Bronsted Acid-Base Reactions PDF

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Joshua Farley Chem 1252L 3/25/2015

Bronsted Acid-Base Reactions Introduction In chemistry, an acid is most often described as a compound that can readily transfer a proton (H+) to an acceptor, known as a base. Acids and bases vary in strength, that is, their ability to dissociate completely in an aqueous solution. Hydrochloric acid, for example, a one of the strongest known acids and dissolves almost completely upon interaction with water. In that reaction, HCl would act as an acid by donating its proton to a water molecule, the base in this case, to produce a hydronium ion (H3O+) and a chloride ion. The hydronium ion becomes known as the conjugate acid to water because it can now donate its extra proton to the conjugate base formed from hydrochloric acid, Cl-. This reversible reaction can be written as: HCl (aq) + H2O (l)  H3O+ (aq) + Cl- (aq) The concentration of hydronium ions in a solution represent the acidity of that solution, while the concentration of hydroxide ions represent the basicity. How acidic a solution is can be determined by finding its pH, the negative logarithm of the concentration of hydronium ions at equilibrium: Eq. 1: The pH of a solution pH = -log [H+]eq; it is important to note that H+ and H3O+ means the same thing in scientific terms. A pH of less than 7 indicates that the solution is acid while a pH higher than 7 indicates a basic solution. A standard pH of 7 represents that a solution is neutral, but is almost never found

Joshua Farley Chem 1252L 3/25/2015

in nature, even in natural water. The absolute strength of an acid is determined by calculating its acid-dissociation constant, Ka, a representation of how much of the acid dissociates. Its value can be found using the following equation: Eq 2: The acid-dissociation constant, Ka, and the base-hydrolysis constant, Kb.

Ka∨b=

[ C ] [ D] [ A ][ B ]

The more positive this value is, the more the acid will dissociate in a solution or the more a base will produce OH-. Acid-base indicators are soluble organic compounds that exist in various colored forms. These forms vary depending on the pH of the solution that it is dissolved in. Indicators can be used to determine the relative pH of a solution, which can then be manipulated to find the equilibrium concentrations and dissociation/hydrolysis constants for each solution. Procedure Forty individual microwells were initially obtained to test numerous samples of each desired solution in order to determine their relative pH values. To do this, three drops of the first substance were placed in five separate microwells. Then, three drops of the second substance were placed in five different microwells. This was repeated for all forty microwells until there were eight sets, each containing the same substances collectively. For the first set of microwells, all of which contained the first substance utilized, one drop of each indicator was used separately among the wells. This was repeated for all eight sets and the final product was eight sets of microwells, each containing 5 individual microwells that had identical starting substances in

Joshua Farley Chem 1252L 3/25/2015

them but all contain a different drop of the indicator. The color of each microwell was observed and compared with the indicator used on it, along with their pH ranges. This allowed for a relative approximation of the pH level of each solution. Table 1: The Bronsted acids used and their conjugate bases Acid Acetic Acid Ammonium

Formula CH3O2H (aq) NH4+

Conjugate Base CH3O2NH3

Table 2: The Bronsted bases used and their conjugate acids Base Acetate Ion Carbonate Ion Sulfate Ion

Formula CH3CO2CO32SO42-

Conjugate Acid CH3CO2H HCO3HSO4-

Table 3: Amphoteric substances used and their conjugates Substance Ammonia Bicarbonate Ion Bisulfate Ion

Formula NH3 HCO3HSO4-

Conjugate Acid NH4+ H2CO3 H2SO4

Conjugate Base NH2CO32SO42-

Table 4: Acid-base indicators used and their pH ranges/colors. Acid-Base Indicator Bromothymol Blue (BTB) Bromocresol Green (BCG) Thymol Blue (TB) Phenolphthalein (PHPH) Alizarin (AZYR)

pH Range/Color Yellow (6.0-7.6) Blue Yellow (3.8-5.4) Blue Red (1.2-2.8) Yellow, Yellow (8.0-9.6) Blue Colorless (8.0-9.8) Magenta Yellow (10.0-12.0) Red

Joshua Farley Chem 1252L 3/25/2015

Analysis Acetic Acid Indicator BTB BCG TB PHPH AZYR pH Range = (2.8-3.8)

Color and associated pH Range Yellow (0.0-6.0) Yellow (0.0-3.8) Yellow (2.8-8.0) Colorless (0.0-8.0) Yellow (0.0-10.0)

