Acids and Bases (courtesy of L. Scheffler, Lincoln High School, 2010)
Acids • React with certain metals to produce hydrogen gas. • React with carbonates and bicarbonates to produce carbon dioxide gas Bases • Have a bitter taste • Feel slippery. • Many soaps contain bases.
Properties of Acids • Produce H+ (as H3O+) ions in water (the hydronium ion is a hydrogen ion attached to a water molecule) • Taste sour • Corrode metals • Good Electrolytes • React with bases to form a salt and water • pH is less than 7 • Turns blue litmus paper to red “Blue to Red A-CID”
Some Common Acids HC2H3O2 acetic acid in vinegar HCl hydrochloric acid stomach acid H3C6H5O7 citric acid fruits H2CO2 carbonic acid soft drinks H32PO4 phosphoric acid soft drinks
Properties of Bases • Generally produce OH- ions in water • Taste bitter, chalky • Are electrolytes • Feel soapy, slippery • React with acids to form salts and water • pH greater than 7 • Turns red litmus paper to blue “BasicBlue”
Some Common Bases NaOH sodium hydroxide lye KOH potassium hydroxide liquid soap Ba(OH)2 barium hydroxide stabilizer for plastics Mg(OH)2 magnesium hydroxide “MOM” Milk of magnesia Al(OH)3 aluminum hydroxide Maalox (antacid)
Arrhenius Definition Arrhenius Acid - Substances in water that increase the concentration of hydrogen ions ((H+ or hydronium ions H3O+). Base - Substances in water that increase concentration of hydroxide ions (OH-). Categorical definition – easy to sort substances into acids and bases Problem – many bases do not actually contain hydroxides
Bronsted-Lowry Definition Acid - substance that donates a proton. Base - substance that accepts a proton. HA + B HB+ + A- Ex HCl + H2O H3O+ + Cl- Acid Base Conj Acid Conj Base • A “proton” is really just a hydrogen atom that has lost it’s electron! • The classification depends on how the substance behaves in a chemical reaction
Conjugate Acid Base Pairs • Conjugate Base - The species remaining after an acid has transferred its proton. • Conjugate Acid - The species produced after base has accepted a proton. • HA & A- - conjugate acid/base pair • A- - conjugate base of acid HA • B & HB+ - conjugate acid/base pair • HB+ - conjugate acid of base :B
A Brønsted-Lowry acidis a proton donor A Brønsted-Lowry baseis a proton acceptor conjugatebase conjugateacid base acid
Examples of Bronsted-Lowry Acid Base Systems • Note: Water can act as acid or base • Acid Base Conjugate AcidConjugate Base • HCl + H2O H3O+ + Cl- • H2PO4- + H2O H3O+ +HPO42- • NH4+ + H2O H3O+ + NH3
Lewis Definition • Lewis • Acid - an electron pair acceptor • Base - an electron pair donor
Brønsted-Lowry vs. Lewis • All B/L bases are Lewis bases BUT, by definition, a B/L base cannot donate its electrons to anything but a proton (H+) • While B/L is most useful for our purposes, Lewis allows us to treat a wider variety of reactions (even if no H+ transfer occurs) as A/B reactions
Acid Strength • Strong Acid - Transfers all of its protons to water; - Completely ionized; - Strong electrolyte; - The conjugate base is weaker and has a negligible tendency to be protonated. • Weak Acid - Transfers only a fraction of its protons to water; • - Partly ionized; - Weak electrolyte; - The conjugate base is stronger, readily accepting protons from water • As acid strength decreases, base strength increases. • The stronger the acid, the weaker its conjugate base • The weaker the acid, the stronger its conjugate base
Acid Dissociation Constants Dissociation constants for some weak acids
Base Strength • Strong Base - all molecules accept a proton; - completely ionizes; - strong electrolyte; - conjugate acid is very weak, negligible tendency to donate protons. • Weak Base - fraction of molecules accept proton; - partly ionized; - weak electrolyte; - the conjugate acid is stronger. It more readily donates protons. • As base strength decreases, acid strength increases. • The stronger the base, the weaker its conjugate acid. • The weaker the base the stronger its conjugate acid.
