Unit 8: Acids & Bases IB Topics 8 & 18
NOTES - Unit 8: Acids & Bases
PART 1: Acid/Base Theory & Properties
BrØnsted-Lowry: a theory of proton transfer
o A Bronsted-Lowry ACID is a ______.
o A Bronsted-Lowry BASE is a ______.
Conjugate pairs: Acids react to form bases and vice versa. The acid-base pairs related to each other in this way are called conjugate acid-base pairs. They differ by just one proton.
HA + B « A- + BH+
Ex) List the conjugate acid-base pairs in the following reaction: CH3COOH(aq) + H2O(l) « CH3COO-(aq) + H3O+(aq)
Ex) Write the conjugate base for each of the following.
a) H3O+ b) NH3 c) H2CO3
Ex) Write the conjugate acid for each of the following.
a) NO2- b) OH- c) CO32-
Amphoteric / ______substances: substances which can act as Bronsted-Lowry acids and bases, meaning they can accept or donate a proton (capable of both). These features enable them to have a “double-identity:”
1) To act as a Bronsted-Lowry acid, they must be able to dissociate and ______.
2) To act as a B-L base, they must be able to accept H+, which means they must have a lone pair of electrons.
Water is a prime example – it can donate H+ and it has two lone pairs of electrons.
· Auto-ionization of water: H2O + H2O « H3O+ + OH-
· Water reacting as a base with CH3COOH: CH3COOH(aq) + H2O(l) « CH3COO- (aq) + H3O+ (aq)
· Water reacting as an acid with NH3: NH3(aq) + H2O(l) « NH4+(aq) + OH-(aq)
Ex) Write equations to show HCO3- reacting with water with (a) HCO3- as an acid and (b) HCO3- as a base.
Lewis: a theory of electron pairs
o A Lewis ACID is an ______.
o A Lewis BASE is an ______.
Lewis acid-base reactions result in the formation of a covalent bond, which will always be a ______bond (a.k.a. ______bond) because both the electrons come from the base.
Example: (note – the “curly arrow” is a convention used to show donation of electons.)
Example: (note – boron has an incomplete octet, so it is able to accept an electron pair)
Example: (note – metals in the middle of the periodic table often form ions with vacant orbitals in their d subshell, so they are able to act as Lewis acids and accept lone pairs of electrons when they bond with ligands to form complex ions. Ligands, as donors of lone pairs, are therefore acting as Lewis bases)
Typical ligands found in complex ions include H2O, CN- and NH3. Note that they all have lone pairs of electrons, the defining feature of their Lewis base properties.
Table 8.1: Acid-base theory comparison
Bronsted-Lowry / Proton donor / Proton acceptor
Lewis / Electron pair acceptor / Electron pair donor
Ex: For each of the following reactions, identify the Lewis acid and the Lewis base.
a) 4NH3(aq) + Zn2+(aq) « [Zn(NH3)4]2+(aq)
b) 2Cl-(aq) + BeCl2 (aq) + « [BeCl4]2- (aq)
c) Mg2+(aq) + 6H2O(l) « [Mg(H2O)6]2+(aq)
Ex: Which of the following could not act as a ligand in a complex ion of a transition metal?
a) Cl- b) NCl3 PCl3 d) CH4
Properties of acids and bases
For acids and bases here, we will use the following definitions:
· Acid: a substance that donates H+ in solution
· Base: a substance that can neutralize an acid to produce water --- includes metal oxides, hydroxides, ammonia, soluble carbonates (Na2CO3 and K2CO3) and hydrogencarbonates (NaHCO3 and KHCO3)
· Alkali: a soluble base. When dissolved in water, alkalis all release the hydroxide ion, OH-. For example:
o K2O(s) + H2O(l) ® 2K+(aq) + 2OH-(aq)
o NH3(aq) + H2O(l) « NH4+(aq) + OH-(aq)
o CO32- (aq) + H2O(l) « HCO3-(aq) + OH-(aq)
o HCO3-(aq) « CO2(g) + OH-(aq)
· Neutralization: net ionic equation =
Acid-Base Indicators
Acid-Base indicators change color reversibly according to the concentration of H+ ions in solution.
Many indicators are derived from natural substances such as extracts from flower petals and berries. ______, a dye derived from lichens, can distinguish between acids and alkalis, but cannot indicate a particular pH. For this purpose, ______was created by mixing together several indicators; thus universal indicator changes color many times across a range of pH levels.