Eq. 3: Finding hydronium ion concentration at equilibrium by rearranging equation 1. [H3O+]eq = 10-pH [H3O+]eq Range = (10-3.8, 10-2.8) = (1.6E-4 M, 1.6E-3 M) R CH3CO2H (aq)  I 0.2 M C -x E (0.2 M – x) [H+] = x; x = 1.6E-4 M and 1.6E-3 M

Ka =

1.6E-4 M ¿ ¿ ¿2 ¿ x2 =¿ (0.2 M−x)

Ka =

1.6E-3 M ¿ ¿ ¿2 ¿ 2 x =¿ (0.2 M−x)

CH3CO20 +x x

H+ 0 +x (1.6E-4 M, 1.6E-3 M)

Joshua Farley Chem 1252L 3/25/2015

Ka Range = (1.3E-7, 1.3E-5)

Ammonium Indicator BTB BCG TB PHPH AZYR pH Range = (6.0-7.6)

Color and associated pH Range Green (6.0-7.6) Blue (5.4-14.0) Yellow (2.8-8.0) Colorless (0.0-8.0) Yellow (0.0-10.0)

[H3O+]eq = (10-7.6, 10-6.0) = (2.5E-8 M, 1.0E-6 M)

R NH4+ (aq)  I 0.2 M C -x E (0.2 M – x) [H+] = x; x = 2.5E-8 M and 1.0E-6 M

Ka =

2.5 E−8 M ¿ ¿ ¿2 ¿ 2 x =¿ (0.2 M−x)

Ka =

1.0 E−6 M ¿ ¿ ¿2 ¿ x2 =¿ (0.2 M−x)

Ka Range = (3.1E-15, 5.0E-12)

NH3 0 +x X

H+ 0 +x (2.5E-8 M, 1.0E-6 M)

Joshua Farley Chem 1252L 3/25/2015

Acetate Ion Indicator BTB BCG TB PHPH AZYR pH Range = (5.4-6.0)

Color and associated pH Range Yellow (0.0-6.0) Blue (5.4-14.0) Yellow (2.8-8.0) Colorless (0.0-8.0) Yellow (0.0-10.0)

[H3O+]eq = (10-6.0, 10-5.4) = (1.0E-6 M, 3.9E-6) 1.0E-14 1.0E-14 , [OH-]eq = 3.9E-6 1.0E-6 ) = (2.6E-9 M, 1.0E-8 M) ¿ R CH3CO2- (aq)  I 0.2 M C -x E (0.2 M – x) [OH-] = x; x = 2.6E-9 M and 1.0E-8 M 2.6E-9 M ¿ ¿ 2 ¿ Kb = ¿ x2 =¿ (0.2 M−x) 1.0E-8 M ¿ ¿ 2 ¿ Kb = ¿ x2 =¿ (0.2 M−x) Kb Range = (3.4E-17, 5.0E-16)

CH3CO2H 0 +x X

OH0 +x (2.6E-9 M, 1.0E-8 M)

Joshua Farley Chem 1252L 3/25/2015

Carbonate Ion Indicator BTB BCG TB PHPH AZYR pH Range = (12.0-14.0)

Color and associated pH Range Blue (7.6-14.0) Blue (5.4-14.0) Blue (9.6-14.0) Magenta (9.8-14.0) Red (12.0-14.0)

[H3O+]eq = (10-14.0, 10-12.0) = (1.0E-14 M, 1.0E-12 M) 1.0E-14 1.0E-14 , [OH-]eq = 1.0 E−12 1.0E− 14 ¿ R CO32- (aq)  I 0.2 M C -x E (0.2 M – x) [OH-] = x; x = 0.010 M and 1.0 M 0.010 M ¿ ¿ 2 ¿ Kb = ¿ x2 =¿ (0.2 M−x) 1.0 M ¿ ¿ 2 ¿ Kb = ¿ x2 =¿ (0.2 M−x) Kb Range = (5.3E-4, 1.25)

) = (0.010 M, 1.0 M)

HCO30 +x X

OH0 +x (0.010 M, 1.0 M)

Joshua Farley Chem 1252L 3/25/2015

Sulfate Ion Indicator BTB BCG TB PHPH AZYR pH Range = (7.6-8.0)

