Book: UTC: 18 Flashcards
Why the proton is bonded to a water molecule, as H3O+, in all aqueous acid-base systems.
The proton (H+) has a charge density so high that it forms a covalent bond with the O atom in H2O, making hydronium, or H3O+.
The Arrhenius definitions of an acid and a base.
This definition says an acid is a substance with H in its formula that dissociates in water to yield H3O+, and a base is a substance with OH in its formula that dissociates in water to yield OH-.
Why all reactions of a strong acid and a strong base have the same ∆Hº_rxn.
Neutralization involves the same reaction no matter what according to the Arrhenius definition of acids and bases: H+ + OH- -> H2O, so ∆Hº_rxn is always what it is for that reaction when it comes to acid-base reactions: -55.9 kJ.
How the strength of an acid (or base) relates to the extent of its dissociation into ions in water.
Because acids and bases are electrolytes, their strength correlates with electrolyte strength: stronger electrolytes dissociate completely, and weak electrolytes dissociate slightly.
How relative acid strength is expressed by the acid-dissociation constant K_a.
The acid-dissociation constant K_a gives the equilibrium expression for the dissociation of a weak acid. Essentially, it tells how much an acid dissociates into H+ via the equation K_a = [H3O+] [A-] / [HA]
How the autoionization of water is expressed by K_w.
Water itself dissociates slightly into H3O+ and OH-, and the equilibrium constant for this process is given by K_w = [H3O+][OH-] / [H2O]^2, but since H2O is a liquid, K_w is actually just [H3O+][OH-]. It turns out K_w remains constant for water, so that this value is always 1E-14.
Why [H3O+] is inversely related to [OH-] in any aqueous solution.
Because K_w is constant, [H3O+], for example, may be written as [H3O+] = K_w / [OH-], or vice versa, so there is an inverse relationship between the two quantities.
How the relative magnitudes of [H3O+] and [OH-] define whether a solution is acidic, basic, or neutral.
Both ions are present in any aqueous solution, but acidic solutions are defined as having a high concentration of protons, while bases are defined as having a high concentration of OH-, so [H3O+] and [OH-] determine acidity vs. basicity as follows: acidic if [H3O+] > [OH-], neutral if [H3O+] = [OH-], and basic if [H3O+] < [OH-].
The Bronsted-Lowry definitions of an acid and a base and how an acid-base reaction can be viewed as a proton-transfer process.
This definition says that an acid is any species that donates an H+ ion, while a base is any species that accepts an H+ ion, so reactions of acids and bases involve the transfer of an H+ ion, or a proton, from one species to another.
How water acts as a base (or as an acid) when an acid (or a base) dissolves in it.
If an acid is poured into water, H2O will “accept” the proton and result in H3O+ and A- for HA + H2O, so we have HA + H2O ⇌ H3O+ + A-. Conversely, for bases, we have: B + H2O ⇌ BH+ + OH-.
How a conjugate acid-base pair differ.
Conjugate acid-base pairs are molecules of an acid and a base that differ by just one H+, or by one proton.
How a Bronsted-Lowry acid-base reaction involves two conjugate acid-base pairs.
A Bronsted-Lowry acid-base reaction involves the following process: A_1 + B_2 ⇌ A_2 + B_1, where 1 and 2 denote conjugates (i.e. A_1 and B_1 are conjugates).
Why a stronger acid and base react (K_c > 1) to form a weaker base and acid.
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How the percent dissociation of a weak acid increases as its concentration decreases.
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How a polyprotic acid dissociates in two or more steps and why only the first step supplies significant [H3O+].
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