Acids and Bases
What are acids? • Why are they important? • What are bases? • Why are they important? • Which is dangerous? Acids or Bases?
16.1 acids and bases • acids from the Latin acidus (sour) • bases (a.k.a. alkalis) are known for their bitter tastes, and slippery feel; most soaps are basic
the arrhenius model • Svante Arrhenius • acids produce H+ ions in aqueous solutions • bases produce OH- (hydroxide) ions in aqueous solutions
16.1 Example • HCl(g)H+(aq) + Cl-(aq) or • HCl(g) +H2O H3O+(aq) + Cl-(aq)this is a strong acid • NaOH(s) Na+(aq) + OH-(aq)this is a strong base
the Brønsted-Lowry model • The Arrhenius concept is limited because the only base is hydroxide (OH-) • Bronsted and Lowry both formulated a better explanation • acids donate protons (H+) • bases accept protons (H+)
16.1 • assume our acid is HA; then… • HA + H2O H3O+ + A- • The HA lost an H+ (proton); it acted acidic • The water gained a proton; it acted as a base
16.1 • when an acid loses an H+ it becomes a conjugate base • when a base gains an H+ it becomes a conjugate acid
+ + + H H O H O Cl Cl H H H The Bronsted-Lowry concept • In this idea, the ionization of an acid by water is just one example of an acid-base reaction. conjugate acid conjugate base acid base conjugate acid-base pairs • Acids and bases are identified based on whether they donate or accept H+. • “Conjugate” acids and bases are found on the products side of the equation. A conjugate base is the same as the starting acid minus H+.
16.1 • a conjugate acid-base pair differs just by one single H+: • HCl/Cl- • H2O/H3O+ • NH4+/NH3
Conjugate Acid base pairs Identify the acid, base, conjugate acid, conjugate base, and conjugate acid-base pairs: HC2H3O2(aq) + H2O(l) C2H3O2–(aq) + H3O+(aq) conjugate base conjugate acid acid base conjugate acid-base pairs OH–(aq) + HCO3–(aq) CO32–(aq) + H2O(l) base acid conjugate base conjugate acid conjugate acid-base pairs
Practice HF(aq) + SO32–(aq) F–(aq) + HSO3–(aq) (a) conjugate base conjugate acid acid base conjugate acid-base pairs (b) CO32–(aq)+HC2H3O2(aq)C2H3O2–(aq)+HCO3–(aq) base acid conjugate base conjugate acid conjugate acid-base pairs (c) H3PO4(aq) + OCl–(aq) H2PO4–(aq) + HOCl(aq) conjugate base conjugate acid acid base conjugate acid-base pairs
16.2 acid strength • Find the Conjugate acid or base of the following • Find the acid of NO3- • Find the base ofHCO3-
16.2 acid strength • strong acids give up every H they can, they are completely ionized • HCl H+ + Cl- • weak acids hold on to most of their H, they are only partially ionized. • H2CO3 H+ + HCO3-
The strong acids are : • HCl, HBr, HI, HNO3, H2SO4, HClO4 • All other acids are weak • Ex. H3PO4 is weak because it does not completely ionize and it is not one of the 6 strong acids.
16.2 acid strength strong basescompletely ionize and form OH- Strong bases include group 1(Alkali metal) hydroxides, Sr(OH)2, Ca(OH)2 and Ba(OH)2. All others are weak. Strong NaOH Weak CuOH NaOHNa+ + OH- CuOHCu+ + OH- Weak bases do not completely ionize.
16.2 strong weak
16.2 strong weak
Identify if the following are strong or weak acids/bases and state whether a light bulb will be bright, dim or dark. • NaOH • HC2H3O2 • HF • HCl • H2C2O4 • Fe(OH)2 • CH3CH2OH
16.2 acid ionization flash.swf
16.2 • diprotic acids have 2 protons that can be lost (e.g. H2SO4) • H3PO4 is triprotic. • Oxyacids: the proton to be lost is hooked up to an O (e.g. HNO3, HOCl, H2SO4, H3PO4) The more O’s in the formula, the stronger the acid.
