1. An Arrhenius
1.1. Acid - dissociates in water to release hydrogen ions
1.2. Base - dissociates in water to release hydroxide ions
2. A Brønsted-Lowry
2.1. Acids (H+)
2.1.1. A proton donor
2.1.2. Strong Acids ( completely dissociate into ions in solution)
2.1.2.1. HCl, HBr, HI, HNO3, H2SO4, HClO3, HClO4
2.1.3. Ka
2.1.3.1. pKa = -log(Ka)
2.1.4. Weak Acids
2.1.4.1. Only partially dissociate
2.2. Bases (OH-)
2.2.1. A proton acceptor
2.2.2. Strong Bases (completely dissociate into ions in solution)
2.2.2.1. LiOH, NaOH, KOH, RbOH, CsOH, Ca(OH)2,
2.2.3. Weak Bases
2.2.3.1. Kb
2.3. Water
2.3.1. Can act as either a acid or base
2.3.2. Amphoteric; helps dissociate
2.3.2.1. Makes water amphiprotic
2.3.3. Kw = Ka * Kb
2.3.3.1. pH
2.3.3.1.1. Acidic
2.3.3.1.2. Dynamic Equilibrium
2.3.3.1.3. [H+] > [OH-]
2.3.3.1.4. Neutral
2.3.3.1.5. [H3O+] = 1 * 10^-7 M
2.3.3.1.6. Basic
2.3.3.1.7. The "p" in pH tells us to take the -log of a quantity (in this case, hydrogen ions)
2.3.3.1.8. Calculated by:
2.3.3.1.9. pOH = -log[OH-]
2.3.4. Kw = 1.00 * 10^-14
2.3.4.1. pKw = -log(Kw)
2.3.4.2. pH + pOH = pKw = 14.00
3. form
3.1. Salts
3.1.1. form Neutral Solution
3.1.1.1. cation from strong base, anion from strong acid
3.1.2. form Basic Solution
3.1.2.1. cation from strong base, anion from weak acid,
3.1.2.1.1. form OH-
3.1.3. form Acidic Solution
3.1.3.1. cation from weak base, anion from strong acid
3.1.3.1.1. form H+
4. Chemical Equilibrium
4.1. A state where the rate of the forward reaction (forming products) is equal to the rate of the reverse reaction (forming reactants).
4.1.1. Reaction is now in Equilibrium: reactant and product amounts do not change over time
4.1.1.1. Maintain a fixed ratio
4.1.1.1.1. Can be expressed as the Equilibrium Constant, K
4.2. Can be graphed by:
4.2.1. [Reactants] and [Products] vs Time
4.2.2. Rate of Forward and Reverse Reactions {kf[reactants], kr[products]} against Time
4.3. General Reaction
4.3.1. aA + bB <--> cC + dD
4.3.1.1. Kc = {[C]^c[D]^d}/{[A]^a[B]^b}
4.3.1.1.1. Kc for the reverse reaction is the inverse of that for the forward reaction
4.3.1.2. "aA + bB" - "Left of the reaction"
4.3.1.3. "cC + dD" - "Right of the reaction"
4.4. At Equilibrium
4.4.1. If K is greater than 1:
4.4.1.1. At equilibrium, products will predominate
4.4.1.2. The reaction is product favored
4.4.1.2.1. If Q<K, ΔG<0
4.4.1.3. The equilibrium lies to the right
4.4.1.4. K >>>1, equilibrium "Lies to the right"
4.4.2. At equilibrium, Q=K, ΔG = 0
4.4.3. If K is less than 1:
4.4.3.1. At equilibrium, reactants will predominate
4.4.3.2. The reaction is reactant favored
4.4.3.2.1. If Q>K, ΔG>0
4.4.3.3. The equilibrium lies to the left
4.4.3.4. K <<< 1, the equilibrium "LIes to the left"
4.4.4. ΔG = RTln(Q/K)
4.5. Uses
4.5.1. Le Châtelier's Principle: system shifts to reduce disturbance
4.5.1.1. "If a system at equilibrium is disturbed by a change in temperature, pressure, or a component concentration, the system will shift its equilibrium position so as to counteract the effect of the disturbance."
4.5.1.2. Increase concentration
4.5.1.2.1. Shift to opposite side of reaction
4.5.1.3. Decrease concentration
4.5.1.3.1. Shift to opposite side of reaction
4.5.1.4. "Increasing overall pressure (or decreasing volume) to a system at equilibrium causes the system to shift its equilibrium position as to decrease overall pressure (or increase volume)." (At constant temperature***)
4.5.1.4.1. Decrease pressure
4.5.1.4.2. Increase Volume
4.5.1.4.3. Increase pressure
4.5.1.4.4. Decrease volume
4.5.1.5. Change In Temperature
4.5.1.5.1. If the temperature of a system at equilibrium is increased, the system reacts as if we added a reactant to an endothermic reaction or a product to an exothermic reaction. The equilibrium shifts in the direction that consumes the "excess reagent," namely, heat.
4.5.1.5.2. "Adding heat to a system at equilibrium causes the system to shift its equilibrium position so as to absorb the heat."
4.5.1.5.3. How Temperature affects Keq
5. form
5.1. Buffers: Resist pH Change
5.1.1. equation
5.1.1.1. Henderson-Hasselbalch
5.1.1.1.1. pH = pKa + log[A-]/[HA]
5.1.2. Formed from
5.1.3. Used in
5.1.3.1. Titration Curves
5.1.4. Optimum
5.1.4.1. Buffering capacity: want to be +/- 1 of pH high concentration
6. Formulas for ΔG
6.1. ΔG0 = ΔH0 - TΔS0
6.1.1. therefore
6.1.1.1. -RTlnKeq = ΔH0 - TΔS0
6.2. ΔG0 = -RTlnKeq
6.2.1. therefore
6.2.1.1. lnKeq = (-ΔH/RT) + (ΔS0/R)