Feb 24 Lecture --thermodynamics

using/releasing energy with heat

heat=q

• exothermic reaction is one where heat is released from a high energy set of reactions

• if it is exiting it is <0

• low energy system converting to a high-energy system is endothermic (heat entering the reaction)

• if it is entering it is >0

forms of energy

• potential energy-stored energy

  • coiled spring

  • chemical bonds (it’s stored in them)

  • cyclist at the top of a hill

• kinetic energy-energy in motion

  • cyclist descending hill

  • releasing coiled spring

• work energy—force acting over a distance

  • work=force X distance

1st law of thermodynamics

• conversion of energy requires that the total energy change in the system and the surroundings must be 0

  • change energy universe=0=change energy system + change energy surroundings

• change in energy=energy in final state-initial energy state

  • E_final>E_initial if energy is entering (+)

  • E_final<E_initial if energy is leaving (-)

convert among common units of energy

• Joule (J) is the amount of energy needed to move a 1-kg mass a distance of one meter

  • 1J=1N*m=1Kg*m²/s²

• calorie (cal) is the amount of energy needed to raise the temp of one gram by 1 deg C

  • Kcal=energy needed to raise 1000 g of water 1 deg C

  • food cal=Kcal

• energy conversion

  • 1 cal=4.18 J

  • 1 Cal=1000 cal

  • 1 kilowatt-hour(kWh)=3.60X10^6 J

measuring heat entering/leaving

heat is a type of kinetic energy

• KE=1/2mv²

heat and temp

• heat is quantified w/ specific heat capacity

  • amount of heat energy required to raise the temp of one gram of a substance 1 deg C

  • C_s

  • units are J/(g*deg C)

• molar heat capacity is the amount of heat required to raise the temp of one mole of substance 1 deg C

• q=(m) X (C_s) X(deltaT)

zeroth law of thermodynamics

heat capacity

important equations

• deltaE=E_final-E_initial

• DeltaE=q(heat)+w(work)

• q=mass of substance X C_s X deltaT

• w=-pdeltaV