In a chemical reaction, the mass of the products is equal to the mass of the reactants consumed. Are the moles of product equal to the moles of reactant consumed? Explain your answer.

According to the law of conservation of mass, the total mass of the reactants before the reaction is equal to the total mass of the products after the reaction. This law applies to the mass of the substances, not to the number of moles.

Moles are a measure of the quantity of a substance in terms of the number of particles (atoms, molecules, or ions). One mole of a substance contains Avogadro's number (\(6.022 \times 10^{23}\)) of particles.

When we consider a chemical reaction, the mass is conserved, but the moles of reactants and products may not always be conserved. The balance of moles depends on the stoichiometry of the reaction, i.e., the stoichiometric coefficients associated with each reactant and product in the balanced chemical equation.

For example, let's consider the combustion reaction of hydrogen gas and oxygen gas: \[ 2H_{2(g)} + O_{2(g)} \rightarrow 2H_{2}O_{(g)} \] Here, 2 moles of hydrogen gas react with 1 mole of oxygen gas to form 2 moles of water vapor. The mass of the reactants and products are equal: \[ (2 \times 2.016 \ g) + (1 \times 32.000 \ g) = 2 \times (2.016 + 16.000) \ g = 36.032 \ g \] The masses are conserved; however, the number of moles of reactants (2 moles of H2 + 1 mole of O2 = 3 moles) is not equal to the number of moles of the products (2 moles of H2O). In conclusion, the mass of products is equal to the mass of reactants in a chemical reaction due to the law of conservation of mass, but the moles of products do not necessarily equal the moles of reactants. The number of moles depends on the stoichiometry of the balanced chemical equation.