Color and associated pH Range Blue (7.6-14.0) Blue (5.4-14.0) Yellow (2.8-8.0) Colorless (0.0-8.0) Yellow (0.0-10.0)

[H3O+]eq = (10-8.0, 10-7.6) = (1.0E-8 M, 2.5E-8 M) 1.0E-14 1.0E-14 , [OH-]eq = 2.5 E−8 1.0E−8 ) = (4.0E-7 M, 1.0E-6 M) ¿ R SO42- (aq)  I 0.2 M C -x E (0.2 M – x) [OH-] = x; x = 4.0E-7 M and 1.0E-6 M 4.0 E−7 M ¿ ¿ 2 ¿ Kb = ¿ x2 =¿ (0.2 M−x) 1.0E−6 M ¿ ¿ 2 ¿ Kb = ¿ x2 =¿ (0.2 M−x) Kb Range = (8.0E-13, 5.0E-12)

HSO40 +x X

OH0 +x (4.0E-7 M, 1.0E-6 M)

Joshua Farley Chem 1252L 3/25/2015

Ammonia Indicator BTB BCG TB PHPH AZYR pH Range = (12.0-14.0)

Color and associated pH Range Blue (7.6-14.0) Blue (5.4-14.0) Blue (9.6-14.0) Magenta (9.8-14.0) Red (12.0-14.0)

[H3O+]eq = (10-14.0, 10-12.0) = (1.0E-14 M, 1.0E-12 M)

Bicarbonate Ion Indicator BTB BCG TB PHPH AZYR pH Range = (8.0-9.6)

Color and associated pH Range Blue (7.6-14.0) Blue (5.4-14.0) Green (8.0-9.6) Light Pink (8.0-9.8) Yellow (0.0-10.0)

[H3O+]eq = (10-9.6, 10-8.0) = (2.5E-10 M, 1.0E-8 M) Bisulfate Ion Indicator BTB BCG TB PHPH AZYR pH Range = (1.2-2.8) [H3O+]eq = (10-2.8, 10-1.2) = (1.6E-3 M, 6.3E-2 M)

Color and associated pH Range Yellow (0.0-6.0) Yellow (0.0-3.8) Orange (1.2-2.8) Colorless (0.0-8.0) Yellow (0.0-10.0)

Joshua Farley Chem 1252L 3/25/2015

Discussion Based on the experimental data, the two acids in table one can be ranked from weakest to strongest by comparing the pH ranges and hydronium ion concentrations at equilibrium. The 0.2 M acetic acid solution revealed a pH of between 2.8 and 3.8, while the ammonium solution had a range between 6.0 and 7.6. Because the pH of acetic acid is lower, its hydronium ion concentration is going to be higher than that of the ammonium solution, making it more acidic. Similarly, the three bases in table 2 can be ranked from weakest to strongest base by comparing their pH values. A solution with a high pH (i.e. a pH above 7.0) is considered a basic solution and the higher the pH, the lower the hydronium ion concentration. In contrast, a lower hydronium ion concentration at equilibrium indicates a high hydroxide ion concentration, indicating a basic solution. The acetate ion solution exhibited a more acidic concentration than the other two bases and had a pH range between 5.4 and 6.0, making it the weakest base of the list. The sulfate ion solution ranks as the second strongest base from the table because of its pH range between 7.6 and 8.0. The strongest base tested was the carbonate ion, which revealed a maximum pH range of 12.0-14.0, meaning that very few hydronium ions (possibly a negligible amount) were present in the solution. In order to understand exactly what is occurring in a solution, we must examine the dissociation and hydrolysis reactions associated with each of the substances. Because these reactions are reversible, we can use the calculated Ka or Kb values to determine which side of the reaction is favored. If the constant is greater than one then the products are favored, but if the constant is less than one then more reactants will be present than products. The carbonate ion was the only substance with a potential K constant higher than one. Based on that information,

Joshua Farley Chem 1252L 3/25/2015

none of the acids or bases analyzed were strong acids or strong bases. The acid-dissociation reactions from the substances in table 1 are as follows: Acetic Acid: CH3CO2H (aq) + H2O (l)  CH 3CO2- (aq) + H3O+ (aq) Ammonium: NH4+ (aq) + H2O (l)  NH3 (aq) + H3O+ (aq)

