A complete chemical reaction is one where all of at least one of the reactants is fully consumed and converted into products. An example involves the reaction of potassium chloride (KCl) with sodium sulfate (Na₂SO₄), but with specific molar ratios to ensure a complete reaction.
Here's a breakdown:
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The Reaction: The reaction between potassium chloride and sodium sulfate produces potassium sulfate and sodium chloride:
2KCl (aq) + Na₂SO₄ (aq) → K₂SO₄ (aq) + 2NaCl (aq)
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Achieving Completeness: For this reaction to be considered "complete" (where all of at least one of the reactants is used up), a specific stoichiometry is required. The reference highlights a situation where the reaction isn't complete. If you have 1 M KCl reacting with 1 M Na₂SO₄, and only 1 M of chloride ions (from KCl) is available, the KCl will be the limiting reactant, and the reaction will stop when all the KCl is consumed, leaving some Na₂SO₄ unreacted.
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Example of Completion: To make this reaction "complete", we need to adjust the amounts of reactants used. For example, if we have exactly the right stoichiometric amounts, say, 2 moles of KCl reacting with 1 mole of Na₂SO₄, the reaction will proceed until at least one of the reactants is completely used up. If you start with slightly more KCl than needed for the Na₂SO₄, the Na₂SO₄ will be completely consumed. The critical factor for a complete reaction is ensuring the reactants are present in the correct stoichiometric ratios, or with one reactant in slight excess.
- If we reacted 2 moles of KCl with 1 mole of Na₂SO₄ and the reaction goes to completion, then all of the sodium sulfate (Na₂SO₄) would be converted to products, leaving perhaps some excess KCl.
In essence, a complete reaction occurs when the limiting reactant is entirely consumed, driving the reaction forward until it reaches its theoretical yield based on that limiting reactant.
