Frank Brescia

EXPERIMENT 26 – Synthesis of Aspirin

You will be given approximately 3.0 g of salicylic acid. Set aside a small sample (appx. 20 mg) in a corked and labeled 13x100 mm test tube for later ferric chloride analysis and weigh the rest into a clean, dry 125 mL Erlenmeyer flask. In the fume hood, add 6.0 mL acetic anhydride via pipet. To this white slurry, add 3-4 drops concentrated sulfuric acid. Heat the flask in the hot water bath with periodic swirling for a period of about seven minutes, then allow it to cool to room temperature on the benchtop.

EXPERIMENT 14 – Gravimetric Determination of a Sulfate

The goal of most quantitative chemical analysis measurements is to estimate the relative abundance of an analyte in a chemical sample. For solid and liquid samples, a very common expression of analyte content is based on the mass fraction. This fraction is commonly expressed as a percentage (w/w %), as parts per million (ppm) or as parts per billion (ppb), depending on the concentration level of the analyte. Classical chemical analysis methods are excellent for the determination of analyte concentrations in the range of 1−100 w/w %. In order to estimate the analyte mass fraction of any solid sample, we typically need two measurements, one to estimate the sample mass and one to estimate the analyte mass.

EXPERIMENT 17 – Formula of a Hydrate

Introduction: Many salts form compounds in which a definite number of moles of water are combined with each mole of the anhydrous salt. Such compounds are called hydrates. The water is an integral part of the crystalline structure of the compound. When salt crystallizes from an aqueous solution, the number of water molecules bound to the metal ion are characteristic of the metal and are in definite proportion. In the formula of a hydrate a dot is used to separate the formula of the anhydrous salt from the number of molecules of water. For example, the formula of calcium sulfate dihydrate is written as CaSO4 · 2 H2O not as H4CaSO6. Heat can transform a hydrate into an anhydrous salt. The water can often be seen escaping as steam.

EXPERIMENT 46 – Determination of an Equilibrium Constant

Introduction: In the previous week, we qualitatively investigated how an equilibrium shifts in response to a stress to re-establish equilibrium. This week we will quantitatively assess the equilibrium constant for the same reaction: the reaction of iron(III) cation complexing with a thiocyanate anion (SCN – ) to form the iron(III) thiocyanate complex, Fe(SCN) 2+ (Equation 1). Its equilibrium expression is as shown in Equation 2.

EXPERIMENT 43 – Effect of Temperature on Rate of Reaction

Demonstration: Before class, prepare three identical 250 mL flasks containing approximately 100 mL of water. One should contain water that has been chilled with ice cubes, but with no ice present; one should be room temperature tap water; the last should have hot water in it, but not steaming, or too hot to handle without hand protection. Place a tablet in each flask and note the difference in rates of reaction. The reactions may be timed for more quantitative data. Students may be asked for reasons for differences in rates. (One usually will suggest a temperature difference, although others may suggest that the liquids are not all water, etc.)