And the base-hydrolysis reactions of table 2 are: Acetate Ion: CH3CO2- (aq) + H2O (l)  CH3CO2H (aq) + OH- (aq) Carbonate Ion: CO32- (aq) + H2O (l)  HCO3- (aq) + OH- (aq) Sulfate Ion: SO42- (aq) + H2O (l)  HSO4- (aq) + OH- (aq) The substances in table 3 are amphoteric substances, meaning that they can act as Bronsted bases by accepting protons from the solvent (water), or they can act as Bronsted acids by donating their protons. Based on the pH values and hydronium ion concentrations calculated, the amphoteric substances can be ranked in order of increasing acidity: Ammonia (pH = 12.0-14.0)  Bicarbonate Ion (pH = 8.0-9.6)  Bisulfate Ion (pH = 1.2-2.8) Or increasing basicity: Bisulfate Ion (pH = 1.2-2.8)  Bicarbonate Ion (pH = 8.0-9.6)  Ammonia (pH = 12.0-14.0) The lower pH values represent a more acidic solution, meaning that bisulfate ions almost always donate their protons rather than accepting them to form hydroxide. Because more hydronium ions are made in a bisulfate ion solution, it can be implied that its K a value is going to

Joshua Farley Chem 1252L 3/25/2015

exceed its Kb value, but will still not be greater than one because it is both a weak acid and a weak base. Ammonia produced a very basic solution containing little to no hydronium ions. Ammonia has a greater tendency to accept protons for the lone pair of electrons on the central nitrogen atom, therefore forming ammonium (NH4+). Because ammonia was found to have the same pH range as the carbonate ion (12.0-14.0), it can be concluded that ammonia also has a Kb value range of between 5.3E-4 and 1.25. Finally, the bicarbonate ion solution produced a slightly basic solution, meaning that it has a greater Kb value due to its tendency to act as a base. Despite this, bicarbonate could still act primarily as an acid when mixed into a very basic solution to create a state of equilibrium. Table 5: Actual Ka and Kb values for compounds in tables 1 and 2 at 25.0 degrees Celsius; acquired from google. Substance Acetic Acid Ammonium Acetate Ion Carbonate Ion Sulfate Ion

Actual Ka / Kb Value Ka = 1.8E-5 Ka = 5.5E-10 Kb = 5.6E-10 Kb = 1.8E-4 Kb = 8.3E-13

Calculated Constant Range (1.3E-7, 1.3E-5) (3.1E-15, 5.0E-12) (3.4E-17, 5.0E-16) (5.3E-4, 1.25) (8.0E-13, 5.0E-12)

The acetic acid, carbonate ion solution, and sulfate ion solution tests were all relatively accurate, but there were obvious discrepancies in the ammonium ion and acetate ion tests. The calculated Kb range for the acetate ion came out to be way lower, almost a negligible amount, in comparison to the actual value. This may have been caused by a problem in pH observations, where the color of the solution didn’t exactly match the colors on the pH scale, causing us to have to make an estimation. The ammonium values are much closer, indicating more minor discrepancies that may have been caused by small factors such as misreadings and excess acid-

Joshua Farley Chem 1252L 3/25/2015

base indicator in the solutions. Overall, the experiment was successful in determining the pH, hydronium/hydroxide concentrations, and the acid-dissociation/base-hydrolysis constants consistent with each of the acids and bases tested. Conclusion Acids and bases are an important part of our everyday lives. We drink them, we avoid them, our body uses them for biochemical pathways, they help us perform tasks and create new substances. This experiment exhibited the basic relationships between the well-known pH scale and the hydronium or hydroxide concentrations of a solution. The hydronium and hydroxide concentrations are inversely related, that is, if a solution contains a large amount of hydronium ions then it will have very little hydroxide ions and be very acidic. Using calculated equilibrium concentrations acquired from an ICE table, the acid-dissociation constant can be found to determine if the reaction favors the reactants or products sides (shows how much the acid dissociates from its proton). Similarly, base-hydrolysis constants can be found by first converting the concentration of hydronium ions to the concentration of hydroxide ions—[OH-][H3O+] = 1.0E-14. An ICE table can then made created in the same fashion to find its constant and determine how much the base will accept protons. As seen in the analysis section, all of the acids and bases tested had constants less than one, signifying that they are all weak in terms of the relative strength of acids and bases. These concepts help us in understanding everyday phenomenon, such as why blowing into a straw in a cup of water will make it more acidic (the formation of carbonic acid) or why hydrochloric acid can burn straight through skin and bone....