Energy-Everything we do in chemistry deals with energy-With the exception of the sun, most energy in our lives comes from chemical reactions.

-We are a walking chemical reaction.Definitions:Thermodynamics: The study of energy and its transformations.Thermochemistry: Relationships between chemical reactions and energy changes involving heat.Energy: The capacity to do work or transfer heat.Work: Energy used to cause an object to move against a force.Heat: Energy used to cause the temperature of an object to increase.Two forms of EnergyKinetic Energy: Energy of motion, it's in objects like atoms, matter, and the transfer of heat. Ex. Ball rolling, car movingKinetic energy = 1/2mv^2. m= mass, v= velocityPotential Energy: The energy an object posseses by virtue of its position or composition- "Stored Energy"-Potential Energy=mgh (g=gravitational constant=9.8m/s^2 on earth), m=mass, h=height-Examples: Ball sitting on a hill, gas in car, electrostatic energyUnits of Energy-SI derived unit: Joule, and other unit: calorie (lower c)

-calorie: Amount of energy needed to raise temperature of 1g water by 1ºC.-1calorie=4.184J-Calorie: 1000calories=1kcalSystem Vs. SurroundingsSystem: Part of the universe we are interested in studyingSurroundings: The rest of the universe-Example: dissolve NaOH in waterThe system is the NaOH. The surroundings include water, and flask, etc.Energy Transfer-Transferred in 2 ways

-1. Moving object through force (work)-2. Cause temperature change. <---

-In #1, work is done. So work=force*distance, and force=mass*gravity -The heavier the weight the more work is done, and the higher the weight the more work is done.

-In #2: energy is transferred as heat, and heat is the energy transferred from a hotter to colder object.pg. 163:First Law of Thermodynamics-Better known as the Law of Conservation of Energy. Energy cannot be created nor destroyed. There's a constant amount of energy in the universe, it's just being transformed into one type or another. -Any energy lost by a system is absorbed by the surroundings, and vice versa.Internal Energy (of a system)-The sum of all the kinetic and potential energies of all components of a system.-Very difficult to measure internal energy.

-Abbreviated as E.-NaOH soln: you'd consider electrostatic energy, the charge of electrons, chemical energy in ionic bonds, molecular movement, internal vibrations, energy of nuclei, etc.

-What we can measure is ∆E (change in energy). ∆=Alt-J-Finding ∆E:

-Can be determined w/o knowing values for initial and final energy.-It has sign. (+ or -) -If ∆E is +, the system gained energy from surroundings.

Calculating ∆E.

-A system exchanges energy with its surroundings in 2 ways: heat and work.

-∆E=Q+W (q is heat, and w is work).-q and w have signs.

-Q is + when heat transfers from surrounedings into the system.-Q is - when heat is lost to surroundings.

-w is + when work is done on the system by the surroundings.-w is - when work is done by the system on the surroundings.kk

Endothemic processes-Processes that absorb heat from surroundings, sign is +Endothermic: Processes that absorb heat from surrounding, sign is +

Exothermic Processes

-Specifically for heat.-Pfocesses that give off heat to surroundings. Sign is -.

-A value of heat to it.Enthalpy

-Abbreviation H.-Basiclally heat energy-Relationship defined as H = E + PV, E is internal energy, PV is work known as PV work.-Again we can find ∆AH, the change is enphalpy.∆H = ∆E + PV-For most chemical reactions and processes, the PV is very small, done

at constant temperature, so the ∆H is the best absorbed.Sign of ∆

-If it's positive, there is + enthalpy and heat is trasnfeered from the surroundings to the system and the process is endothermic.-If ∆H is -and there is- enthalpy and heat is transferred from the system to the surroundings and the process is exothermic.

To start reactions, you need activation energy. Separating atoms by putting energy in and breaking the bonds. When the bonds have been broken, then atoms can recombine.exothermic, so more heat out than inOverall, heat is given off. Also means that there is greater enthalpy in reactants than products. Diagram of this Exothermic reactionDifference between initial line and final line is the heat given off.

Book style: And Now, an endothermic reactionHere's more heat in than out. It is absorbed and therefore, it's an endothermic reaction. There's greater enthalpy in the products.Thermo-chemical laws1. Value for ∆H in a Reaction is directly proportional to the mol values of all reactants and products.The 890 kJ heat are given off when 1mol CH4 burns. It's directly related to the mol ratios. This is also a conversion factor!2. The ∆H in a reaction is equal in magnitude (but opposite in sign) to the ∆H of the reverse reaction. In other words, take the reactants and make them the products, and vice versa. 3. Enthalpy content of a substance varies as its physical state changes. Water and water vapor don't have the same enthalpy content. 46. c. If a reaction occurs spontaneously, exothermic ones tend to be

spontaneous. If we go the other way (endothermic reaction), it's not feasible, because you can do it, but it needs a lot of energy to be put in. Heat Capacity- amount of heat required to raise the temperature of a substance by 1ºC or 1K.

-J/ºC or kJ/ºC..."raise the temperature of one mole of a substance..."-Molar heat capacity...."raise the temperature of one gram of substance..."-Specific heat capacity (short: specific heat)Specific heat of water= 4.184J/g ºC (this is also a calorie)-It takes 4.184J of heat to raise temp of 1g water by 1ºC.-Joules, per grams, degree CelciusDifferent substances have different specific heats.-Water or metal? Which has higher specific heat?

-Water. Because metals conduct heat much easier than water. FOr exmaple, if we put them both outside, the metal heats up faster. It takes less heat energy to raise the temp of the metal.

q=mass*∆T*specific heat.(∆T= T_final - T_initial)- change in temperatureCoffee-cup calorimetry-We can calculate the amt. of heat being given off.-That heat's given into surroundings.-Let's control the surroundings with a coffee cup to insulate it. The surroundings are now simply the water.-If we measure heat going into water, that's same amt. heat lost by the solution. Bomb calorimetry- use to show loss of heat from burning soemthing. q = -C*∆T

Most reactions don't take place in one step. -Add up the ∆H's in each step to get the overall ∆H. To do this, each of those reactions need to add up to the overall equation. Not usually easy.1. Correct Side?2. Correct coefficients?3. Cancel out?

-If all the moles of everything cancel out, you're fine.4. Add up?12. Practice problem.

1. Correct side? No, N2O5 needs to be on the right of the arrow. Reverse the second equation.2. Correct coefficients? No, 3rd equation has 1/2N2, which is just N2 in the final equation. Multiply entire 3rd equation by 2.3. Canceled out? No, you can cancel out anything that is present in equal amounts both before and after the arrows. You can cancel:2HNO3: before the arrow in eq2, after arrow in eq3.H2: before the arrow in eq1, after arrow in eq3, but only after these steps:

-Flip the first equation so H2 is after the arrow.-Divide eq1 by 2, so it's not 2H2, and instead it's H2.

4. Add the new ∆H's. