Class 6: Introduction to Acids and Bases Flashcards
Define and identify acids and bases using the Lewis definition.
- Lewis definition of acids and bases:
- Acid: A substance that can donate a proton (H+)
- Base: A substance that can accept a proton (H+)
- Identifying acids:
- Acids contain hydrogen (H) that is bound to a highly electronegative atom (e.g. O, N, Cl)
- The H-X bond is polar, making H+ easy to donate
- Examples: HCl, H2SO4, CH3COOH
- Identifying bases:
- Bases contain atoms with a lone pair of electrons that can accept H+
- Examples: NH3, OH-, CO3(2-)
- The Lewis definition focuses on proton transfer, not creation of H+ or OH-
Describe how the Lewis definition of acids and bases includes all Arrhenius and Bronsted-Lowry acids and bases.
- Arrhenius acids and bases:
- Acids produce H+(aq) in water
- Bases produce OH-(aq) in water
- Included in Lewis definition as special cases
- Brønsted-Lowry acids and bases:
- Acids are proton (H+) donors
- Bases are proton (H+) acceptors
- Directly equivalent to Lewis definition
- Lewis definition is more general:
- Includes Arrhenius and Brønsted-Lowry
- Applies to gas phase and non-aqueous solutions
- Includes compounds without H to transfer
- e.g. BF3 (Lewis acid), NH3 (Lewis base)
- Any Arrhenius or Brønsted-Lowry acid/base can be identified as a Lewis acid/base
- But Lewis includes more compounds beyond just protic cases
So in summary, the Lewis concept generalizes and encompasses the earlier acid-base definitions.
Use curved arrow notation to show the flow of electrons in acid-base reactions.
- Acid donating a proton (H+) to a base:
- Acid: H-X
- Base: :Y
- Reaction: H-X + :Y → X- + H-Y
- Curved arrows: H-X ➝ X-
:Y ➝ H-Y
- Base accepting a proton (H+) from an acid:
- Acid: H-A
- Base: :B
- Reaction: H-A + :B → A- + H-B
- Curved arrows: H-A ➝ A-
:B ➝ H-B
- Lone pairs on Y/B are nucleophilic, attacking H+
- Curved arrows track movement of electron pairs
- From X-H/A-H bonds breaking
- Into new bonds forming
So in summary, curved arrows reveal the flow of electrons, showing proton transfer from acid to base in the making/breaking of covalent bonds.
Memorize the six strong acids.*
- The six strong acids are:
- Hydrochloric acid (HCl)
- Sulfuric acid (H2SO4)
- Nitric acid (HNO3)
- Hydrobromic acid (HBr)
- Perchloric acid (HClO4)
- Hydriodic acid (HI)
- Key points:
- 100% ionized in aqueous solution
- Do not need to memorize their ionization constants
- Exceptions to the general trends in acid strength
- Mnemonics:
- “Hi, Clint, Say No Peace” (first initials)
- “Hot Nachos Can Surely Burn Insides” (first words)
So in essence, commit these six strong acid formulas to memory as they are critical exceptions and do not follow the typical acid strength trends.
*Note: I do not actually have the ability to memorize information myself, as I am an AI assistant without persistent memory. However, I can provide memorization tips for humans.
Describe the autoionization of water and the relationship among Ka, Kb, and Kw.*
- Water autoionizes into hydronium (H3O+) and hydroxide (OH-) ions
- 2H2O ⇌ H3O+ + OH-
- Ion product of water (Kw)
- Kw = [H3O+][OH-] = 1.0 x 10^-14 at 25°C
- For an acid (HA) and its conjugate base (A-):
- Ka = [H3O+][A-]/[HA]
- Kb = [OH-][HA]/[A-]
- Ka and Kb are related by:
- Ka x Kb = Kw
- pKa + pKb = pKw (pK = -log K)
- Stronger the acid, weaker its conjugate base
- Higher Ka, lower Kb
So in essence:
- Water autoionizes to H3O+/OH-
- Ka/Kb measure acid/base strength
- Ka x Kb must equal Kw
- Acid strength inversely related to base strength
Memorize pKa values of important acids (see pKa handout).* (pKa Values to Know and Apply
Define and identify acids and bases using the Lewis definition.
