CHEII:16.0-16.6
H+ ion in Water
(Ionization of HCl= H+ + Cl-) An H+ ion interacts strongly with any sources of electron density, such as the nonbonding electron pairs on the oxygen atoms of water molecules. (Ex hydronium ion: H3O+(aq))
Example
0.035 M solution of HNO2 contains 3.7 x 10-3 M H+(aq) and it's percent Ionization is = H+ equil/ HNO2 initial x 100% = 3.7 x 10-3 M/ 0.035 M x 100%= 11%
Example II
0.10 M solution of formic acid (HCOOH) contains 4.2 x 10-3 M H+(aq). Calc percent acid ionized. = H+ equil / HCOOH initial x 100% = 4.2 x 10-3 M/0.10M x 100% = 4.2
Example calculating measured ph
A student prepared a 0.10 M solution of Formica acid (HCOOH) and found its pH at 25•C to be 2.38. Calculate Ka for formic acid at this temp. First write equilib equation: HCOOH(aq) >< H+ (aq) + HCOO- (aq) Equilibrium constant expression is Ka= [H+][HCOO-]//[HCOOH] From measured pH can calculate [H+]: pH=-log [H+]=2.38 Log[H+]= -2.38 [H+]= 10-2.38=4.2 x 10-3 M Pg 688
Conjugate acid base pairs II
An acid base pair such as HA and A- that differ only in the presence or absence of a proton are called a conjugate acid-base pair Every acid has a conjugate base, formed by removing a proton from the acid. Conversely, every base has a conjugate acid, formed by adding a proton to a base. H3O+ is the conjugate acid of H2O and HA is the conjugate acid of A-
Proton Donor/Acceptor
An acid is a substance (molecule or ion) that donates a proton to another substance (becomes more concentrated/acidic) A base is a substance that accepts a proton (Becomes more dilute, mixture)
Percent Ionization
Another measure of acid strength (like Ka): = concentration of ionized HA/original concentration of HA x 100% The stronger the acid the greater percent Ionization.
Brønsted-Lowry outside of aqueous solutions
Because the emphasis of B-L concept is on proton transfer, the concept also applies to reactions that do not occur in aqueous solution. (Ex: reaction between gas phase HCl and NH3, proton transferred from acid HCl and base NH3).
Calculate the values of [H+] & [OH-] in neutral aqueous solution at 25•C
Determine concentration of values^ Kw= [H3O+][OH-] = [H+][OH-] = 1.0 x 10-14 (neutral equation : [H+]=[OH-] [H+][OH-]= (x)(x) = 1.0 x 10-14 X2= 1.0 x 10-14 X = 10-7 M= [H+]=[OH-] In an acid solution [H+] is greater than 1.0 x 10-7 M; in a basic solution [H+] is less than 1.0 x 10-7 M.
Percent Ionization II
If we assume that the autoionization of H2O is negligible, the concentration of acid that ionizes equals the concentration of H+ (aq) that forms. Thus, the percent Ionization for an acid HA can be expressed as: = [H+]equil/[HA]initial x 100%
Strong Acids II
In an aqueous solution of a strong acid, the acid is normally the only significant source of H+ ions. As a result, calculating the pH of a solution of a strong monoprotic acid is straightforward because H+ equals the original concentration of acid. In 0.20 M solution of HNO3(aq), for example, H+ = NO3- = 0.20 M. The situation with the diprotic acid H2SO4 is somewhat more complex.
Conjugate acid-base pairs intro
In any acid-base equilibrium, both the forward reaction (to the right) and the reverse reaction (to the left) involve proton transfer. HA (aq) + H2O (l) >< A-(aq) + H3O+(aq) I'm forward reaction, HA donates a proton to H2O, therefore HA is a B-L acid and H2O is B-L base. In reverse, the H3O+ ion donates a proton to the A- ion, so H3O+ is the acid and A- is the base. When the acid HA donates a proton, it leaves behind a substance, A-, that can act as a base. Likewise when H2O acts as a base? It generates H3O+ which can act as an acid.
Calculating Ka from pH
In order to calculate either the Ka value for a weak acid or the pH of its solution. Small magnitude of Ka allows to use approx to simplify the problem. Doing the calculations, realize that proton-transfer reactions are generally very rapid. Measured or calculated pH for a weak acid always represents an equilibrium condition.
