles of
bonding and properties of
Same principle for bases pKb > -1
atoms and groups in molecules
∴ water "levels" ranges of pKb an
Since water levels pKb and pKa th
Strong acids Weak acids measured in water
Ka [𝐻 𝑂 ][𝐴 ] As Ka ↑ acid strength ↑ >1 <10-4 Alcohols: (R-OH)
⎯⎯⎯⎯⎯⎯⎯⎯⎯⎯ Acids pKa < -1.9 will protonate et
[𝐻𝐴]
Bases pKb < -1.9 will be protonate
pKa -log10Ka As pKa ↓ acid strength ↑ Around 0 or negative >4 Ammonia pKa of 36 accommodat
Alkanes (saturated hydrocarbons
Effects of Size and
Cl is very electronegativity on Acidity
electronegative More stable anions weaker
and attracts bases ∴ stronger conjugate a
electrons to Electronegativity: more stabl
itself ∴ H is anion
electron Size: negative charge spread
deficient ∴ HCl larger area
can give off a
proton (H+)
Inductive effects on acidity
The pull of the electronegat
The greater the atoms to create greater part
electronegativity the acid. The strong acid wa
of the atom stable conjugate base, reme
which captures stronger the acid. Inductive
the bonding atoms/groups closer to the p
electrons during
dissociation the
Hybridization effects on ac
greater the
extent of
The more s-character the m
dissociation
be more stable
∴ sp more acidic (>) than s
∴ therefore
Acids and Bases part 1 Equilibrium positions of Acid-ba
Dissociation: process in which molecules separate into smaller parts like atoms and ions
See: https://www.youtube.com/watch?v=tXBI2e_77gs "What is Dissociation in Chemistry" by Vincent
Sapone
Arrhenius definition: 1. Weaker acid and base for
Acids are substances that dissociate in water to produce H30+ ions for the compounds to be
Bases are substances that dissociate in water to produce -OH ions (base has to contain hydroxide ion) 2. Weaker acid larger pKa; w
Does not explain ammonia or many other organic acids and bases 3. Weaker acids and bases a
Strong acids and bases dissociate more than weak so reactions lean towards the products (towards the pKa is important for identifying s
right) equation is favoured and if the r
Strength is measured by concentration of hydronium ions (H30+) ∴ pH = -log10[H3o+] between forward and reverse re
• Use arrows to indicate wh
Bronsted-Lowry definition: • If we don't know pKa the
Bronsted acids as compounds that can donate a proton (H+ ion)
Solvent effects on acidity and bas
Bronsted bases are compounds that can accept a proton (H+ ion) (must have Water: am
electron pair ready to accept proton ie be an anion) Acid: H3O+; Base: -OH
Bronsted-Lowry acids compounds that have an H attached to an electronegative atom Any acid stronger than water (pK
When a base accepts a proton it becomes a conjugate acid capable of donating a proton to give H3O+. When we add HCl t
therefore acid's pKa is actually -1
Acid strength
Acid-dissociation constant (Ka): relative strength of the acid
We do not need to know
Expresses the extent of ionization in water; dissociation constants, just how
to compare strengths of acids
and bases through principles of
bonding and properties of
Same principle for bases pKb > -1
atoms and groups in molecules
∴ water "levels" ranges of pKb an
Since water levels pKb and pKa th
Strong acids Weak acids measured in water
Ka [𝐻 𝑂 ][𝐴 ] As Ka ↑ acid strength ↑ >1 <10-4 Alcohols: (R-OH)
⎯⎯⎯⎯⎯⎯⎯⎯⎯⎯ Acids pKa < -1.9 will protonate et
[𝐻𝐴]
Bases pKb < -1.9 will be protonate
pKa -log10Ka As pKa ↓ acid strength ↑ Around 0 or negative >4 Ammonia pKa of 36 accommodat
Alkanes (saturated hydrocarbons
Effects of Size and
Cl is very electronegativity on Acidity
electronegative More stable anions weaker
and attracts bases ∴ stronger conjugate a
electrons to Electronegativity: more stabl
itself ∴ H is anion
electron Size: negative charge spread
deficient ∴ HCl larger area
can give off a
proton (H+)
Inductive effects on acidity
The pull of the electronegat
The greater the atoms to create greater part
electronegativity the acid. The strong acid wa
of the atom stable conjugate base, reme
which captures stronger the acid. Inductive
the bonding atoms/groups closer to the p
electrons during
dissociation the
Hybridization effects on ac
greater the
extent of
The more s-character the m
dissociation
be more stable
∴ sp more acidic (>) than s
