Law of Conservation of Mass

Introduction to the Law of Conservation of Mass

Antoine Lavoisier: Fundamental in establishing the Law of Conservation of Mass. His work confirmed that the total mass remains unchanged in a closed system.

• Notable Experiments:

  • Combustion of Phosphorus and Sulphur:
    • Demonstrated mass consistency when reactions occurred in sealed environments.
    • Utilised quantitative methods to confirm all matter was conserved.
  • Mercury Oxide Experiment:
    • Heated mercury oxide to reveal conservation of oxygen and mercury in a closed flask.
    • Illustrated that reactions involve atom rearrangement, not mass alteration.

• 18th Century Scientific Context:

  • An era characterised by growing empirical evidence refuting outdated theories.
  • Joseph Priestley, a leading proponent of phlogiston theory, observed its disproof through Lavoisier's focus on oxygen.

Introduction to Balancing Equations

• Reactants: Substances existing at the start of a reaction. Example: Hydrogen and oxygen that form water.

• Products: Substances generated by a reaction. Example: Water formed from hydrogen and oxygen.

• Balancing is essential due to the Law of Conservation of Mass:

  • Mass is neither created nor destroyed.
  • Ensures equality in the number of atoms of each element on both sides.

Fundamental Principles

• Conservation of Atoms: List each atom type distinctly to ensure equilibrium.
• Conservation of Mass: Total reactant mass equals total product mass.
• Subscripts remain unchanged: They define compound identity.

Step-by-Step Approach

Conversion and Balancing Method

• Translate word equations into chemical formulas.
• Write the unbalanced equation.
• Tally atoms for each element among reactants and products.
  • Initiate with atoms appearing in the fewest compounds.
  • Balance carbon and hydrogen initially, followed by oxygen.
• Vary coefficients to align the atoms.
• Recount atoms for verification.
• Simplify coefficients to their minimal ratios.

Example Walkthroughs

Methane Combustion

• Equation: CH₄ + O₂ → CO₂ + H₂O
• Balancing Steps:
  • Translate the equation to include chemical formulas.
  • Balance carbon (C), hydrogen (H), then oxygen (O). Results in: CH₄ + 2O₂ → CO₂ + 2H₂O.

Phosphorus Trichloride Synthesis

• Equation: P₄ + Cl₂ → PCl₃
• Balancing Steps:
  • Balance phosphorus (P) initially.
  • Then balance chlorine (Cl).
  • Results in: P₄ + 6Cl₂ → 4PCl₃

Introduction to Stoichiometry and Mass Conservation

• Stoichiometry: Quantitative connections between reactants and products in chemical reactions.
  • Based on the Law of Conservation of Mass, where reactant mass equals product mass.

Importance of Balanced Equations

• Balanced Equations: Preserve mass and atom conservation in a reaction.
  • Analogy: Comparable to a recipe ensuring proper ingredient amounts for desired outcomes.
  • Essential in industries for scaling reactions and obtaining predictable results.

Worked Example: Oxygen Mass Calculation in Hydrogen Combustion

Question: How much oxygen is necessary to completely combust 10 g of hydrogen?

• Balanced Equation:

  • $2H_2 + O_2 \rightarrow 2H_2O$

• Given: 10 g of $H_2$.

• Solution Steps:

  1. Convert mass of $H_2$ to moles: $\frac{10 \text{ g}}{2 \text{ g/mol}} = 5 \text{ moles of } H_2$.
  2. From equation: 2 moles of $H_2$ react with 1 mole of $O_2$.
  3. Calculate moles of $O_2$: $5 \text{ moles of } H_2 \times \frac{1 \text{ mol } O_2}{2 \text{ mol } H_2} = 2.5 \text{ moles of } O_2$.
  4. Convert to mass of $O_2$: $2.5 \text{ moles} \times 32 \text{ g/mol} = 80 \text{ g}$.

Overview of Mass Changes

• Key Concepts:
  • Law of Conservation of Mass: Total mass remains stable during chemical reactions.
  • Stoichiometry: Used to foresee, comprehend, and compute mass changes in reactions.

Mass Changes through Representative Stoichiometry Problems

• Problem Type: Determine mass balance by assessing initial and final amounts.
• Approach Steps:
  • Identify known and unknown amounts.
  • Compose and balance the chemical equation.
  • Utilise molar ratios and stoichiometric coefficients:
    • Relate reactants to products.
    • Perform mass computations to verify mass balance.

Question: Why might mass seem to change in an open system?

Problem Examples

Worked Example: Magnesium Oxide Production

• Steps for Calculation:
  1. Known Quantities: Identify the mass of magnesium and oxygen.
  2. Balanced Equation:
     • Reaction: $2\text{Mg} + \text{O}_2 \rightarrow 2\text{MgO}$
  3. Mass Calculation:
     • Convert mass to moles and verify mass conservation from reactants to products.
  • Outcome Example:
    • Calculation: 48 g of Mg reacts with 32 g of O₂ to produce 80 g of MgO, confirming mass conservation.

Practical Implications

• Systems Examination:
  • Closed Systems: Accurately conserve mass.
  • Open Systems: May display apparent mass losses due to factors like gas escape.
• Real-world Examples:
  • Industrial applications utilise mass conservation for efficiency.
  • Environmental monitoring employs mass balance to track pollutants.

Tips and Tricks for Chemistry Exams

Strategies for Avoiding Mistakes

Balancing Equations

• Common Errors:

  • Miscounting Atoms: Overlooking complex compounds results in incorrect atom counts.
  • Neglecting Verification: Always verify atom counts after balancing.

• Balancing Strategy Checklist:

  1. Identify Elements: Note all elements involved in the reaction.
  2. Initial Tally: Count each atom type on both sides.
  3. Modify Coefficients: Adjust coefficients to balance the atoms.
  4. Reassess Counts: Verify all elements for balance.

Worked Example:

• Example for balancing equation $\text{C}_3\text{H}_8 + \text{O}_2 \rightarrow \text{CO}_2 + \text{H}_2\text{O}$:
  1. Balance Carbon: Start with carbon atoms.
  2. Balance Hydrogen: Follow with hydrogen atoms.
  3. Adjust Oxygen: Increase oxygen to match others.
  • Final Balanced Equation: $\text{C}_3\text{H}_8 + 5\text{O}_2 \rightarrow 3\text{CO}_2 + 4\text{H}_2\text{O}$

Importance of Units and Significant Figures

Quick Guide to Units

Conversion	Multiplier
Grams to Moles	$\text{Molar Mass}$
Moles to Particles	$\text{Avogadro's Number}$

Significant Figures:

• Significance:
  • Crucial for precision. Incorrect application can significantly alter results.
  • Strict adherence to guidelines is necessary for accuracy.

Work Organisation and Exam Techniques

Organisational Strategies

• Time Management Checklist:

  • Allocate time for reading, solving, and reviewing wisely.
  • Modify based on question difficulty and familiarity.

• Problem Approach Flowchart:

  • Read Instructions Clearly: Fully comprehend question requirements.
  • Plan Approach: Choose suitable formulas and data.
  • Execute Solution: Conduct computations carefully.
  • Review: Scrutinise results for possible inaccuracies.

Definitions Recap

• Molar Mass: Mass of one mole of a substance, typically in grams per mole.
• Avogadro's Number: Explanation of constant $6.022 \times 10^{23}$, representing the number of entities in one mole.