Table 8.2: Some common acid-base indicators
Litmus / Pink / blue
methyl orange / Red / yellow
phenolphthalein / Colorless / pink
Acids react with metals, bases and carbonates to form salts
1. Neutralization reactions with bases: acid + base ® salt + water
a. With hydroxide bases
b. With metal oxide bases
c. With ammonia (via ammonium hydroxide)
2. With reactive metals (those above copper in the reactivity series): acid + metal ® salt + hydrogen
3. With carbonates (soluble or insoluble) / hydrogencarbonates: acid + carbonate ® salt + water + carbon dioxide
Strong, Concentrated and Corrosive
In everyday English, strong and concentrated are often used interchangeably. In chemistry, they have distinct meanings:
· strong: completely ______into ions
· concentrated: high number of ______of solute per liter (dm3) of solution
· corrosive: chemically ______
Similarly, weak and dilute also have very different chemical meanings:
· weak: only slightly dissociated into ions
· dilute: a low number of moles of solute per liter (dm3) of solution
Strong and weak acids and bases
Consider the acid dissociation reaction: HA(aq) « H+(aq) + A-(aq)
Strong acid: equilibrium lies to the right (acid dissociates fully) ® reversible rxn is negligible ® exists entirely as ions
Ex:
Weak acid: equilibrium lies to the left (partial dissociation) ® exists almost entirely in the undissociated form
Ex:
Similarly, the strength of a base refers to its degree of dissociation in water.
Strong base ex:
Weak base ex:
NOTE: Weak acids and bases are much more common than strong acids and bases.
Table 8.3: Strong and Weak Acids and Bases you should know
Strong Acids(only six; know 1st three for IB) / Strong Bases
(Grp 1 hydroxides & barium hydroxide) / Weak Acids
carboxylic and carbonic acids / Weak Bases
ammonia and amines
H2SO4, sulfuric acid* / LiOH, lithium hydroxide / CH3COOH, ethanoic acid
and other organic acids / C2H5NH2, ethylamine
and other amines
HNO3, nitric acid / NaOH, sodium hydroxide / H2CO3, carbonic acid
Note CO2(aq) = H2CO3(aq) / NH3, ammonia
Note NH3(aq) = NH4OH(aq)
HCl, hydrochloric acid / KOH, potassium hydroxide / H3PO4, phosphoric acid
HI, hydroiodic acid / Ba(OH)2, barium hydroxide
HBr, hydrobromic acid
HClO4, perchloric acid
*NOTE: Sulfuric acid, H2SO4, is a diprotic acid which is strong in the dissociation of the first H+ and weak in the dissociation of the second H+.
For purposes of IB, only monoprotic dissociations are considered.
Experimental methods for distinguishing between strong and weak acids and bases
1. Electrical conductivity: strong acids and bases will have a higher conductivity (higher concentration of mobile ions)
2. Rate of reaction: faster rate of rxn with strong acids (higher concentration of ions)
3. pH: measure of H+ concentration in sol’n. A 1.0 M sol’n of strong acid will have lower pH than 1.0 M sol’n of weak acid; 1.0 M sol’n of strong base will have higher pH than 1.0 M sol’n of weak base
PART 2: pH, pOH & pKw
The pH Scale
· pH is a value chemists use to give a measure of the acidity or alkalinity of a solution.
· Used because [H+] is usually very small
· pH stands for pouvoire of hydrogen.
o Pouvoire is French for “power.”
o The normal range of the pH scale is 0-14.
o However, it is possible (if the hydronium or hydroxide concentrations get above 1 Molar) for the pH to go beyond those values.
· pH= -log[H+]
· [H+] = 10-pH
· As pH decreases, [H+] increases exponentially (a change of one pH unit represents a 10-fold change in [H+]
Example: If the pH of a solution is changed from 3 to 5, deduce how the hydrogen ion concentration changes.
Calculations involving acids and bases
· Sig figs for Logarithms (see page 631): The rule is that the number of decimal places in the log is equal to the number of significant figures in the original number. Another way of saying this is only numbers after decimal in pH are significant.
Example: [H+] = 1.0 x 10-9 M (2 significant figures) ® pH = -log(1.0 x 10-9) = 9.00 (2 decimal places)
· Ion product constant of water, Kw
o Recall that water autoionizes: H2O(l) « H+(aq) + OH-(aq) (endothermic)
o Therefore Kc =
o The concentration of water can be considered to be constant because so little of it ionizes, and it can therefore be combined with Kc to produce a modified equilibrium constant known as kw. In fact, liquids and solids never appear in equilibrium expressions for this reason.
o Therefore, Kw =
o At 25°C, Kw = 1.00 x 10-14
o In pure water, because [H+]=[OH-], it follows that [H+]=
o So at 25°C, [H+] = 1.0 x 10-7, which gives pH = 7.00
· Kw is temperature dependent
o Since the dissociation of water reaction in endothermic (bonds breaking), an increase in temperature will shift the equilibrium to the ______, thus ______the value of Kw.
