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Acid and Base Dissociation:
Concerning the concepts of acids and bases, it is important to keep in mind the relative dissociative strength of the molecules in solution. Basically, what this means is keeping in mind how easily the molecule can break into its component ions. This breaking apart of the ions is called dissociation.
The terms “STRONG” acid and “STRONG” base refer to the relevant strong driving forces that cause the dissociation of H+ (protons) or OH- (hydroxide ions), respectively. Molecules that form strong acids and bases have a much greater ability to give away H+’s or OH-‘s compared to “WEAK” acids and “WEAK” bases. The difference between a dissociative strong or weak molecule is clear for very strong cases, but when moderate driving forces are present, the differentiation between strong and weak becomes situation conditional. Acids strengths are characterized quantitatively by Ka (acid/base dissociation constants) and pKa (-log Ka), but in most common conversation, referring to a molecule as a strong acid or strong base means that it will 100% fully dissociate into ions when mixed into a solvent such as water. Therefore, it is helpful to group all the halogenic acids (HF, HCl, HBr, HI) and sulfuric acid (H2SO4)as strong acids, while most other common acids are considered weak. In the case of bases, alkali metal bases (LiOH, NaOH…) are strong, and most others are weak.
Conjugates:
Acids and bases have a relationship to each other through conjugate acids and conjugate bases. When an acid has lost its H+, it is now the conjugate base and, likewise, when the base has gained an H+-, it becomes the conjugate acid.
For example:
Acid: HCl , Conjugate Base: Cl-
Acid: CH3COOH , Conjugate Base: CH3COO-
Base: OH- , Conjugate Acid: H2O
Base: NH3 , Conjugate Acid: NH4+
This Week's Experiment:
In this experiment, the strong acid, HCl, and the strong base, NaOH, are used to create serial dilutions which are combined to form salt solutions of NaCl according to the reaction below:
HCl(aq) + NaOH(aq) à NaCl(aq) + H2O(l)
Because HCl has strong dissociative forces, it readily gives up its 1 equivalent of H+ cation, which floats away in solution from the Cl- anion. In the case of NaOH, Na+ readily gives up its 1 equivalent of OH- anion in solution and the two separate according to the reactions below:
HCl(aq)à H+(aq) + Cl-(aq)
NaOH(aq) à Na+(aq) + OH-(aq)
In every case, with every molecule, full ion dissociation occurs. This means that no HCl or NaOH molecules exist, instead everything is broken up into equivalents of H+, Na+, Cl- and OH- in solution. Each ion is stabilized by water molecules that are attracted by intermolecular forces. If weak acids or bases had been used in this experiment, it would be impossible to assume that all of the molecules dissociated, instead only a percent of the molecules would break up. Exactly to what extent the molecules dissociate is characterized by its Ka and pKa.
Assuming full dissociation, the theoretical number of moles of H+(aq) can be determined and thus also a theoretical pH for each solution. Calculate the number moles of H+ using molarity calculations and stoichiometric ratios, then use the pH and pOH equations to find theoretical pH. The common equations used for pH calculations are:
pH = -log [H+] and [H+] = 10^-pH
pOH = -log [OH-] and [OH-] = 10^-pOH
pH + pOH = 14 and [H+]*[OH-] = 10^-14
For an acid example, 10ml of 0.5M aqueous HCl solution would yield 0.005mol of H+(aq) according to:
- L HCl * 0.5mol HCl/L = 0.005 mol HCl
0.005 mol HCl * 1 mol H+/1 mol HCl = 0.005 mol H+
Because pH = -log [H+],
pH (0.5M HCl) = -log 0.005 = 2.3 (pH<7, acidic solution)
A base solution example using 15mL of 0.6M NaOH:
0.015 L NaOH * 0.6mol NaOH/L = 0.009 mol NaOH
0.009 mol NaOH * 1 mol OH-/1 mol NaOH = 0.009 mol OH-
Because pOH = -log [OH-],
pOH = -log 0.009 = 2.0
and
pH = 14 – pOH so pH = 14 – 2.0 = 12.0 (pH>7, basic solution)
All of the above calculations are only possible because we are dealing with a strong acid or strong base!
Serial Dilutions
Firstly, to define terms, the solute is dissolved by the solvent to create the solution.
When conducting “wet” chemistry, serial dilutions are an important tool in experimentation. They allow the researcher to create controlled, systematic dilutions from a stock solution. Each successive solution is diluted so that the concentration of solute is changed by a desired ratio. For example, in 1/10 serial dilutions starting with a stock solution of 5M solute concentration, the serial dilution would be comprised of 5M, 0.5M, 0.05M, 0.005M… Each solution is reduced to one tenth of the previous solutions concentration. For another example, 1/25 serial dilutions starting from a stock solution of 12M solute concentration would be comprised of 12M, .48M, 0.0192M, 0.000768M… The greater the ratio of solute concentration change, the greater the difference in concentration for each successive serial dilution. Among other techniques, serial dilutions are a tool in the chemist’s box of methods, analysis and experimentation.
Theoretical Yield
Stoichiometry is a term that refers to calculations involving chemical equations in such a way as to relate parts of the reaction through ratios based on the coefficients of the balanced equation. A simple example is as follows:
Na+ + Cl- à NaCl
In this example, one mole of Na+ cations could combine with one mole of Cl- anions to form one mol of NaCl. This reaction is said to be “one to one” because all of the parts of the reaction have the same coefficient of 1. The ratio holds true in the case for any amount of cations or anions. Consider if there are 0.0876554 moles of Na+ cations combining with 50 moles of Cl- anion. There are far fewer Na+ cations, so all of them will find Cl- anions. Since one Na+ and one Cl- will make one NaCl, that means that 0.0876554 moles of Na+ cations will form 0.0876554 moles of NaCl and leftover Cl-.
By knowing how many moles of a reactant we are starting with, and using stiochiometry, we can determine the theoretical number of moles of a product that can be formed. This calculated value is called the theoretical yield and it represents the maximum amount of product that can be made. Usually less is made because in real situations, moles can be lost in transferring liquids, mixing and various other handling measures. The actual amount of moles made is called the experimental, or actual, yield. Also, errors resulting in loss of material or hindrance of the reaction can lead to an experimental yield which is less than the theoretical yield.
It’s always good practice to consider the reasons for a difference between the theoretical yield and the experimental yield. These explanations of the errors are information to be included in conclusion sections of reports.
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