DAT- Gen chem

4.1 Moles and Molar Mass

  • Moles in Chemistry

    • Moles: An SI unit used to measure the amount of any substance.

    • 1 mole = Avogadro's number = 6.022imes10236.022 imes 10^{23} particles (atoms, molecules, etc.).

    • This value represents the number of atoms in 12 grams of Carbon-12.

    • Why 12 grams?

    • The theoretical atomic mass of the carbon-12 isotope (6 protons and 6 neutrons) means that the atomic mass (12 grams) of carbon is equal to exactly 1 mole of carbon.

  • Molar Mass: The mass of one mole of a substance expressed in grams per mole.

    • Molar Mass Formula: Molar mass = Atomic mass (in g/mol).

  • Calculating Molar Mass of Water (H₂O):

    • H: 1 g/mol, O: 16 g/mol.

    • Molar mass of H₂O = 2(1) + 1(16) = 18 g/mol.

  • How to Calculate Moles (from mass to moles):

    1. Plug in the given mass (in grams) of the substance.

    2. Calculate the molar mass (M).

    3. Divide the mass by the molar mass to find the number of moles (n):
      n=mMn = \frac{m}{M}.

Example: Calculating Moles

  • Example 1: How many moles are in 348g of caproic acid, C₆H₁₂O₂?

    • Given: 348 g of C₆H₁₂O₂.

    • Molar mass calculation:

      • C: 6(12 g/mol) + H: 12(1 g/mol) + O: 2(16 g/mol) = 116 g/mol.

    • Calculation:

    • n=348g116g/mol=3moln = \frac{348 g}{116 g/mol} = 3 mol of C₆H₁₂O₂.

  • Example 2: How many moles of H₂O are produced when 5 moles of oxygen are reacted?

    • Reaction: 2H<em>2+O</em>2→2H2O2H<em>2 + O</em>2 → 2H_2O

    • Using mole ratio to find moles of H₂O produced from 5 moles of O₂:

    • 5 mol O<em>2×2 mol H</em>2O1 mol O<em>2=10 mol H</em>2O5 \text{ mol } O<em>2 \times \frac{2 \text{ mol } H</em>2O}{1 \text{ mol } O<em>2} = 10 \text{ mol } H</em>2O.

  • Example 3: Given 42.5 grams of NaNO₃ equals 0.5 moles, find the mass when 1 mole of NaNO₃.

    • Calculation:

    • 0.5 mol NaNO3=42.5g0.5 \text{ mol } NaNO₃ = 42.5g

    • To find 1 mole:

      • 1 mol NaNO₃×2=42.5g×2=85g1 \text{ mol NaNO₃} \times 2 = 42.5g \times 2 = 85g.

4.2 Percent Composition

  • Percent Composition: The percentage of the total mass of a compound attributed to a specific element.

    • This can be used to determine molecular and empirical formulas.

  • How to Calculate Percent Composition:

    1. Find the molar mass of the entire compound using each element's molar mass.

    2. Divide the specific component's molar mass by the entire compound's molar mass.

    3. Multiply the decimal by 100 to find the percent composition.

4.3 Empirical and Molecular Formula

  • Molecular Formula:

    • Regular formula showing the number of atoms of each element in a compound.

    • Reflects the actual elemental composition of a compound.

    • Found using empirical formula and molecular weight.

    • Stable, exists as a whole number ratio of the empirical formula.

    • Example: C₂H₆.

  • Empirical Formula:

    • Reduced formula; the simplest or most reduced ratio of atoms in a compound.

    • Represents the simplest elemental composition.

    • Found from mass percentages of elements in a compound.

    • Can be determined from combustion or composition analysis.

    • Example: CH₃.

  • Determining Empirical Formula from Molecular Formula:

    1. Divide the molecular formula by the subscripts' largest common factor.

  • Determining Molecular Formula from Empirical Formula:

    1. Calculate the molar mass of the empirical formula.

    2. Divide the given molecular molar mass by the molar mass of the empirical formula.

    3. Multiply each subscript by the whole number.

4.4 Balancing Equations

  • Chemical equations can be balanced using coefficients to reflect the number of atoms per side.

  • Steps to Balance a Chemical Reaction:

    1. Identify each element and its quantity for both reactants and products.

    2. Choose an element and compare quantities on each side.

    3. Add coefficients as needed.

    4. Repeat as necessary until all elements are balanced.

    5. Check overall balance by counting quantities.

4.5 Calculations From Balanced Equations

  • Moles to moles, moles to grams calculations use molar ratios for conversions.

4.6 Limiting Reagent

  • Limiting Reagent: The reactant that determines how much product is formed; it is completely consumed in a reaction.

  • Excess Reagent: The other reactant not completely consumed.

