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Title: Determine the energy change for the following reaction: Methane + oxygen produces carbon dioxide + water
Description: The energy change for a reaction can be determined by calculating the difference between the energy of the products and the energy of the reactants. In this case, the reaction is the combustion of methane ( � � 4 CH 4 ) with oxygen ( � 2 O 2 ) to produce carbon dioxide ( � � 2 CO 2 ) and water ( � 2 � H 2 O). The balanced chemical equation for the reaction is: � � 4 + 2 � 2 → � � 2 + 2 � 2 � CH 4 +2O 2 →CO 2 +2H 2 O To determine the energy change, we need to consider the enthalpy (heat) of formation of each compound involved. The enthalpy change ( Δ � ΔH) can be calculated using the following equation: Δ � = ∑ � × Δ � � ∘ ( products ) − ∑ � × Δ � � ∘ ( reactants ) ΔH=∑n×ΔH f ∘ (products)−∑m×ΔH f ∘ (reactants) Where: � n and � m are the stoichiometric coefficients from the balanced equation. Δ � � ∘ ΔH f ∘ represents the standard enthalpy of formation for each compound. For this reaction: Δ � � ∘ ΔH f ∘ of � � 4 CH 4 = -74.9 kJ/mol Δ � � ∘ ΔH f ∘ of � 2 O 2 = 0 kJ/mol (since it's an elemental form) Δ � � ∘ ΔH f ∘ of � � 2 CO 2 = -393.5 kJ/mol Δ � � ∘ ΔH f ∘ of � 2 � H 2 O = -285.8 kJ/mol Substituting the values into the equation: Δ � = ( 1 × − 393.5 ) + ( 2 × − 285.8 ) − ( 1 × − 74.9 + 2 × 0 ) ΔH=(1×−393.5)+(2×−285.8)−(1×−74.9+2×0) Δ � = − 787.1 kJ/mol ΔH=−787.1kJ/mol The negative sign indicates that the reaction is exothermic, meaning it releases energy. In this case, the combustion of methane with oxygen to produce carbon dioxide and water releases approximately 787.1 kJ of energy per mole of methane reacted.
Description: The energy change for a reaction can be determined by calculating the difference between the energy of the products and the energy of the reactants. In this case, the reaction is the combustion of methane ( � � 4 CH 4 ) with oxygen ( � 2 O 2 ) to produce carbon dioxide ( � � 2 CO 2 ) and water ( � 2 � H 2 O). The balanced chemical equation for the reaction is: � � 4 + 2 � 2 → � � 2 + 2 � 2 � CH 4 +2O 2 →CO 2 +2H 2 O To determine the energy change, we need to consider the enthalpy (heat) of formation of each compound involved. The enthalpy change ( Δ � ΔH) can be calculated using the following equation: Δ � = ∑ � × Δ � � ∘ ( products ) − ∑ � × Δ � � ∘ ( reactants ) ΔH=∑n×ΔH f ∘ (products)−∑m×ΔH f ∘ (reactants) Where: � n and � m are the stoichiometric coefficients from the balanced equation. Δ � � ∘ ΔH f ∘ represents the standard enthalpy of formation for each compound. For this reaction: Δ � � ∘ ΔH f ∘ of � � 4 CH 4 = -74.9 kJ/mol Δ � � ∘ ΔH f ∘ of � 2 O 2 = 0 kJ/mol (since it's an elemental form) Δ � � ∘ ΔH f ∘ of � � 2 CO 2 = -393.5 kJ/mol Δ � � ∘ ΔH f ∘ of � 2 � H 2 O = -285.8 kJ/mol Substituting the values into the equation: Δ � = ( 1 × − 393.5 ) + ( 2 × − 285.8 ) − ( 1 × − 74.9 + 2 × 0 ) ΔH=(1×−393.5)+(2×−285.8)−(1×−74.9+2×0) Δ � = − 787.1 kJ/mol ΔH=−787.1kJ/mol The negative sign indicates that the reaction is exothermic, meaning it releases energy. In this case, the combustion of methane with oxygen to produce carbon dioxide and water releases approximately 787.1 kJ of energy per mole of methane reacted.