Lewis Acid-Base definition:
An acid accepts an electron while a base donates an electron
Arrhenius Acid-Base definition:
An acid increases the concentration of H+ ions when that are present when added to H20
A base increases the concentration of OH- ions
Bronsted-Lowry Acid-Base definition:
Acids donates a proton (H+) and bases accept a proton
Describe how the Lewis definition of acids and bases includes all Arrhenius and Bronsted-Lowry acids and bases.
It includes both of them, where Arrhenius Acid-Bases and Bronsted-Lowry Acid-Bases are examples of lewis
The Lewis theory of acids and bases is named after the very same chemist responsible for Lewis dot structures.
While the Brønsted-Lowry definition of acids and bases requires that a proton (H+) be either donated or
accepted, the Lewis definition is more extensive and focuses on the transfer of an electron pair:
A Lewis acid is a compound that accepts an electron pair
A Lewis base is a compound that donates an electron pair
- Which molecule is the Lewis acid? The Lewis base?
Lewis acid: ______BF3____________ Lewis base: _________ NH3_____________
Explain how each meets the Lewis definition for acids and bases:
In this reaction, ammonia donates its lone pair of electrons into a bond with boron. Boron
trifluoride accepts the lone pair of electrons, forming the new bond.
Explain how acetic acid (CH3COOH) acts as an Arrhenius acid, a BrønstedLowry acid, and a Lewis acid.
Arrhenius: When dissolved in water, acetic acid dissociates and the hydronium ion concentration is
increased.
Brønsted-Lowry: Acetic acid donates a proton (its hydrogen bonded to oxygen) to water in this
reaction
Lewis: The hydrogen on acetic acid accepts a lone pair of electrons, forming a new bond with the
oxygen in water and breaking its bond with acetic acid.
Common electron sources – Lewis bases
* Nucleophile
* Lone pair of electrons
* Covalent bond (sigma or pi)
Common electron sinks – Lewis acids
- Electrophile
- Partially positive atom
- Carbocation (carbon with a + formal
charge) - H+ (called a proton)
How are the acids in the forward reaction related to the bases in the reverse direction? What is the term for
this relationship?
They are conjugate acid base pairs. Ammonia acts as a base in this reaction. When it forms a bond
with the proton, it becomes ammonium. Ammonium functions as an acid in the reverse reaction.
- Explain why water is more likely to function as an acid than ammonia. (Hint: compare the conjugate base of
ammonia with the conjugate base of water.)
To determine why water functions as an acid in this reaction rather than a base, we need to compare
the conjugate bases in questions 4 and 6 (OH- and NH2-). Oxygen is more electronegative than nitrogen
and can more readily accommodate the negative charge. This means OH- is more stable than NH2. By
stable, we mean lower in potential energy and more, more specifically, lower in free energy. Because
of this, water will function as the acid and ammonia as the base.
to figure out if a base is strong or weak
- Memorize the strong bases:
- Group 1 hydroxides: NaOH, KOH, LiOH
- Group 2 hydroxides: Ca(OH)2, Ba(OH)2
- These bases completely dissociate in water
- All other bases are considered weak bases
- Examples: NH3, CH3NH2, C5H5N (pyridine), etc.
- These bases only partially dissociate in water
- Compare Kb values
- Strong bases have no listed Kb value (assumed 100% dissociated)
- Weak bases have Kb values < 1
- Consider the base strength trend
- Oxides and hydroxides of higher charge density cations are stronger bases
- Amines and nitrogen-containing bases are generally weak
- Use molecular structure
- Ionic hydroxides are generally strong bases
- Bases with N-H bonds are generally weak
- More electronegative substituents make the base weaker
- Degree of dissociation
- Strong bases dissociate 100% in water
- Weak bases have low degree of dissociation (< 5%)
The key is to memorize the common strong bases, and then use Kb values, periodic trends, molecular structure, and degree of dissociation to identify weak bases.
Oxygen is more electronegative than carbon and therefore can stabilize the negative charge better
