IONIZATION CONSTANT OF WEAK ACIDS
The two important variables for determining the strength of acids isthe total amount of hydrogen ions released and the ionizationconstant. Weak acids dissociate partially in solution to yield fewhydrogen ions per unit time and therefore, have a low rate ofreaction. The dissociation Ka constant determine the rateof dissociation of an acid. This experiment focused on determiningKa using titration curves and experimental values ofnormal titrations. The Ka obtained using the two methodswere the same.
An acid is asubstance that dissociates in water and releases hydrogen ions,protons. Weak acids are substances that dissociate partially insolution. The hydrogen yielded from the dissociation of weak acids islower compared to those produced by strong acids which completelydissociate. The PH of weak acids is close to the PH 7, from 4 to 6.9.Weak acids have a poor conductivity of electric current due to therelatively small number of mobile ions, hydrogen ion, available toconduct. The rate of reaction of the acid with metals, alkalis, metaloxides, metal carbonates and other chemicals increases withconcentrations of hydrogen ions. Therefore, weak acids have lowerreaction rates than strong acids (Hamid, Q., Shannon, J., &Martin, J. 2005).
The two importantvariables for determining the strength of acids is the total amountof hydrogen ions released and the pKa of the reaction. pKais used to tell how much an acid has dissociated to at aparticular PH. The ionization constant, Ka, is thequantitative measure of the strength of an acid.
HA(aq) +H2O(l)↔H3O+(aq)+ A–(aq)
The position ofequilibrium of the equations is determined by the nature of the acidHA. The acid (HA) and the base (A–) form a conjugateacid-base pair. The acid donates its protons to the base (H2O)in the forward reaction while the acid (H3O+)donates a proton to the base (A–) in the reverse reaction(Stanton et al. 2010). The mathematical equation for theequilibrium constant is as follows
The values of Kaspan a wide range and therefore, they are better accommodated by alogarithmic scale defined in a similar manner to PH:
The pKais inversely proportional to Ka. The acid with greaterionization has a high rate of reaction with bases and salts, andtherefore, the endpoint of the reaction is reached within a shortperiod of time (Clugston, M., & Flemming, R. 2000). The reactionbetween the weak acid and base is complete when it reaches theequivalence point. Midway to the endpoint, half of the acid hasreacted to form the conjugate base A– and H2Oand the concentration of the acid and base are equal. Sodiumhydroxide is used to determine moles of the acids that have reactedand thus helps in knowing the dissociation of the acid.
Preparation and standardization of sodium hydroxide
0.5g of KHP wasmeasured and dispensed in a dry 125 ml Erlenmeyer flask. The KHP wasdissolved in 50 ml of deionized water, and two drops of indicatorwere added. It was titrated with NaOH solution from the burette untilthe endpoint. The amount of NaOH used in the titration was recordedin a table. The above procedure was repeated until three reliabletrials were achieved. The data was used to calculate the molarity ofthe base.
Determination of themolarity and pH of the acetic acid solution
10 ml of 150 mlacetic acid of unknown molarity was transferred into a beaker.Two-point calibration of the PH meter was done. After calibration,the PH of the CH3COOH solution was measured and the [H3O+]determined. A pipette was used to transfer 25.00 ml of CH3COOH into aconical flask, and 2 drops of the indicator added. The standardizedsodium hydroxide was titrated against the acetic acid to theendpoint. The procedure was repeated until three reliable trials wereobtained. The results were used to calculate the molarity of theacetic acid and ionization constant.
Two-pointcalibration of the PH electrode was performed using PH=4 and PH=7buffers. The PH sensor was calibrated. A reservoir was set up,calibrated and filled with the standardized NaOH solution. Airbubbles were eliminated by flushing NaOH into a waste beaker. Thedrop counter was also calibrated in order to record precise volume oftitrant in ml. The reservoir was refilled with NaOH solution and25.00 ml CH3COOH solution transferred to 400 ml beaker using avolumetric pipette. A magnetic stir bar was placed in the beaker withCH3COOH solution and the beaker placed on the magnetic stirrer belowthe ring stand. The PH sensor was inserted. The positions of the dropcounter and burette were adjusted to align with the center of themagnetic stirrer. The apparatus were further arranged according tothe instructions in the Lab manual. Titrations were done to producethree reliable trials.
25.00 ml of aceticacid was titrated with 1.0 M NaOH. The acetic acid was back-titrateduntil the solution was just slightly acidic and sodium hydroxideadded to reach the endpoint with the help of phenolphthalein and PHmeter. After the endpoint had been achieved, more 25.00 ml of 1.0MCH3COOH was added to form a half-titrated solution.
Mass of KHP (g)
Moles of KHP that reacted
Volume of NaOH (ml)
Molarity of sodium hydroxide (M)
Average Molarity of sodium hydroxide
Table 2: Titrationof acetic acid, and determination of Ka
PH of acetic acid before titration
Volume of NaOH (ml)
Average volume of NaOH (ml)
Moles of NaOH that reacted
Molarity of acetic acid
Ka of acetic acid
Table 3: Titrationcurves
Figure2: PH against volume trial 1
Figure 1: PH againstvolume trial 2
At the equivalencepoint, the moles of the titrant, sodium hydroxide, are equal to themoles of acetic i.e. stoichiometric equivalence point. Theconcentration of hydroxyl ions at the equivalence point is higherthan that under normal conditions. The PH increased with the additionof more sodium hydroxide, but the rate of increase was rapid towardsthe equivalence point. The increase in PH was due to the increase ofhydroxyl ions contributed by the dissociating sodium hydroxide. Atthe equivalence point, all the acid reacted with sodium hydroxide toform the basic salt, NaCH3COO. The values of theionization constants obtained through the normal titration andtitration curves were in agreement.
Hamid, Q., Shannon, J., & Martin, J. (2005). Physiologicbasis of respiratory disease. Hamilton: B. C. Decker.
Pilling, G. (1999). Salters higher chemistry. Oxford:Heinemann.
Stanton, B., Zhu, L., & Atwood, C. H. (2010). Experimentsin general chemistry featuring measurement: Guided inquiry,self-directed, and capstone. Belmont, CA: Brooks/Cole, CengageLearning.