4FO512 CalculationsBB

Chemical Calculations Overview

In forensic chemistry, a thorough understanding of chemical calculations is crucial. This document covers key concepts such as the mole, mass, concentrations, and how to manipulate equations involving these concepts.

Contents

• The Mole

• Rearranging Equations

• Applying the Equations

• Concentrations

The Mole

• Definition: The mole (mol) is the unit for measuring the amount of substance.

• Avogadro’s Number: One mole of any substance contains approximately (6.022 \times 10^{23}) entities (atoms, molecules, etc.), known as Avogadro’s number.

• Molecular Mass: The mass of one mole of a compound in grams is equal to its molecular mass (RMM). For water, the molecular mass is (18 \text{ g mol}^{-1}); thus, 1 mole of water weighs 18g.

Moles, Mass, and RMM

• Relationships: The relationship between mass, moles, and RMM can be summarized in the following equations:

  • To calculate moles: [ \text{Moles} = \frac{\text{Mass (g)}}{\text{RMM (g mol}^{-1})} ]

  • To calculate RMM: [ \text{RMM} = \frac{\text{Mass (g){\text{Moles (mol) ]

  • To calculate Mass: [ \text{Mass (g)} = \text{Moles} \times \text{RMM (g mol}^{-1}) ]

Problem-Solving with Moles

1. Calculating Moles of Compounds: Given the mass of a compound, you can find the number of moles. For example:

   • 51g of ammonia (NH3)

   • 28.4g of sodium sulfate (Na2SO4)

   • 2.96g of magnesium nitrate (Mg(NO3)2)

2. Calculating Mass Based on Moles: To find the mass from the number of moles:

   • 2 mol of magnesium sulfate (MgSO4)

   • 0.2 mol of potassium carbonate (K2CO3)

   • 0.6 mol of ammonium phosphate ((NH4)3PO4)

Chemical Reactions and Molar Ratios

• Example Reaction: [\text{CH}_3\text{COCH}_2\text{COOH} \rightleftharpoons \text{CH}_3\text{COCH}_3 + \text{CO}_2] (3-oxobutanoic acid to propanone and carbon dioxide). This reaction illustrates how diseases, such as diabetes, can result in increased levels of propanone due to fat breakdown.

• Calculating Products: Steps for calculating the mass of propanone produced from 125 mg of 3-oxobutanoic acid:

  1. Calculate the RMM of 3-oxobutanoic acid.

  2. Determine the moles of 3-oxobutanoic acid.

  3. Use molar ratios to find the moles of propanone produced.

  4. Calculate the mass of propanone.

SI Units in Measurements

• Area: [\text{Area} = \text{length} \times \text{length} \text{ (m}^{2}\text{)}]

• Volume: [\text{Volume} = \text{length} \times \text{length} \times \text{length} \text{ (m}^{3}\text{)}]

  • Common conversions: 1 dm³ = 1000 cm³.

Concentrations

• Definition: Concentration denotes the amount of solute in a given volume of solution.

• Types of Concentrations:

  • Mass concentration (g dm⁻³)

  • Molar concentration (mol dm⁻³)

• Formulas:

  • Molar concentration: [\text{Molar Concentration} = \frac{\text{Moles of Solute}}{\text{Volume of Solution (dm}^{3})}]

  • To calculate volume: [\text{Volume} = \frac{\text{Moles of Solute}}{\text{Molar Concentration}}]

  • To calculate moles: [\text{Moles} = \text{Molar Concentration} \times \text{Volume}]

Practice Problems

1. Determine the concentration of sodium chloride (NaCl) in various solutions.

2. Find the number of moles and mass of hydrochloric acid (HCl) in a given volume.

3. Calculate the mass of magnesium hydroxide needed for a 250 cm³ of a 3 mol/dm³ solution.

4. Determine the concentration of magnesium nitrate in a certain volume.

Summary

Through this material, students will gain an understanding of the mole and its importance in chemical calculations. Key skills include the ability to manipulate equations involving moles, mass, and molar ratios, as well as conversion between different units of measurement.