o As Kw increases, so do the concentrations of H+(aq) and OH-(aq) ® pH decreases
o However, since hydronium and hydroxide concentrations remain equal, water does not become acidic or basic as temperature changes, but the measure of its pH does change.
o Example: Fill in the rest of the table below
Table 8.4: Kw is temperature dependent
Temp (°C) / Kw / [H+] in pure water / pH of pure water0 / 1.5 x 10-15 / 0.39 x 10-7 / 7.47
10 / 3.0 x 10-15
20 / 6.8 x 10-15 / 0.82 x 10-7 / 7.08
25 / 1.0 x 10-14
30 / 1.5 x 10-14 / 1.22 x 10-7 / 6.92
40 / 3.0 x 10-14 / 1.73 x 10-7 / 6.77
50 / 5.5 x 10-14
· H+ and OH- are inversely related
o Because the product [H+] x [OH-] is constant at a given temperature, it follows that as one goes up, the other must go down (since Kw = [H+][OH-])
Table 8.5 Solutions are defined as acidic, basic, or neutral based on the relative concentrations of H+ and OH-
Type of sol’n / Relative concentrations / pH at 25°CAcid
Neutral
Alkaline
o Example: A sample of blood at 25°C has [H+]=4.60 x 10-8 mol dm-3. Calculate the concentration of OH- and state whether the blood is acidic, neutral or basic.
§ How would you expect its pH to be altered at body temperature (37°C)?
· pH and pOH scales are inter-related
o pOH=
· From the relationship: KW = [H+][OH-]
-log KW = -log([H+][OH-])
-log KW = (-log[H+]) + (-log[OH-])
pKW = pH + pOH
· at 25°C, KW = 1.0 x10-14, thus 14.00 = pH + pOH at 25°C
· Given any one of the following we can find the other three: [H+],[OH-],pH and pOH
· Example: Lemon juice has a pH of 2.90 at 25°C. Calculate its [H+],[OH-], and pOH.
PART 3: Weak Acids & Bases
Strong acids and bases: pH and pOH can be deduced from their concentrations: since we assume strong acids and bases dissociate completely, pH and pOH can be calculated directly from the initial concentration of solution.
Example: Calculate the pH of a 0.10 M solution of NaOH at 298 K.
Example: Calculate the pH of a 0.15 M solution of HNO3 at 298 K.
Dissociation constants express the strength of weak acids and bases
Since equilibrium for weak acids and bases lies far to the left (they do not dissociate fully), concentrations of ions in solution cannot be determined by the initial concentrations without knowing the extent of dissociation.
Consider the equilibrium expression for the dissociation of any weak acid in water:
o Ka is known as the acid dissociation constant.
o It has a fixed value for a particular acid at a specified temperature.
o Since the value of Ka depends on the position of equilibrium of acid dissociation, it gives us a direct measure of the strength of an acid.
o The higher the value of Ka at a particular temperature, the greater the dissociation and so the stronger the acid.
o Note: because Ka is an equilibrium constant, its value does not change with the concentration of the acid or in the presence of other ions.
Consider the equilibrium expression for the dissociation of any weak base in water:
Kb is known as the base dissociation constant. It has the same characteristics as those described above for Ka.
Calculations involving Ka and Kb
The values of Ka and Kb enable us to compare the strengths of weak acids and bases and to calculate ion concentrations present at equilibrium (and therefore the pH and pOH values). Keep the following in mind:
o The given concentration of an acid or base is its initial concentration (before dissociation occurs).
o The pH (or pOH) of a solution refers to the concentration of H+ ions (or OH-ions) at equilibrium.
o The concentration values substituted into the expressions for Ka and Kb must be the equilibrium values for all reactants and products.
o When the extent of dissociation is very small (very low value for Ka or Kb) it is appropriate you use the approximations [acid]initial » [acid]equilibrium and [base]initial » [base]equilibrium.
Calculation of Ka and Kb from pH and initial concentration
Example: Calculate Ka at 25°C for a 0.0100 mol dm-3 solution of ethanoic acid, CH3COOH. It has a pH value of 3.40 at this temperature.
Example: Calculate Kb at 25°C for a 0.100 mol dm-3 solution of methylamine, CH3NH2. Its pH value is 11.80 at this temperature.
Calculation of [H+] and pH, [OH-] and pOH from Ka and Kb
A real, but ugly example: Calculate the pH of a 0.10M solution of HNO2 (Ka = 4.0 x 10-4)