Question

I have to write out a lab plan for an upcoming experiment! and complete the table below

Experiment 9: Introduction to Acids, Bases, and pH Learning

Objectives

» to gain experience with a wide variety of acids and bases

» to learn how to predict the principal species in a solution

»to learn how to measure acidity, basicity, and pH

» to predict what will happen when different acid and base solutions are mixed

Be sure to record all data in your own notebook. Materials Lots of acids and bases in various concentrations, litmus paper, well plates and pH meters.

Introduction to Acids and Bases Acids and bases are all around us. They are present in the fruits and vegetables we eat, in the sodas we drink, and in the various commercially available products that we use to unclog our kitchen sinks! Acids are characterized with respect to their ability to produce hydronium ions (H3O + ). Bases are characterized with respect to their ability to produce hydroxide ions (OH– ). Both acids and bases can be classified as “strong” or “weak,” which are allusions to whether the species is a strong electrolyte (placing many ions in solutions, see (1) below) or a weak electrolyte (placing a small number of ions in solution, see (2), below). (1) HCl(aq) + H2O()  H3O + (aq) + Cl (aq) Production of hydronium ion by a strong acid (2) NH3 (g) + H2O() NH4 + (aq) + OH (aq) Production of hydroxide ion by a weak base Strong acids produce large amounts of hydronium ions because their reaction with water goes essentially to completion, thus producing a stoichiometric number of ions. Weak acids react with water in the same type of proton-transfer process, but the equilibrium constant is not as large and only a minimal number of ions are created. The same can be said for strong and weak bases with respect to the amount of hydroxide formed. The chemical structure of the acid or base plays a large role in its ability to lose or attract a proton. 9-2 Introduction to Acids, Bases and pH In order for you to predict the principal species in solution, you must know if a particular acid or base is strong or weak. It is best to just memorize the strong ones: Strong Acids HCl hydrochloric acid HBr hydrobromic acid HI hydroiodic acid HNO3 nitric acid HClO4 perchloric acid H2SO4 sulfuric acid Strong Bases NaOH sodium hydroxide KOH potassium hydroxide Strong acids and strong bases are strong electrolytes; their principal species in solution are the fully dissociated ions – H3O + , Br– NO3 – , ClO4 – , HSO4 – , Na+ , K+ , OH– , etc. IT IS CRITICAL THAT YOU MEMORIZE THE NAMES AND FORMULAS OF THESE STRONG ACIDS AND BASES AND ALSO KNOW WHAT IONS THEY PRODUCE. DO IT NOW. Pretty much anything else that has a name that ends with “acid” – hydrofluoric acid, acetic acid, formic acid, nitrous (rather than nitric) acid, phosphoric acid, etc. – is a weak acid. Just about anything else containing nitrogen in it is a weak base. Weak acids and weak bases are weak electrolytes. Their principal species in solution is just themselves. The lower the pH of a solution, the more acidic it is. Conversely, the higher the pH of a solution, the more basic it is. A pH of 7 indicates that the concentrations of H3O + and OH– are equal; such solutions are called “neutral.” Acids and bases react with each other to form neutral solutions with salts and water as the products. (Can you write the reaction equation between HCl and NaOH?) Salts in solution will also form acidic, basic or neutral solutions depending on the properties of the ions. This will be apparent in some of the work carried out in today’s experiment.

Procedures Part 1: Acid, Base, or Neutral?

Types of substances in this table include neutral salts, strong acids, strong bases, weak acids (might be a salt), and weak bases (might be a salt). Water is present in large quantities in any aqueous system, so you need not list H2Omolecules as a principal species.

