Week5

Week 5 Overview

• Sections Covered: 3.1 Chemistry in Context, 3.2 Determining Empirical and Molecular Formula, 3.3 Molarity, 3.4 Other Units for Solution Concentration

Section 3.1: Learning Objectives

• Calculate formula masses for covalent and ionic compounds.

• Define the mole and relate it to Avogadro’s number.

• Explain mass-mole relationships and perform related calculations.

Formula Mass

• Definition: The formula mass is the sum of the average atomic masses of all atoms in a substance's formula.

• Covalent Substances: Exist as distinct molecules with a formula mass often referred to as molecular mass.

• Example: Chloroform (CHCl3) average mass = 119.37 amu.

Additional Examples

• Aspirin (C9H8O4): Average mass = 180.15 amu.

Formula Mass for Ionic Compounds

• Nature of Ionic Compounds: Composed of cations and anions in ratios leading to neutral bulk matter.

• Formula Mass: For ionic compounds, it is not typically referred to as molecular mass.

• Example: Table salt (NaCl) consists of a 1:1 ratio of sodium and chloride ions; formula mass = 58.44 amu.

The Mole

• Definition: A mole is the amount of a substance that contains the same number of entities as atoms in 12 grams of carbon-12.

• Purpose: Provides a link between the mass of a sample and the number of particles.

Avogadro’s Number

• Value: 6.02214179 × 10^23 entities.

• Context: Indicates different masses for 1 mole of different elements.

Molar Mass

• Definition: The mass in grams of one mole of a substance, expressed in g/mol.

• Concept: Molar mass = atomic or formula mass in amu.

• Example: 1 mole of carbon-12 has a mass of 12 g.

One Mole of Different Compounds

• Examples include:

  • 1-octanol (C8H17OH): 130.2 g

  • Mercury(II) iodide (HgI2): 454.9 g

  • Methanol (CH3OH): 32.0 g

  • Sulfur (S8): 256.5 g

Calculations

• Grams to Moles: Convert grams of potassium to moles using its molar mass.

• Example: 4.7 g of potassium is approximately 0.12 mol using 39.10 g/mol.

Moles to Grams

• Example: 9.2 × 10^–4 mol of argon has a mass of 0.037 g using 39.95 g/mol.

Mass to Atoms Calculation

• Example: 24.0 g copper wire contains approximately 3.24 × 10^17 atoms using its molar mass.

Molar Mass Calculation Example

• Glace and Calculation: Calculate moles in a 28.35 g sample of glycine (C2H5O2N) which equals 0.378 mol after calculating its molar mass of 75.07 g/mol.

Carbon Atoms in Saccharin Calculation

• Given 40.0 mg of saccharin (C7H5NO3S) with a molar mass of 183.18 g/mol, calculate the number of carbon atoms in this quantity.

• Result: 1.31 × 10^23 molecules leads to approximately 9.17 × 10^23 carbon atoms.

Section 3.2: Learning Objectives

• Percent Composition: Compute the percent composition of a compound.

• Empirical and Molecular Formulas: Determine these using mass data.

Percent Composition from Mass

• Definition: The percentage by mass of each element in a compound.

• Example Computation: A 10.0 g sample with 2.5 g H and 7.5 g C leads to 25% H and 75% C.

Empirical Formulas

• Method: Derived from mass data by converting masses to moles, finding the smallest mole ratio, and adjusting if necessary.

Handling Fractional Subscripts

• Conversion Method: If ratios yield fractions, multiply all by the same number to achieve whole numbers.

Empirical Formulas from Percent Composition

• Procedure: Assume a 100 g sample, convert to moles, find ratios.

Example Calculation

• Ethanol's gas with 27.29% C and 72.71% O gives CO2 as the empirical formula.

Derivation of Molecular Formulas

• Process: Obtained by multiplying the empirical formula by n based on the molecular mass compared to empirical mass.

• Example: Empirical formula CH2O with a molecular mass of 180 amu yields C6H12O6.

Section 3.3: Learning Objectives

• Solutions: Fundamental properties, calculating concentrations using molarity, and dilution calculations.

Solutions Overview

• Definition: A homogeneous mixture with uniform composition.

• Components: Solvent (larger concentration) and solute (lower concentration).

Molarity Definition

• Formula: Molarity (M) = moles of solute per liter of solution.

Calculating Molarity Example

• Found in soft drinks (e.g., 0.133 mol of sucrose in 355 mL leads to concentration calculations).

Dilution Basics

• Definition: Making a solution less concentrated by adding solvent.

• Dilution Equation: n1 = n2, allowing for concentration-volume calculations pre and post dilution.

Section 3.4: Learning Objectives

• Concentration Units: Learn various units such as mass percentage and ppm.

• Mass Percentage Calculation: Ratio expressed as a percentage of solute mass to solution mass.

Volume Percentage Explained

• Useful for liquid solutes in liquid solvents with measurements expressed in %.

PPM and PPB

• Very low concentration measurements often reported in ppm and ppb, especially relevant in health contexts.