# Strong and Weak Acids Lab

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Strong and Weak Acids Lab

Purpose

The purpose of this lab is to compare properties of strong and weak acids and bases as well as properties of dilutions of strong acids.

Theory

pH is a measure of the acidity of a water solution. The acidity of a water solution is determined by the relative number of hydrogen ions (H+) or hydroxyl ions (OH-) present. Acidic solutions have a higher relative number of hydrogen ions, while basic solutions have a higher relative number of hydroxyl ions. In water solutions, the product of the molar concentrations1 of hydrogen and hydroxyl ions is equal to a dissociation constant (Kw). Knowing the value of the constant and the concentration of hydrogen ions makes it possible to calculate the concentration of hydroxyl ions, and vice versa. At 25°C, the value of Kw is 1 x 10-14. pH is defined to be the negative logarithm of the hydrogen ion concentration pH = - log[H+] At 25°C, a neutral solution has a pH of 7.0, while solutions with pH < 7 are acidic and solutions with pH > 7 Some indicators can be used to determine pH because of their colour changes somewhere along the change in pH range.

Conductivity; Acids and bases in aqueous solutions will conduct electricity because they contain dissolved ions. Therefore, acids and bases are electrolytes. Strong acids will be strong electrolytes because they are fully ionized while weak acids are partially ionized.

Acids will react quickly with metal depending on how strong the concentration of the acid is. The weaker the acid the slower the reaction will occur.

Procedure

For part one of the lab, six small beakers were laid out, each contained one of the six acid/bases solution. Six strips of Ph paper were taken and one was dipped in each beaker, and the colour the Ph paper changed, it was compared to the Ph chart on the container. Afterwards each beaker has a Ph meter dipped in the solution and a Ph reading was taken.

For part two of this lab, 5 beakers each contained one of the acid solution detailed in the theory were placed under a lightbulb and the brightness of the light bulb was observed.

For part three of this lab, a piece of Magnesium, zinc and copper were placed in three separate wells in a spot plate. This was a repeated process, and then 3M of HCl was poured over the first three metal pieces, while 3M of acetic solution was poured over the other three metals. The reaction rate of the metals was observed.

For part four of this lab, three test tubes were filled with 2ml of different concentrations (0.01M, 0.1M,1.0M) of HCl and then a strip of Magnesium was added in each one test tube. The reaction rate of the different concentrations of the HCl and strips of Mg was observed.

For part of this lab, an Erlenmeyer flask was filled with 10ml of 02M HCl, while a solution of 0.1M of NaOH was titrated into it until the acid was completely neutralized. This process was repeated with the same concentration and volume of acetic acid.

Data/Results

Measure of pH

Observations: When the pH paper was dipped into the solutions it would change color according to the level of pH. The pH meter was reading a pH level even if there was nothing to measure. All of the solutions were clear.

Sample calculation for the theoretical pH

value of 0.01M and 0.1M of Ethanoic acid

 CH3COOH H+ CH3COO- Initial 0.01M 0 0 Change -x x X Equilibrium 0.01M x X Acid/Base 0.01M HCl 0.1M HCl 0.01M Acetic Acid 0.1M Acetic Acid 0.01M NaOH 0.01M NaOH pH(paper)±0.5 2 1 4 3 8 12 pH(meter)±0.2 1.5 0.6 2.5 2.5 9.1 13.2 Theoretical 2 1 3.37 2.9 12 13

Ka=1.8x10-5==      -7[pic 1][pic 2][pic 3]

x=4.24x10-4=[H+]

pH =+] = -log[4.24x10-4] = 3.37[pic 4]

pH =3.37

 CH3COOH H+ CH3COO- Initial 0.1M 0 0 Change -x x x Equilibrium 0.1M x x

 Acid/Base 0.01M HCl 0.1M HCl 0.01M Acetic Acid 0.1M Acetic Acid 0.01M NaOH 0.01M NaOH Paper % uncertainty 25% 50% 15% 17% 4% 4% Meter % uncertainty 10% 20% 6% 7% 2% 2% % error of the paper 0% 0% 19% 3% 33% 8% % error of the meter 25% 40% 26% 14% 24% 2%

Ka=1.8x10-5==      -6[pic 5][pic 6][pic 7]

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