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Chemistry Chapter 15
Acid-Base Equilibria: Proton Transfer in Biological Systems
| Term | Definition |
|---|---|
| H₃O⁺ | When describing aqueous solutions, the terms hydrogen ion, proton, and hydronium ion all refer to the same species. Thus, the symbols H⁺ and H⁺(aq) are simply abbreviated forms of this species. (4.5) |
| Arrhenius Acids | Compounds that produce hydronium ions when dissolved in water. (15.1) |
| Arrhenius Bases | Compounds that produce hydroxide ions when dissolved in water. (15.1) |
| Brønsted-Lowry Acids | Substances that donate H⁺ ions. (15.1) |
| Brønsted-Lowry Bases | Substances that accept H⁺ ions. (15.1) |
| Neutralization Reaction | Reaction of strong acid and strong base that yields salt + H₂O. (15.1) |
| Strong Acids/Bases | Complete dissociation. A dilute solution of a strong acid contains no HA molecules. (15.1) |
| Weak Acids/Bases | Dissociate only partly in water, incomplete dissociation (⥂). In a dilute solution of a weak acid, most HA molecules are undissociated. (15.1) |
| Kₐ | The acid-dissociation constant. Stronger acid = higher [H₃O⁺] = larger Kₐ value. (15.1) |
| HA(aq) + H₂O (l) ⇌ A⁻ (aq) + H₃O⁺ (aq) | Derivation of Ka begins with Kc for this general reaction. Kc = [H₃O⁺][A⁻]÷[HA][H₂O]. In general, [H₂O] is treated like a constant and removed. (15.1) |
| Oxoacids (HNO₃, H₂SO₄, HClO₄) | Strong acids. (O ⪰ (O-H) + 2). (15.2) |
| Hydrohalic Acids (HCl, HBr, HI) | Strong acids except for HF. HF is not a strong acid because F has a high electron affinity and is very electronegative. Therefore, it is a strong bond that requires a lot of energy to break. (15.2) |
| Weak Acids | HF, HCN & H₂S (H not bonded to O or halogen), Oxoacids where O=(O-H) or (O-H)+1, Organic Acids -RCOOH. (15.2) |
| Strong Bases | Group 1 & 2 oxides and hydroxides. M₂O or MOH (M = Group 1 Metal). MO or M(OH)₂ (M = Group 2 Metal). (15.2) |
| Weak Bases | Have nitrogen with lone electron pair. Ammonia and Amines (R₃₋ₓNHₓ). (15.2) |
| Kₐ | The extent to which an acid ionization reaction takes place is reflected in the value of its equilibrium constant, K: the greater the value of K, the higher the concentrations of products and the lower the concentrations of reactants at equilibrium. The equilibrium constants that describe acid ionization reactions are given the symbol Ka, where the subscript “a” indicates that the equilibrium involves ionization of an acid. (15.2) |
| Kₐ < 1 | Only a fraction of their molecules donate H⁺ ions in most aqueous solutions. Thus, when Kₐ < 1 the compound is a weak acid. (15.2) |
| Why is water not included in Kₐ expression? | Ionization reactions often consume an insignificant fraction of the water molecules present in most aqueous solutions. (15.2) |
| Kₐ Value Ranking | Hydrofluoric acid >Nitrous acid > Formic acid > Acetic acid > Hypochlorous acid. The larger the value of Kₐ, the greater the concentration of H3O+ ions produced by a given concentration of acid. The numerator contains the term [H₃O⁺] meaning Kₐ will increase when more protons have dissociated. (15.2) |
| Oxygen's Effect on Acid Strength | The second most electronegative element, after fluorine. Oxygen atoms bonded to the central atoms in molecules of oxoacids attract electron density toward themselves and away from the H atoms in the molecules’ O―H groups. This allows these H atoms to interact more strongly with the O atoms in molecules of H₂O. (15.2) |
| Why is HNO₃ a stronger acid than HNO₂? | More oxygen atoms = more delocalization. Greater delocalization of the charge on a nitrate ion helps make nitric acid stronger than nitrous acid. Hypohalous acids PRIMARILY get stronger as their central atom increases in electronegativity (HClO>HBrO>HIO). In other words, central atom's electronegativity creates the largest change and then oxidation number differentiates further. (15.2) |
| Charge Delocalization | A stabilizing force that spreads energy over a larger area rather than keeping it confined to a small area. Delocalization refers to the electrons being distributed throughout the entire area of a molecule rather than attached to a single molecule. (Review) |
| Weak Base Interaction with Water | Molecules like Ammonia, and other amines, are considered bases because the lone pair on Nitrogen is able to pull a Hydronium ion away from a water molecule. This leaves behind a hydroxide ion. Although the hydroxide isn't originally part of the amine, the amine's reaction with water causes it to exist, dissociated in the solution. Brønsted-Lowry model. Amines with more R-groups are typically stronger bases (more electron density around the central Nitrogen) and have a larger Kb value. (15.2) |
| Conjugate Pairs | An acid forms its conjugate base when it donates a H⁺ ion, and a base forms its conjugate acid when it accepts a H⁺ ion. (Ex. NO₂⁻ ions and HNO₂ molecules are Brønsted-Lowry bases and acids, respectively). Conjugate pairs exist when structures differ ONLY because of the donated/accepted H⁺ ion. (15.3) |
| H₂O and H₃O⁺ | Acid(aq) + H₂O(l) ⇌ conjugate base(aq) + H₃O⁺(aq). A conjugate pair. (15.3) |
| H₂O and OH⁻ | Base(aq) + H₂O(l) ⇌ conjugate acid(aq) + OH⁻(aq). A conjugate pair. (15.3) |
| Reciprocality of Conjugate Pairs | Strong acids have very weak conjugate bases and strong bases have very weak conjugate acids. (15.3) |
| Leveling Effect | H₃O⁺, the conjugate acid of H₂O is the strongest acid that can exist in water. Stronger acids just create hydronium. OH⁻ is the strongest base that can exist in water. Stronger bases just hydrolyze to create hydroxide. (15.3). |
| Kₐ₁>Kₐ₂>Kₐ₃ | Polyprotic acids steadily grow weaker as their hydrogens ionize. This is reflected in the value of Kₐ, which is a measure of products (H⁺) over reactants. (15.7) |
| Conceptual Acid Strength | The stronger the H-A bond, the weaker the acid. The greater the bond polarity, the stronger the acid. The greater the electronegativity difference, the more the electrons are pulled toward A. The hydrogen shares less electron density, and thus is easier to separate. (15.2) |
| Periodic Trends | Across a period, acid strength increases (EN increases down a period). Down a group, acid strength increases. The length of the bond increases down a group and thus bond strength decreases. 🠕Bond Length 🠗Bond Strength 🠕Acid Strength. Moving DOWN is typically more influential, because we are comparing the same type of acid. (15.2) |
| Oxoacid Strength | For acids with the same number of oxygen atoms: as EN of central atom increases, acidity increases. (15.2) |
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