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Title: Determine the energy change for the following reaction: Methane + oxygen produces carbon dioxide + water
Description: The energy change for a reaction can be determined by calculating the difference between the energy of the products and the energy of the reactants. In this case, the reaction is the combustion of methane ( � � 4 CH 4 ) with oxygen ( � 2 O 2 ) to produce carbon dioxide ( � � 2 CO 2 ) and water ( � 2 � H 2 O). The balanced chemical equation for the reaction is: � � 4 + 2 � 2 → � � 2 + 2 � 2 � CH 4 +2O 2 →CO 2 +2H 2 O To determine the energy change, we need to consider the enthalpy (heat) of formation of each compound involved. The enthalpy change ( Δ � ΔH) can be calculated using the following equation: Δ � = ∑ � × Δ � � ∘ ( products ) − ∑ � × Δ � � ∘ ( reactants ) ΔH=∑n×ΔH f ∘ (products)−∑m×ΔH f ∘ (reactants) Where: � n and � m are the stoichiometric coefficients from the balanced equation. Δ � � ∘ ΔH f ∘ represents the standard enthalpy of formation for each compound. For this reaction: Δ � � ∘ ΔH f ∘ of � � 4 CH 4 = -74.9 kJ/mol Δ � � ∘ ΔH f ∘ of � 2 O 2 = 0 kJ/mol (since it's an elemental form) Δ � � ∘ ΔH f ∘ of � � 2 CO 2 = -393.5 kJ/mol Δ � � ∘ ΔH f ∘ of � 2 � H 2 O = -285.8 kJ/mol Substituting the values into the equation: Δ � = ( 1 × − 393.5 ) + ( 2 × − 285.8 ) − ( 1 × − 74.9 + 2 × 0 ) ΔH=(1×−393.5)+(2×−285.8)−(1×−74.9+2×0) Δ � = − 787.1 kJ/mol ΔH=−787.1kJ/mol The negative sign indicates that the reaction is exothermic, meaning it releases energy. In this case, the combustion of methane with oxygen to produce carbon dioxide and water releases approximately 787.1 kJ of energy per mole of methane reacted.
Description: The energy change for a reaction can be determined by calculating the difference between the energy of the products and the energy of the reactants. In this case, the reaction is the combustion of methane ( � � 4 CH 4 ) with oxygen ( � 2 O 2 ) to produce carbon dioxide ( � � 2 CO 2 ) and water ( � 2 � H 2 O). The balanced chemical equation for the reaction is: � � 4 + 2 � 2 → � � 2 + 2 � 2 � CH 4 +2O 2 →CO 2 +2H 2 O To determine the energy change, we need to consider the enthalpy (heat) of formation of each compound involved. The enthalpy change ( Δ � ΔH) can be calculated using the following equation: Δ � = ∑ � × Δ � � ∘ ( products ) − ∑ � × Δ � � ∘ ( reactants ) ΔH=∑n×ΔH f ∘ (products)−∑m×ΔH f ∘ (reactants) Where: � n and � m are the stoichiometric coefficients from the balanced equation. Δ � � ∘ ΔH f ∘ represents the standard enthalpy of formation for each compound. For this reaction: Δ � � ∘ ΔH f ∘ of � � 4 CH 4 = -74.9 kJ/mol Δ � � ∘ ΔH f ∘ of � 2 O 2 = 0 kJ/mol (since it's an elemental form) Δ � � ∘ ΔH f ∘ of � � 2 CO 2 = -393.5 kJ/mol Δ � � ∘ ΔH f ∘ of � 2 � H 2 O = -285.8 kJ/mol Substituting the values into the equation: Δ � = ( 1 × − 393.5 ) + ( 2 × − 285.8 ) − ( 1 × − 74.9 + 2 × 0 ) ΔH=(1×−393.5)+(2×−285.8)−(1×−74.9+2×0) Δ � = − 787.1 kJ/mol ΔH=−787.1kJ/mol The negative sign indicates that the reaction is exothermic, meaning it releases energy. In this case, the combustion of methane with oxygen to produce carbon dioxide and water releases approximately 787.1 kJ of energy per mole of methane reacted.