Chem110 Notes - Physical and Chemical Properties, Reactions, Solutions

Physical and Chemical Properties

  • Physical Properties: Can be measured without altering the chemical formula. Examples include:

    • Melting Point
    • Boiling Point
    • Color
    • Odor
    • Density
    • Physical State (solid, liquid, gas)

  • Physical Changes: Changes in the form of a material without changing its chemical formula. Examples include:

    • Melting
    • Boiling
    • Dissolving solids in liquids
    • Filtration
    • Distillation

  • Chemical Changes: Changes that result in the formation of new chemical compounds. Examples include:

    • Rusting of iron (oxidation)
    • Burning of gasoline (combustion)
    • Photosynthesis

Chemical Equations

  • Definition: A chemical equation uses chemical formulas to represent reactions.
  • Stoichiometric Coefficients: Indicate the number/moles of each particle involved in the reaction.
  • Physical Phase Indication: The physical phase of substances is denoted by letters in parentheses (s, l, g, aq).

Special Conditions in Reactions

  • Special conditions for reactions can be indicated above or below the arrow in the equation.
  • Catalysts (e.g., platinum, ruthenium) may be required to speed up reactions.
  • Some reactions require specific solvents.

Material Balance

  • Conservation of Mass: Atoms cannot be created or destroyed during a chemical reaction; thus,
    • The number and type of atoms in reactants must equal those in products (This balances the equation).

Interpreting Chemical Equations

  • Example:
    • 2H<em>2(g)+O</em>2(g)→2H2O(g)2 H<em>2(g) + O</em>2(g) \rightarrow 2 H_2O(g)
    • Represents the reaction of hydrogen and oxygen to form water.

Balancing Chemical Equations

  1. Identify reactants and products; write unbalanced equation.
  2. Count atoms on both sides.
  3. Adjust coefficients to balance.
  4. Rewrite balanced equation.

Stoichiometry

  • Determines the amount of products from given reactants.
  • Examples include:
    • Moles of H₂O from 0.780 mol of CH₄.
    • Grams of O₂ needed to produce 14.0 g of water.

Solutions

  • Definition: A homogeneous mixture of two or more components.
  • Components:
    • Solutes: Dissolved species.
    • Solvents: The medium dissolving the solutes (e.g., sugar in tea).
  • Liquid solutions are primarily discussed in Chem 110.

Concentration

  • Definition: Proportion of solute to solvent.
  • Types of Concentration:
    • Weight Fraction: Mass of solute to mass of solution.
    • Molality: Moles of solute to mass of solvent.
    • Molarity: Moles of solute to volume of solution (most common).
    • Dilute Solutions: Low solute concentration.
    • Concentrated Solutions: High solute concentration.
    • Saturated Solutions: Maximum solute dissolved.
    • Mole Fraction: Moles of solute to total moles.

Preparing Solutions in the Lab

  • Weigh the solute and transfer to volumetric glassware.
  • Add solvent to dissolve solute and fill to mark.

Calculating Molarity in the Lab

  • Example: Dissolving 14.8 g of CuSO₄ in 100.00 mL of water.
    1. Determine moles of solute.
    2. Convert volume to liters.
    3. Calculate molarity.

Molarity as a Conversion Factor

  • Molarity can be used for dimensional analysis in conversions of grams to volume.

Dilution

  • High concentration stock solutions can be diluted.
  • The number of solute molecules is conserved; only the solution's volume changes.

Dilution Calculations

  • Rearranging molarity to find moles before and after dilution using the dilution equation:
    • M<em>1V</em>1=M<em>2V</em>2M<em>1V</em>1 = M<em>2V</em>2

Mass Fraction

  • Definition: Proportion of each element by mass.

Mass Fraction from Chemical Formulas

  • Example for calculating mass fraction of elements in compounds (e.g., ammonium nitrate) using the chemical formula.

Mass of an Element in a Compound

  • Determine the mass of each component in a sample using the compound's molar mass.

Combustion Analysis

  • Used to find the empirical formula by analyzing combustion reactions.
    • Steps include burning a known mass, measuring CO₂ and H₂O produced, and relating it to C and H content.

Empirical vs. Molecular Formulas

  • Empirical Formula: Simplest formula with smallest whole-number ratios of elements.
  • Molecular Formula: Shows actual numbers of atoms in a compound, multiples of the empirical formula.

Determining Empirical Formulas from Mass Fractions

  • Example: Determine empirical formula from given mass percentages by calculating moles and their ratios.

Determining Molecular Formulas

  • If the empirical formula and molar mass are known, the molecular formula can be determined using the formula mass to find the multiplier of the empirical formula.
    • Example: For ascorbic acid, determine its molecular formula given its empirical formula and molar mass.