Chem Conservation of Energy Notes

Thermochemistry Energy: The ability to do work Work = Force x distance (you’ll learn about more of that in physics!)  Units for energy: Joule (J): 1000 J = 1 kJ calorie (cal): 1000 cal= 1 kcal= 1 Calorie (in your food) 1 cal = 4.184 J  st The 1 Law of Thermodynamics: The Law of Conservation of Energy In any chemical reaction of physical process energy can be converted from one form to another, but it is neither created nor destroyed A few important vocab terms: System: what we are observing (battery, atmosphere, etc) Insulated: a system is insulated if it cannot transfer past the boundary Surrounding: everything outside of the system State of the system: the temperature, pressure, volume, concentration, and phase of the system. If any of these change, the system changes. Boundary: a separation between the system and the surroundings Forms of energy: **all can be converted into other forms of energy and chemical reactions usually involve at least heat and chemical** Two main types of energy Kinetic energy: energy due to MOTION KE=1/2mv2 -Heat energy is kinetic energy -The hotter something is…the faster it goes Potential energy: depends on position or composition Examples: -gravitational potential energy -elastic (due to compression) -chemical potential energy: stored in chemical bond -Chemical potential energy: in a reaction, some bonds are broken (-ΔPE) and some bonds are formed (+ΔPE) The symbol for heat is q Heat is easy to measure so when we consider a change in energy (ΔE), we are usually discussing heat lost or gained by a system. The amount of heat is equal to the total kinetic energy of ALL of the molecules of a system. The degree of heat, or temperature, is related to the AVERAGE kinetic energy of the molecules of a system. HEAT ALWAYS FLOWS FROM WARMER TO COOLER OBJECTS UNTIL THE TEMPERATURE IS THE SAME THROUGHOUT BOTH OBJECTS. The two objects at the same temperature are then considered to be at thermal equilibrium. There are three types of heat flow: 1. Conduction: the transfer of heat by collisions between the particles in a substance. Solids made of particles of loosely held electrons (like metals) are good conductors 2. Convection: the transfer of heat in a gas or liquid by means of the currents in the heated fluid (like water in a pot or air in a room) Exothermic reactions  Reactions that involve the release of energy  Energy flows from the system to the surroundings  Feels hot because your hand is part of the surroundings  Heat leaves the system (-q)  Ex. Combustion Endothermic reactions  Reactions that absorb energy  Energy flows from the surroundings to the system  Feels cold because your hand is part of the surroundings  Heat is gained by the system (+q)  Ex. Cold packs 3. Radiation: the transfer of heat by way of electromagnetic radiation (like energy from the sun or a heat lamp) Specific heat Specific heat: the amount of heat need to raise 1.0 g of a substance 1°C (so it does not depend on the amount of substance you have). Units are in J/g°C or cal/g°C. Other things about specific heat: **The unit for heat called the calorie comes from the specific heat of liquid water  it takes 1.00 cal of heat to raise the T of 1.00 g of H2O(l) by 1oC  remember, 1 cal = 4.184J  H2O = 1.00 cal/goC = 4.184 J/goC **The larger the specific heat, the “harder” (the more energy it takes!) it is to heat or cool something Heat content q=mcΔT mass (in grams) Change in temperature Specific heat Example problem: How much heat is needed to raise the T of 27g of copper, whose c = 0.387 J/goC, from 50oC to 70oC?     q = mc T q = heat content, m=mass, c=specific heat, ∆T=change in temperature. q = (27g)(0.387 J/goC)(20oC) q = 209 J Example problem: How much heat is released if 250mL of water is cooled from 80oC to 25oC? (1.0 g = 1.0 mL)  q = mc T  q=(250g)(4.184 J/goC)(-55oC)  q = -57,530 J = -57.5 kJ