Common Strong Acids/Bases Strong Acids Hydrochloric Acid Nitric Acid Sulfuric Acid Perchloric Acid Strong Bases Sodium Hydroxide Potassium Hydroxide *Barium Hydroxide *Calcium Hydroxide *While strong bases they are not very soluble
A/B Behavior & Chemical Structure • Binary Acids • Hydrogen and another element • Polyprotic Acids • Have more than 1 Hydrogen to give away • Oxyacids • have O in compound • Carboxylic Acids • have –COOH in compound
Wait, water can go both ways? • amphoteric substances can behave as either an acid or base depending on what they react with. • water and anions with protons (H+) attached are the most common amphoteric substances
Autoionization of Water H2O + H2O OH- + H3O+ @ 25 oC the concentrations for both [H3O+] and [OH-] = 1.00 x 10-7 and [H3O+] [OH-] = 1.00 x 10-14 = Kw
Since [H3O+] [OH-] = 1.00 x 10-14 = Kw • when [H3O+]=[OH-] the solution is neutral • when [H3O+]>[OH-] the solution is acidic • when [H3O+]<[OH-] the solution is basic
The pH scale is a way of expressing the strength of acids and bases. Instead of using very small numbers, we just use the NEGATIVE power of 10 on the Molarity of the H+ (or OH-) ion.Under 7 = acid 7 = neutral Over 7 = base
pH calculations – Solving for H+ If the pH of Coke is 3.12, [H+] = ??? Because pH = - log [H+] then - pH = log [H+] Take antilog (10x) of both sides and get 10-pH =[H+] [H+] = 10-3.12 = 7.6 x 10-4 M *** to find antilog on your calculator, look for “Shift” or “2nd function” and then the log button
Calculating the pH pH = - log [H+] (Remember that the [ ] mean Molarity) Example: If [H+] = 1 X 10-10pH = - log 1 X 10-10 pH = - (- 10) pH = 10 Example: If [H+] = 1.8 X 10-5pH = - log 1.8 X 10-5 pH = - (- 4.74) pH = 4.74
pH calculations – Solving for H+ • A solution has a pH of 8.5. What is the Molarity of hydrogen ions in the solution? pH = - log [H+] 8.5 = - log [H+] -8.5 = log [H+] Antilog -8.5 = antilog (log [H+]) 10-8.5 = [H+] 3 X 10-9 = [H+]
pOH • Since acids and bases are opposites, pH and pOH are opposites! • pOH does not really exist, but it is useful for changing bases to pH. • pOH looks at the perspective of a base pOH = - log [OH-] • Since pH and pOH are on opposite ends pH + pOH = 14
The pH Scale pH [H3O+ ] [OH- ] pOH
pH testing • There are several ways to test pH • Blue litmus paper (red = acid) • Red litmus paper (blue = basic) • pH paper (multi-colored) • pH meter (7 is neutral, <7 acid, >7 base) • Universal indicator (multi-colored) • Indicators like phenolphthalein • Natural indicators like red cabbage, radishes
pH indicators • Indicators are dyes that can be added that will change color in the presence of an acid or base. • Some indicators only work in a specific range of pH • Once the drops are added, the sample is ruined • Some dyes are natural, like radish skin or red cabbage
pH and acidity The pH values of several common substances are shown at the right. Many common foods are weak acids Some medicines and many household cleaners are bases.
Titration & Titration Curves • Titration: the adding of one solution of an known concentration into another solution • standard solution: a solution with a known concentration • Titration curve: a graph showing pH vs volume of acid or base added • The pH shows a sudden change near the equivalence point • The Equivalence point (a.k.a. stoichiometric point) is the point at which the moles of OH- are equal to the moles of H3O+
pH cm3 base added Strong acid-strong base Titration Curve • At equivalence point, Veq: Moles of H3O+ = Moles of OH- • There is a sharp rise in the pH as one approaches the equivalence point • With a strong acid and a strong base, the equivalence point is at pH =7 • Neither the conjugate base or conjugate acid is strong enough to affect the pH
Weak acid-strong base Titration Curve • The increase in pH is more gradual as one approaches the equivalence point • With a weak acid and a strong base, the equivalence point is higher than pH = 7 • The strength of the conjugate base of the weak acid is strong enough to affect the pH of the equivalence point
Buffered Weak Acid-Strong Base Titration Curve • The initial pH is higher than the unbuffered acid • As with a weak acid and a strong base, the equivalence point for a buffered weak acid is higher than pH =7 • The conjugate base is strong enough to affect the pH
Neutralization An acid will neutralize a base, giving a salt and water as products Examples Acid Base Salt water HCl + NaOH NaCl + H2O H2SO4 + 2 NaOH Na2SO4 + 2 H2O H3PO4 + 3 KOH K3PO4 + 3 H2O 2 HCl + Ca(OH) 2 CaCl2 + 2 H2O A salt is an ionic compound that is formed from thepositive ion (cation)of thebaseand thenegative ion (anion) of the acid
Acid-Base Reactions 0 • Acid: • produces hydrogen ions (H+) in solution (Arrhenius) • proton donor (Lewis) • Base: • produces hydroxide ions (OH-) in solution (Arrhenius) • proton acceptor (Lewis) • The reaction ALWAYS forms water and an ionic compound(mostly aqueous, known as a salt). • The actual definition of a salt is the ionic product of an acid-base neutralization reaction)
Acid-Base Neutralization • Chemically the reaction looks like this: Acid + Base Salt + Water • A classic example: HCl + NaOH NaCl + H2O Hydrochloric Acid Water Sodium Hydroxide (Lye) Sodium Chloride (Table Salt)
Acid-Base Reactions 0 • Example HNO3(aq)+ KOH (aq) Molecular: HNO3(aq)+ KOH (aq) Total Ionic: H+(aq) + NO3- (aq) + K+ (aq) + OH-(aq) H2O (l) + K+(aq) + NO3- (aq) Net Ionic: H+(aq) + OH-(aq) H2O (l) H2O (l) + KNO3 (aq)
Acid-Base Neutralization • Here’s the equation again: • HCl + NaOH NaCl + H2O • Chemically, this is a double replacement reaction: • The H traded its Cl for an OH • The Na traded its OH for a Cl.
Buffer Solutions - Characteristics • A solution that resists a change in pH. It is pH stable. • A weak acid and its conjugate base form an acid buffer. • A weak base and its conjugate acid form a base buffer. • We can make a buffer of any pH by varying the concentrations of the acid/base and its conjugate.