Solution Type of Substance Principal Specie(s) in Solution Prediction Measured pH

a 0.1 M HCl Strong acid H3O + , Cl– acidic

b 0.1 M H3PO4 Weak acid

c 0.1 M CH3CO2H Weak acid

d 0.1 M NH4Cl Salt (acidic) NH4 + , Cl–

f 0.1 M NaCl Salt (neutral) Na+, Cl– neutral

g 0.1 M NaO2CCH3 Salt (basic)

h 0.1 M NH3 Weak base NH3

i 0.1 M NaOH Strong base

When you get to lab, discuss your predictions with your partner and try to come to a consensus on your best guesses. Then measure the pH of each solution using a pH meter (Read the pH Meter Blue Pages! The use of these meters will also be demonstrated for you.) Record your predictions and observations in tabular form in your notebook. Note any surprises, as you will need to explain them later!

Answer 1

Solution                 Type of Substance Principal Specie(s) in Solution          Prediction Measured pH

a 0.1 M HCl    Strong acid H3O + , Cl– acidic                pH = -log[H+] = -log(0.1) = 1

b 0.1 M H3PO4 Weak acid         H3PO4 H+ +H₂PO₄^2-                            1.58

c 0.1 M CH3CO2H               Weak acid             CH3COOH CH3COO- + H+                  2.8

d 0.1 M NH4Cl               Salt (acidic) NH4 + , Cl–                pH < 7 (since it is acidic salt)

f 0.1 M NaCl              Salt (neutral)    Na+, Cl–l                   pH =7(NEUTRAL)

g 0.1 M NaO2CCH3             Salt (basic)                         Na+, CH3COO-              pH > 7(BASIC SOLUTION)

h 0.1 M NH3                       Weak base      NH3+H2O NH4+ + OH-     pH = 11.13 approximately(basic)

i 0.1 M NaOH                      Strong base                         Na+ and OH-                  pH > 7(BASIC SOLUTION)

b) pH of 0.1M H3PO4: To calculate the pH of weak acid, we must use ICE table.

H3PO4 H⁺ + HPO₄^-2

0.1M          0         0

-0.1x         +0.1x   +0.1x

0.1(1-x)      0.1x     0.1x

Ka = [H+]x [HPO₄^-2 ] / [H3PO4]

6.9x10^-3 = 0.1 x² / (1-x)                         as 1>>x , hence 1-x is approximately equal to 1.

By simlifying :

x = 0.26

Hence[H+] = 0.026

pH = -log(0.026) = 1.58

This is for the first dissociation of the phoshoric acid.

c)Calculation of pH of acetic acid:

CH3COOH(aq) + H2O(l) -----> CH3COO-(aq) + H3O+(aq)

Ka IS DISSOCIATION CONSTANT WHICH IS CONCENTRATION OF PRODUCT/CONCENTRATION OF REACTANT.

Ka= [CH3COO-] * [H3O+] / [CH3COOH]

[H2O] is not included because its change in concentration is negligible.
After partial dissociation, [CH3COOH] IS STILL APPROXIMATELY 0.1M.
[CH3COO-] = [H3O+]
Ka of CH3COOH=1.74 * 10^-5

1.74 * 10^-5 = [H30+]^2 / 0.1
[H30]^2 = 1.74*10^-5 * 0.1
= 1.74*10^-6
[H30+]= 1.32*10^-3

pH= -log[H30+]
pH= -log(1.32*10^-3)
pH= 2.88

pH of 0.1M acetic acid is 2.8

h) pH of 0.1M NH3

NH3 +H2O NH4+ + OH-

0.1M               0           0

-0.1x             +0.1x    +0.1x

0.1(1-x)          0.1x      0.1x

Kb of NH3 = 1.8X10^-5

Kb = [NH4+] X [ OH-]/ [NH4OH]

1.8X10^-5 = 0.1 x² / (1-x )                              as 1>>x , hence 1-x = 1

By simplifying:

x = 1.34x10^-2

[OH-] = 0.1 x = 0.1 * 1.34x10^-2

        = 1.34 *10^-3

Hence pOH= -log(1.34 *10^-3)

                 = 2.87

pH = 14- pOH

     = 11.13