Chemistry: Formulas and Composition Study Notes

Introduction

  • Engagement between students and instructor.
  • Overview of the day's lesson and the chapter to be covered.

Chapter Overview

  • This is the last chapter before Exam 2.
  • The chapter covers essential concepts of formulas and compositions in chemistry.
    • Objectives include understanding:
    • Molecular and empirical formulas
    • Percent composition
    • Calculation methods for these concepts.

Chemical Formulas

  • Two main types of chemical formulas:
    1. Molecular Formula
    • Provides the exact number of atoms of each element in a compound.
    • Example of glucose: C₆H₁₂O₆ (exact count of each atom).
    1. Empirical Formula
    • Represents the smallest whole number mole ratio of the elements in a compound.
    • Example: Glucose (C₆H₁₂O₆) has an empirical formula of CH₂O.
    • Explanation: Divide each subscript in the molecular formula by the greatest common factor (6 in this case).
  • Notably, ionic compounds are always expressed as empirical formulas because they do not exist as discrete molecules but rather as ionic arrangements.

Example Comparisons

  • Glucose:
    • Molecular Formula: C₆H₁₂O₆
    • Empirical Formula: CH₂O
  • Benzene:
    • Molecular Formula: C₆H₆
    • Empirical Formula: CH (ratio = 1:1 for C and H).

Ionic Compounds

  • Discussed how ionic compounds, like sodium chloride (NaCl), should be represented by their empirical formula (NaCl represents a 1:1 ratio of sodium to chloride ions due to their ionic nature).
  • Crystal structures discussed:
    • Arrangement in a 3D space where positive and negative ions surround each other.

Moles and Chemical Composition

Moles Concept

  • Used to define quantities of substances in terms of atoms or molecules.
  • Example calculations provided:
    • Oxygen moles in Aluminum Oxide (Al₂O₃): 3 moles of O in 1 mole of Al₂O₃.
    • Oxygen in Calcium Phosphate (Ca₃(PO₄)₂): 8 moles O in 1 mole of the compound.

Percent Composition

  • Definition: Mass percentage of each element in a compound.
  • Formula to calculate percent composition of element A:
    \text{Percent Composition} = \frac{n \times \text{(Molar Mass of Element A)}}{\text{(Molar Mass of Compound)}} \times 100
  • Example calculation with glucose:
    • Molar Mass of Glucose: 180.156 g/mol
    • Calculation for Carbon (C): 40% carbon in glucose, Hydrogen (H): approximately 6.71%, Oxygen (O): approximately 53.28%.
  • The total percentages should equal 100%.

Finding Empirical Formula from Percent Composition

  • Procedures outlined clearly:
    1. Assume a 100 g sample to convert percentage to grams.
    2. Calculate the moles of each element based on their percent composition.
    3. Find the smallest number of moles and simplify the mole ratio to obtain the empirical formula.
    • Example: Given Carbon at 40.92%, Hydrogen at 4.58%, and Oxygen at 54.5%:
    • C: 40.92g, H: 4.58g, O: 54.5g.
    • Convert percentages to masses in a 100 g sample, then to moles, find the overall mole ratio, and simplify.

Calculating Empirical Formula of Compounds

  • Empirical formula calculations demonstrated through sample problems with given percent compositions:
    • For example, ascorbic acid with components calculated as:
    • C: 40.92 g; H: 4.58 g; O: 54.5 g.
    • Assumed to have a total mass of 100 g for ease of calculation.
  • Key steps emphasized:
    1. Assume 100 g for calculations.
    2. Convert to grams, calculate moles.
    3. Find ratios and simplify to derive the empirical formula.

Questions and Engagement

  • Several student questions regarding methods and rationale in deriving empirical formulas.
  • Discussion on when to transition from decimal ratios to whole number representations in empirical formulas, emphasizing significant figures in calculations.