16.2 • organic acids carbon-containing chemicals with acid properties; most common are carboxylic acids –C–O–H || O • they are weak (e.g. acetic acid (vinegar), formic acid (ant bites), citric acid, vitamin C, aspirin)
16.3 water as an acid and a base • Water is amphoteric (acts as both acid and base) • It can accept a proton (to become hydronium), or • donate a proton (to become hydroxide ion)
16.3 H2O + H2O H3O+ + OH– • Because of this even purewater has someacids and basesin it
How ‘acid’ or ‘basic’ is it? Acidic Solutions = Hydronium ion concentration is greater than hydroxide ion concentration Basic Solutions = Hydroxide concentration is greater than hydronium ion concentration
16.3 • a solution is acidic if [H3O+] > [OH-] • a solution is basic if [H3O+] < [OH-] • a solution is neutral if [H3O+] = [OH-] • The brackets [ ] mean concentration What are the units for concentration? M, moles/liter
16.4 the pH scale • What does pH mean? • the pH scale is an easy way to see how acidic or basic a solution is • mathematically: pH = –log[H3O+] • A pH of 7 is 10-7 mol/liter or 10-7 M H3O+ • if you know the [H3O+] you can figure out the pH of the solution by: entering the [H3O+], taking the log, and changing the sign (+/-)
example • what is the pH if the [H3O+] = 0.001 M • pH = 3 • pH if the [H3O+] = 0.000021 M • pH = 4.68 • pH if the [H3O+] = 0.000000059 M • pH = 7.23
16.4 • every change of one pH unit = 10x acidic concentration change • pH > 7 basic • pH = 7 neutral • pH < 7 acidic
16.4 • to find the [H3O+] from pH, do the opposite • enter pH; change sign; take inverse log (10^…) • so if pH is 8.53 • change sign (-8.53) • take inv log 10^(-8.53) = (3.0 x 10-9 M) = 0.000000003 M
Water • Water ionizes- falls apart into ions. • H2O H+ + OH- • Called the self ionization of water. • Only a small amount. • [H+ ] = [OH-] = 1 x 10-7M • A neutral solution. • In water Kw = [H+ ] x [OH-] = 1 x 10-14 • Kw is called the ion product constant.
Ion Product Constant • H2O H+ + OH- • Kw is constant in every aqueous. solution [H+] x [OH-] = 1 x 10-14 M2 • If [H+] > 10-7 then [OH-] < 10-7 • If [H+] < 10-7 then [OH-] > 10-7 • If we know one, we can determine the other. • If [H+] > 10-7 acidic [OH-] < 10-7 • If [H+] < 10-7 basic [OH-] > 10-7
pH and pOH These are the most useful…. • pH = - log [H+] • pOH = - log [OH-] • pH + pOH = 14 • [H+] x [OH-] = 1.0 x 10 -14
pH and pOH Calculations pH + pOH = 14 [H+] [OH-] = 1 x 10-14
1. If the pOH of a solution is 5.25, what is pH? 2. What is the pH of a solution with a pOH of 4? 3. What is the [H+] of a solution with a [OH-] of 1.0 x10 – 4 4. If the pH is 2.7, what is the [H+] ? 5. What is the [OH-] of a solution with a pH of 10? = 5. x 10 -6
Strong acids have the same [H+] as the acid (so do bases with [OH-] • Calculate the pH of a 0.25 M HCl solution. • Calculate the pH of a 0.10 M NaOH solution. So [H+]= 0.25 M HCl H+ + Cl- pH= - Log (0.25) = 0.60 So [OH-] = 0.10 M NaOH Na+ + OH- pOH= - Log (0.10) = 1, so pH = 14 -1 = 13
16.5 measuring pH pH can be measured in several ways one way involves indicators, things which change colors in the presence of different H3O+ concentrations
16.5 some indicators can be put on paper, we can read colors to determine pH
16.5 a really accurate way of doing it is with a pH meter
Neutralization Reactions • Acid + Base Salt + water • Salt = an ionic compound • Water = HOH • HNO3 + KOH • HCl + Mg(OH)2 • H2SO4 + NaOH • Really just double replacement.
Titration • When you add the same number of moles of acid and base, the solution is neutral. • By measuring the amount of a base added you can determine the concentration of the acid. • If you know the concentration of the base. • This is a titration.
Titration equations • Ma x Va x # of H+ = Mb x Vb x # of OH- moles of H+ = moles of OH-
More Practice • If it takes 45 mL of a 1.0 M NaOH solution to neutralize 57 mL of HCl, what is the concentration of the HCl ? • If it takes 67 mL of 0.500 M HF to neutralize 15mL of NaOH what was the concentration of the NaOH ? • How much of a 0.275 M HCl will be needed to neutralize 25mL of .154 M NaOH?
16.8 buffered solutions a buffered soln resists a change in pH even when a strong acid or base is put in it especially important to living critters whose bio systems require a cnst pH how does it work?
16.8 a soln contains a weak acid HA, and its conjugate base A- when we dump in extra acid or base these two gobble it up: the A- reacts with the xs protons to become HA the HA reacts with the xs OH- to become A- overall there is little change in pH!!!