  • How to Determine Limiting Reagent:

    1. Balance the equation.

    2. Convert all given information to moles.

    3. Pick one reactant, calculate required amount of the other reactant to confirm limiting reagent.

    4. If excess exists, the chosen is limiting; if not, the other is limiting.

4.7 Density

  • Density: A measure of mass per unit volume of a substance.

    • Density of water is approximately 1 g/mL.

    • An intrinsic property constant at given temperature/pressure for pure substances.

4.8 Theoretical and Percent Yield

  • Types of Yields:

    • Actual Yield: The real quantity of product from a reaction.

    • Theoretical Yield: The maximum expected product from an ideal reaction.

    • Percent Yield: rac{\text{Actual Yield}}{\text{Theoretical Yield}} \times 100 ext{%} .

  • Important Concepts:

    • Actual yields are typically lower than theoretical yields.

    • Theoretical yield is calculated using the limiting reagent.

5.1 Solution Properties

  • Solution: Liquid mixture with solute uniformly distributed.

    • Solute: Minor component dissolved in solvent.

    • Solvent: Major component, usually liquid.

  • Solution Types:

    • Unsaturated: Less than the maximum solute concentration.

    • Saturated: Maximum solute concentration.

Solubility: The ability of solute to dissolve in solvent (e.g., salt in water).

5.2 Polarity and Classification of Matter

  • Polarity: Separation of electric charge; polar molecules possess dipole moments.

    • Non-polar molecules are uniformly charged.

  • Intermolecular Forces:

    1. London Dispersion Forces: Weak, occur in all molecules.

    2. Dipole-Dipole Interactions: Stronger, occur in polar molecules.

    3. Hydrogen Bonding: Very strong between H and F, O, N.

5.3 Concentration Calculations

  • Molarity: Moles of solute per liter of solution.

    • Molality: Moles of solute per kg of solvent.

  • Types of Solutions:

    • Diluted Solution: Low solute concentration.

    • Concentrated Solution: High solute concentration.

    • Saturated Solution: Maximum solute concentration.

5.4 Colligative and Non-Colligative Properties

  • Colligative Properties: Depend on solute concentration, not identity.

    • Boiling Point Elevation: ΔT<em>b=K</em>b×m\Delta T<em>b = K</em>b \times m(molality)

    • Freezing Point Depression: ΔT<em>f=K</em>f×m\Delta T<em>f = K</em>f \times m.

5.5 Net Ionic Equations

  • Molecular Equations: All substances as molecules.

  • Ionic Equations: Ionic compounds as separate ions.

  • Net Ionic Equations: Show only compounds/ions directly involved in reactions.

12.1 Oxidation-Reduction

  • Oxidation: The process of losing electrons.

  • Reduction: The gain of electrons.

  • Oxidizing Agent: Gains electrons, undergoes reduction.

  • Reducing Agent: Loses electrons, undergoes oxidation.

12.2 Redox Reactions (Acidic)

  • Half Reactions: Parts of redox reactions showing reduction or oxidation.

  • Steps to Balance Redox in Acidic Medium:

    1. Write half reactions.

    2. Balance other atoms.

    3. Balance oxygen by adding H2O.

    4. Balance hydrogen by adding H+.

    5. Equalize charges by adding electrons.

    6. Combine the balanced half reactions.

12.4 Electrochemical Cells

  • Electrochemical Cells: Convert chemical energy to electrical energy via redox reactions.

    • Galvanic Cells: Spontaneous, anode has excess electrons.

    • Electrolytic Cells: Non-spontaneous, require energy.

    • Standard Cell Potential: Sum of oxidation and reduction potentials (E cell).

12.5 Standard Cell Potential

  • E cell: Calculated from the reduction potentials.

    1. Determine half reactions.

    2. Look up table values.

    3. Balance half reactions.

    4. Add potentials for overall cell potential.

12.6 Electrolysis Calculations

  • How to calculate moles, energy change, and cell potential related to electrolysis.

11.1 Dynamic Equilibrium and Equilibrium Constants

  • Chemical reactions can reach equilibrium when forward and reverse rates equalize.

  • Equilibrium Constant (K_eq): Ratio of products to reactants at equilibrium; excludes pure solids/liquids.

11.2 Le Chatelier's Principle

  • Changes (concentration, pressure, temperature) will shift equilibrium to restore balance.

11.3 Precipitation Reactions

  • Solubility Product Constant (K_sp): Indicates how much solute can dissolve in water.

    • Qsp: Current solubility state of the solution; relationships to Ksp.

11.4 Acid/Base Equilibria

  • Amphoteric Species: Can act as both acid and base.

  • Autoionization of Water: Water molecules transfer protons.

  • Effects of strong acids/bases on equilibrium.

12.7 Titrations

  • Titration: Method to determine concentration of an acid/base.

    • Use indicators for noting endpoints.

    • Half Equivalence Point: Midpoint in titration; pH equals pK_a.

    • Ecological significance in acid-base reactions.