31-Dilution final

Molarity

  • Molarity (M) is defined as the number of moles of solute in 1 liter of solution.

    • Formula: M = moles of solute / liters of solution

  • To calculate the number of moles or grams of a solute in a given volume of a solution with a known concentration:

    • ( n = n = C × V = (L × mol/L) × (g/mol)

      where:

    • ( n ) = number of moles

    • ( C ) = concentration (M)

    • ( V ) = volume (L)

Example Calculation

  • To determine the mass of oxalic acid needed to prepare 50.00 mL of a 0.125 M solution:

    • Given the molar mass (MW) of oxalic acid: 90.035 g/mol.

    • Volume of solution = 50.00 mL = 0.050 L

    • Moles of C2H2O4 = ( [0.125 \text{ mol/L}] \times [0.050 L] = 0.00625 \text{ mol} )

    • Mass of C2H2O4 = ( n \times MW = 0.00625 \text{ mol} \times 90.035 \text{ g/mol} = 0.562 \text{ g} )

Dilution

  • Dilution is the process of lowering the concentration of a solution by adding more solvent.

    • It does not change the number of moles of solute in the solution.

    • The relationship between moles and volume before and after dilution can be expressed as:

      • ( n_{initial} = n_{final} )

      • ( V_{initial} \times M_{initial} = V_{final} \times M_{final} )

    • Example: Calculation of dilution concentrations using the formula above shows that the number of particles remains constant pre- and post-dilution.

Process of Diluting a Solution

  • Steps for dilution include:

    1. Use distilled water to fill to desired volume (V_initial).

    2. Multiply the volume by the molarity of standard solution (M_initial).

    3. Divide by total volume of diluted solution (V_final) to find the molarity of the diluted solution (M_final).

Solubility and Dissolution

  • Dissolution of a solute involves the solute integrating into the solvent but can lead to further dissociation into ions in certain cases.

    • Nonelectrolytes: These do not ionize upon dissolving (e.g., glucose, ethanol).

    • Electrolytes: These dissolve and produce ions that allow conductivity in solutions.

      • Strong Electrolytes: Fully dissociate into ions (e.g., NaCl).

      • Weak Electrolytes: Partially dissociate into ions (e.g., HF, PbCl2).

Chemical Reactions

  • Chemical reactions involve the exchange of partners among atoms, and they can be categorized based on the nature of the exchanges:

    1. Molecular partner exchange (dissolution).

    2. Atom exchanges in synthesis, decomposition, substitution, and metathesis.

    3. Proton transfers typical of acid-base reactions.

    4. Electron transfers indicative of oxidation-reduction reactions.

Net Ionic Equations

  • An overall ionic equation specifies reactants and products as ionic substances, represented as dissolved ions.

    • In net ionic equations, spectator ions are omitted, only showing species actively involved in the reaction. For example:

      • ( ext{Na}^+ + ext{Cl}^- + ext{Ag}^+ + ext{NO}_3^- \rightarrow ext{AgCl(s)} + ext{Na}^+ + ext{NO}_3^- )

Precipitation Reactions

  • In precipitation reactions, soluble reactants yield a product with limited solubility.

    • The solid formed is referred to as a precipitate; for example:

      • (Pb(NO₃)₂ + 2NaI → PbI₂ + 2NaNO₃ )

Influence on Solubility

  • Factors influencing solubility include ionic character, charge on ions, and atomic radii:

    1. Higher ionic character increases the solubility influenced by a large ΔH.

    2. Lower charges on ions results in weaker ion-ion interactions, which may allow for greater solubility.

    3. Elements close in the periodic table tend to form more covalent bonds, affecting their behavior in complex ions (e.g., PbCl, PbN2O6, Al2S6O9).