Unit 5 - Reaction Energy

LANGUAGE DEVELOPMENT

• exothermic reaction - releases energy (heat/light/electricity)

• endothermic reaction - absorbs energy in the form of heat

• enthalpy - energy changes that occur during a reaction —> total heat content in a system

• hess’s Law - if a process can be expressed as the sum of two or more steps, the enthalpy change for the overall process is the sum of the ΔH values for each step

• collision theory - in order for a chemical reaction to occur, particles must collide

• activation energy - the minimum amount of energy needed to start a chemical reaction

• catalyst - changes the rate of a chemical reaction without being consumed during the reaction

• reaction rate - the speed at which a chemical reaction takes place

5.1 - Energy in Chemical Bonds

• exothermic reaction (-)

  • releases energy in the form of:

    • heat

    • light

    • sound

    • decrease in temperature

• endothermic reaction (+)

  • absorbs energy in the form of:

    • heat

• enthalpy

  • energy changes occuring during a reaction

    • change in enthalpy —> energy released or absorbed as heat by a system

      • constant pressure

      • not directly measured

      • equation: H(products) - H(reactants)

      • kJ/mol

      • extensive property - depends on the energy of a substance’s particles

        

• thermochemical equation - includes the value of change in H —> right side of the equation

  • physical states of reactants and products

  • adding energy to the equation

  • standardize enthalpy values:

    • STP (standard temperature and pressure) —> 25 degrees Celsius (298 K) and 1 atm

    • formation of 1 mole of the substance

  • burning different types of fuels/combusting them

• synthesis reactions - compound is produced from pure elements

BOTH EQUATIONS MUST RESULT IN THE SAME AMOUNT OF PRODUCTS

• calorimeters

  • well-insulated —> minimize energy transfer

  • constant pressure

    • measures change in enthalpy under constant pressure

    • chemical reactions

    • energy transfer between 2 substances

    • measure temp before/after reaction to determine energy flow

  • bomb calorimeter

    • combustion

    • large amounts of gaseous products

    • large changes in thermal energy

    • contained at fixed volume —> increased pressure

• enthalpy of a reaction

  • heat absorbed at constant pressure —> extensive

• Hess’s Law: Below

• Hess’s Law is based on the law of conservation of energy

  • pure carbon

    • graphite vs diamond —> graphite is standard and diamond is highly pressurized at a hot temperature

    • graphite: enthalpy is 0

    • diamond: can’t be measured because its not at constant pressure

  • assumes energy difference between reactants/products is independent of the route taken to get from one to the other

    

• calculating enthalpy of a reaction

  • 

• bond energies

  • endothermic - breaking bonds

    • total bond energy is greater in the reactants

  • exothermic - forming bonds

    • total bond energy is greater in the products

  • predict the stability of the products of a reaction

  • less bond energy = more potential energy

  • less potential energy = a more stable reaction

• computational chemist

  • use computer simulations to solve chemical problems

    • mathematical models

    • examine relationships among chemical reactions

    • how atoms interact with reactant concentrations, wind, air, temperatures, and seasonal changes in sunlight

    • sub microscopic level

    • improve productivity/efficiency of industrial processes

      • predict how reacting molecules combine under different conditions

      • super computers and computing clusters —> require massive amounts of data

5.2 - reaction rates

• collision theory - particles must collide for a reaction to occur

  • break bonds in reactants

  • get the different particles to interact

  • correct orientation/enough energy —> effective

• increase reaction rates

  • concentration

    • increased = total number of effective collisions increases

    • low: one reaction

    • medium: 4 reactions

    • high: more than 20 reactions

  • temperature

    • particles have more energy and the temperature increases

  • surface area

    • when phases come together (area of contact)

    • heterogenous reaction

    • reaction rate depends on surface area

  • nature of reactants

    • different substances vary greatly in their tendencies to react

    • chemical structure

• activation energy - hill

  • higher than reactants and products

  • transition states —> activated complex

    • brief existence

    • higher frequency of collisions —> more energy to form activated complex

    

• calculating energy requirements

  • Change in E forward: energy of products - energy of reactants

  • Change in E reverse: energy of reactants - energy of products

  • Ea: energy of activated complex - energy of reactants

  • Ea’: energy of activated complex - products

• lowering activation energy

  • catalyst - substanec changing the rate of a chemical reaction without being consumed during the reaction

    • don’t appear in final reaction or transitioned into another state or combined with another element

  • lower activation energy = required to start the reaction

    • photosynthesis and cellular respiration

      • proceed rapidly at relatively low temperatures

• catalyzing changes

  • first catalyst speeds up a reaction

    • removes hydrogen atoms from long chains of carbon atoms

    • double bonds between carbon/polyetheline molecules

      • react more with other compounds

    • reduces activation energy and breaks apart plastic at the double bonds —> hydrocarbon chains at various lengths

      • short enough to be recycled

      • expensive —> recyclable/reusable

• rate law - expresses dependence of reaction rate on the concentration of the reactants

  • R = k [A]^n * [B]^m

• concentration/reaction rate

• chemical kinetics

  • to study how fast chemical reactions happen

    • what affects this speed

  • measures:

    • reaction rate

    • concentration changes over time

    • how rate depends on concentration, temperature, and catalysts

  • predicts: how long a reaction takes to occur

  • reaction efficiency

  • designing safer/faster processes for industrial explosives