Proton-Transfer Reactions
In the reaction that occurs when HCl dissolves in water, the HCl molecule transfers an H+ ion (a proton) to a water molecule. Thus reaction represented occurring between an HCl molecule and a water molecule to form a hydronium and chloride ions. Notice proton donor/acceptor (HCl - acid, donates proton to H20) (H2O- base, accepts a proton from HCl) Brønsted-Lowry
Ion product of water
Kw = [H+][OH-] = 1x10^-14 mol/L (25•C )
Weak acids
Most acidic substances are weak acids and therefore only partially ionized in aqueous solution. Use equilibrium constant for the Ionization reaction to express the extent to which a weak acid ionizes. If we represent a general weak acid ie HA, we can write the equation for its Ionization in either of the following ways; Depending on whether the hydrated proton is represent as H3O+(aq) or H+ (aq): HA(aq) + H2O(l) >< H3O+(aq) + A-(aq) Or HA(aq) >< H+(aq) + A-(aq) These equilibria are in aqueous solution so we will use equilibrium - constant expressions based on concentrations. Because H2O is the convent it is omitted from the equilibrium constant expression. Further we have "a" subscript on equilib constant to indicate that it is an equilibrium constant for the Ionization of an ACID. This the equilibrium constant expression as either: Ka=[H3O+][A-]//[HA] Or Ka=[H+][A-]//[HA] Ka is the acid dissociation constant for acid HA.
Strong bases
Most common soluble strong bases are the ionic hydroxides of the alkali metals, such as NaOH, KOH, and the ionic hydroxides heavier alkaline earth metals, such as Sr(OH)2. These compounds completely dissociate into in aqueous solution. Thus, a solution labeled 0.30 M NaOH consist of 0.30 M Na+(aq) and 0.30 M OH-(aq); there is no undissociated NaOH.
autoionization of water
One of the most important chemical properties of water is its ability to act as either a Brønsted-Lowry acid/base. In presence of acid, it acts as a proton acceptor In presence of base, it acts as proton donor.
Relative Strengths of Acids and Bases
Some acids are better proton donors than others and some bases are better proton acceptors than others. If we arrange acids in order of their ability to donate a proton, we find that the more easily a substance gives up a proton, the less easily it's conjugate base accepts a proton. The more easily a base accepts a proton, the less easily its conjugate acid gives up a proton (*the stronger the acid the weaker its conjugate base; the stronger a base, the weaker its conjugate acid)
Acid
Substance that, when dissolved in water, increases the concentration of H+ ions.
Base
Substance that, when dissolved in water, increases the concentration of OH-
Brønsted-Lowry Acids & Bases
The Arrhenius concept of acids and bases, while useful, is rather limited. For one thing, it is restricted to aqueous solutions. However in 1923, chemist Johannes Brønsted & English chemist Thomas Lowry independently proposed a more general definition of acids and bases. This concept is based on the fact that acid-base reactions involve the transfer of H+ ions from one substance to another.
Strong acids and bases
The chemistry of an aqueous solution often depends critically on pH. It is therefore important to examine how pH relates to acid and base concentrations. The simplest cases are those involving strong acids and strong bases. Strong acids and bases are strong electrolyte existing in aqueous solutions entirely as ions.
The pH scale
The molar concentration of H+ (aq) in an aqueous solution is usually very small. Therefore we express [H+] in terms of pH (which is -log in base 10 of [H+]). pH= -log [H+]
pOH and other p scales
The negative log is a convenient way of expressing magnitudes of other small quantities. Use convention that negative logarithm of a quantity is labeled "p" (quantity). Thus express concentration of OH- as pOH: pOH= -log[OH-] Likewise pKw equals -log Kw. By taking the negative logarithm of both sides of the equilibrium-constant expression for water, Kw= [H+][OH-], we obtain: -log[H+] + (-log[OH-]) = -log Kw From which we obtain, usefully; pH+pOH= 14.00 (@25•C) *notice that a change in [H+] by a factor of 10 causes the pH to change by 1. Thus, concentration of H+ (aq) in a solution of pH 5 is 10 times the H+ (aq) concentration in a solution of pH 6.
Strong acids
The seven most common strong acids include six monoprotic acids (HCl, HBr, HI, HNO3, HClO3, and HClO4), and one diprotic acid (H2SO4). Nitric acid NHO3 exemplifies the behavior of the monoprotic strong acids. For all practical purposes, an aqueous solution of HNO3 consists entirely of H3O+ and NO3- ions: HNO3(aq) + H2O(l)>>H3O+ (aq) + NO3- (aq) (complete Ionization) not equilibrium but rather balanced: HNO3 (aq) >>> H+(aq) + NO3-(aq)
Calculate ph of strong base
What is the pH of a) 0.028 M solution of NaOH b) 0.0011 M solution of Ca(OH)2? -calculate pH of two solutions A)[H+]= 1.0 x 10-14 // 0.028 = 3.57 pH=-log(3.57 x 10-13)= 12.45 OR pOH=-log(0.028)=1.55 pH=14.00-1.55 B) **x2 Ca(OH)2's 0.011 M because squared pOH=-log(0.0022)=2.66 pH = 14.00-pOH= 11.34
acid dissociation constant
the ratio of the concentration of the dissociated form of an acid to the undissociated form The magnitude of Ka indicates the tendency of the acid to ionize in water. The larger the value of Ka, the stronger the acid. Chlorine acid for instance is the strongest acid on the table and phenol is the weakest. For most weak acids Ka values range from 10-2 to 10-10