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Unit 14 Power Point
 

Unit 14 Power Point

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    Unit 14 Power Point Unit 14 Power Point Presentation Transcript

    • Activation Energy and Equilibrium Unit 14 Insight Schools Physical Science
    • Getting Chemical Reactions Going
      • Energy is usually needed to break the bonds and get the reaction going.
      • This can be a spark,
      • flame, heat, or even a feather.
    • Getting Chemical Reactions Going
      • Activation energy is the term for the energy needed to break the bonds so new bonds can form
      Activation Energy Energy Released Net
    • Endothermic and Exothermic Reactions Exothermic Reactions Endothermic Reactions Energy Energy is Released to the surroundings Energy is Absorbed from the surroundings Signs Gets Hot Flame Light Gets Cold Dark color Examples Combustion of Methane Cold Pack Photosynthesis
    • Factors That Influence Reaction Rates Affect on Reaction Rate Catalyst Increases the Reaction Rate Concentration of Reactants Increase in Concentration increases the reaction rate because of an increase in particle collisions. Surface Area More surface area increases the reaction rate. (Smaller particles give a greater overall surface area when all the particles surfaces are added together.) Temperature Increase in Temperature increases the reaction rate because of an increase in particle collisions. Stirring Increases the Reaction Rate by increasing the number of collisions.
    • Factors That Influence Reaction Rates Affect on Reaction Rate Catalyst Concentration of Reactants Surface Area Temperature Stirring
    • The Affect of Catalysts
    • Exothermic vs Endothermic
    • Equilibrium
      • The point in a chemical reaction where the rate at which reactants and products are being produced is equal.
    • Equilibrium Concentration
      • Symbolized by [ ] around the compounds or ions
      • [NaCl] [Na + ] and [Cl - ]
      H 2 O
    • Changes in the Equilibrium Concentration
      • Symbolized by [ ] around the compounds or ions
      • [NaCl] [Na + ] and [Cl - ]
      • [NaCl] [Na + ] and [Cl - ]
      • [NaCl] [Na + ] and [Cl - ]
      H 2 O H 2 O evaporating H 2 O added
    • Changes in the Equilibrium Concentration
      • [NaCl] [Na + ] and [Cl - ]
      • [NaCl] [Na + ] and [Cl - ]
      H 2 O evaporating H 2 O added
    • Le Chatelier’s Principle
    • Le Chatelier’s Principle
      • If something is done that disturbs a system at equilibrium, that system responds in a way to counteract what was done.
    • Temperature
      • Increase temperature:
      • favors the endothermic reaction.
      • Decrease temperature:
      • favors the exothermic reaction.
    • Pressure/Volume Changes
      • Increase pressure: favors the reaction producing fewer moles of gas.
      • Decrease pressure: favors the reaction producing more moles of gas.
    • Examples of Changes in Equilibrium CO 2 CO 2 CO 2 CO 2
    • Concentration Changes
      • Increase [reactant]: favors the forward reaction
      • Increase [product]: favors the reverse reaction