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47 Cards in this Set
- Front
- Back
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Arrhenius definition |
Acid-base reaction that is based on H+ and OH- in WATER! |
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Bronsted-Lowry definition |
Based on any king of reaction in which H+ is transferred from one species to another, Bronsted Lowry acid base reaction |
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Lewis defintion |
Electron pairs; most general type of acid-base reactions |
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Arrhenius Acids |
Aqueous solution; H+ is viewed as a "proton" |
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Arrhenius Bases |
has OH- as part of the molecule and releases it in aqueous solution Hydroxides are Arrhenius bases |
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Bronsted-Lowry Acid-Base Theory |
As long as you have H+ donor (acid) and H+ acceptor (base), you don't need to be in water. H-A + :B <=> :A- + H-B+ |
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Bronsted lowery explains arrhenius, lewis explains bronsted-lowry and then some |
these get more specific |
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HCl molecular compound |
When dissolved in water, it forms "hydrochloric acid" |
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Rate law formula for single reactant |
Rate = k[A]^n |
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Rate law formula for multiple reactants |
Rate = k[A]^m [B]^n |
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Strong Acid |
A strong electrolyte; practically all of the acid molecules ionize (essentially 100% dissociation) Complete dissociation means that molarity of H+ in aqueous solution is equal to initial concentration of acid |
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Strong Base |
A strong electrolyte; practically all the base molecules form OH- ions either through dissociation or reaction with water (essentially 100% dissociation) |
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Weak acid |
A weak electrolyte; only a small percentage of the molecules ionize; equilibrium |
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Weak base |
A weak electrolyte; only a small percentage of bse molecules form OH- ions, either through dissociation or reaction with water; equilibrium - usually Amine containing bases |
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Examples of Strong Acids Table 15.3 |
Hydrochloric (HCl), Hydrobromic (HBr), Hydroiodic (HI), Nitric (HNO3), Perchloric (HClO4), Sulfuric (H2SO4) (Note: just the 1st H+ from sulfuric acid is strong and comes off) SEE BOOK NEEDS UPDATE |
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Examples of Weak Acids |
If it is not a STRONG acid, it is a WEAK acid
Triprotic phosphoric acid, each subsequent H+ is weaker HNO2??? check book |
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General Trends in Acidity |
Stronger an acid is at donating H, the weaker teh conjugate base is at accepting |
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Cations vs. anion and acids? |
water series from slides (NEEDS UPDATE) |
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Autoionization of Water |
Water is amphoteric so it can act as an acid or a base; with PURE water it will do this with itself - this is a process called autoionization - endothermic process (Requires energy to split water into ions; thus if you increase temp. the equil. shifts to the products side (right side) [Hydronium] and [hydroxide] = 10-7 at 25 deg. C |
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Ion product of Water (Kw) |
Kc = [H30+][OH-] (*no H2O because liquid) water is special so Kw is it's own constant at 25C Kw = 1.00E-14 |
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Acid Ionization Constant Ka KNOW AND LOVE THIS!!!! |
Acid strength is measured by size of equilibrium when it reacts with H2O Equilibrium constant for this reaction is the acid ionization constant: Ka = [H30+][A-]/ = [H+][A-]/ [HA] [HA] - Large Ka = stronger acid |
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Acidic and Basic SOlutions |
- Aqueous solutions ALWAYS have H3O+ and OH-
- neutral solutions have equal concentrations - Acids: [H+] > [OH-] - Bases: [OH] > [H+] |
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Measuring Acidity: pH |
Acidity or Basicity of a solution is expressed as pH. "p" refers to "power of" pH = - log[H30+] Example: pH of water = - log[10-7]= 7 |
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Sig figs and logs |
integer part is just place holder; make sure sig figs of number that log is taken of = sig figs in answer. Same for pH answers...
Example: pH of 7.0; 2 sig figs in answer number after decimal ONLY counts with logarithms |
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pH range goes from 0 to 14 in aqueous solution! |
litmus paper?? |
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pOH is another way of expressing acidity/bascitiy |
in PURE water pH = pOH pH + pOH = 14 (*AT 25C) |
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pK is another way to express strength of an acid or base; pKa = - log (Ka) = 10^-pKa |
Stronger acid, smaller pKa Stronger base, smaller pKb |
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Finding pH of a strong acid or strong base solution |
These completely dissociate into ions in water Example: Acids HCl + H2O <=> H3O+ + ClOH Bases KOH + H2O <=> K+ and OH- and H2O |
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Percent ionization *** STUDY THIS for EXAM; how to calculate |
% ionization = [ionized H+ at equil]/[original concentration] |
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Finding pH of mixtures and acids |
1. Usually can ignore weaker acid's contribution because [H30+] contribution is negligible |
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Strong Bases |
Stronger base, more likely to accept H+
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Weak Base |
- Much less than 1% ionization in water |
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Base Ionization Constant (Kb) * KNOW AND LOVE ME TOO |
Base strength is measured by size of the eq. constant when it reacts with H20 |
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Note: conjugate base of strong acid is so weak that it is essentially neutral?? Strong Acid + Strong Base in equal molar mixture; the pH of mixture is 7 |
Strong Acid + Strong Base = Neutral solution |
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Relationship of Ka and Kb for Weak acid + conjugate base *calculation problems with this |
Henderson Hasselbach for Buffered Solutions pH = pKa + log [conjugate base/acid] pOH = pKb + log [conjugate acid/base] |
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Metal Cations as Weak Acids (not really Bronsted-Lowry realm) |
Cations of small, highly charged metals are weakly acidic
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Classifying Salt solutions as Acidic Basic or Neutral |
- If only one of the ions is from strong ACID or strong BASE, the solution shifts in favor of strong
- if both, solution will be neutral |
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Examples from above slide: Al(NO3)3 forms acidic solution |
Group 1 and 2 make things acidic Transition metals do not form salts |
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Ionization in Polyprotic Acids |
These ionize in steps; Ka1 > Ka2 > Ka3 etc.
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Ionization in H2SO4 *this one is an ODD BIRD |
Ionization for these must be considered with ICE Tables
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Binary Acids |
Element + non-metal,
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Strengths of Oxyacids, H-O-Y |
The more electronegative the Y atom the stronger the oxyacid
More acidic closer to top; less as you move down |
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Lewis-Acid Base theory |
Lone pairs of e- is key
Donating lone pair, Lewis Base Accepting lone pair, Lewis Acid |
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Lewis Base |
MUST have a lone pair |
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Lewis Acid |
MUST seek a lone pair, they are electron defficient examples: H+ has empty 1s orbital B in BF3 has empty 2p orbital and incomplete octet (very strong lewis acid) small charged metal cations, Atoms attached to highly electronegative atoms |
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Lewis Acid-Base reactions |
Examples:
Electron Donor |
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Bronsted lowry |
cannot just have acid, must also have base interacting with acid
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