Chapter 3 Mass Relationship in Chemical Reaction

Atomic Mass: The average atomic mass is calculated using the weighted average of isotopes.
- Example: Carbon's atomic mass is 12.01 amu, derived from isotopes 12C and 13C.
Avogadro's Number: 1 mole = 6.022 x 10²³ particles.
- Important for converting between particles and moles.
Molar Mass: Reflects the mass of one mole of a substance (in grams equivocal to atomic mass in amu).
- Example: 1 mole of 12C equals 12 g.
Molecular Mass vs. Formula Mass: Molecular mass is the sum of atomic masses in a molecule (in amu), while formula mass pertains to ionic compounds.
Percent Composition: Determines the mass percent of each element in a compound.
Empirical vs. Molecular Formula: Empirical formula reflects the simplest ratio of atoms while the molecular formula indicates the actual number of atoms in a molecule.

Definition: Atomic mass in amu quantifies mass of individual atoms.
- E.g., average mass derived from natural isotopes of an element.
Example Calculation: For Copper (Cu), average mass = (0.6909 x 62.93 amu + 0.3091 x 64.9278 amu) = 63.55 amu.

Definition: A mole is a counting unit in chemistry representing 6.022 x 10²³ particles of a substance.
Conversions: Grams to moles and vice versa depend on molar mass.
- E.g., 42 g of O = 42 g O x (1 mol O / 32 g O) = 1.31 mol O.

Conversions Between Grams and Moles: Use molar masses to convert between grams and moles effectively.
When converting moles to grams or grams to moles, factor in the compound's molar mass—this is a crucial step in stoichiometry.

Calculation Steps: Using the formula:
% of element = (mass of element in 1 mole of compound / molar mass of compound) × 100%
Example: For H₃PO₄, molar mass = 97.99 g, mass percentages are calculated to yield %H = 3.086%, %P = 31.61%, and %O = 65.31%.

Empirical Formula: Simplest whole-number ratio from the actual atomic ratios.
Determining the Empirical Formula: Convert grams to moles, then divide each by the smallest number of moles to get ratios.
Check reduction to simplest form.

Reactants and Products: Balance equations ensuring conservation of mass.
- E.g., Mg + O₂ → MgO as a basic representation of product formation.
Balancing Equations: Requires adjusting coefficients (not subscripts) to ensure equal numbers of each atom on both sides.

Mole Ratios: Utilize coefficients in balanced equations for calculating amounts of reactants/products based on available quantities.
Example: Given 3.4 moles of Cl₂, based on the reaction 2 Na + Cl₂ → 2 NaCl, calculate moles of NaCl produced.

Mass Measurements: Convert grams of substances to moles using molar mass, crucial for quantitative chemical analysis.
Example Calculation: Given 19.7g of MgCl₂, use stoichiometry to find mass of AgCl produced in the reaction with AgNO₃.

When calculating moles or grams, always 1) identify the known values, 2) utilize molar mass for conversion, 3) make sure equations are balanced for accurate stoichiometric calculations.
Always verify results—check sums, conversions, and ratios to ensure mathematically consistent answers.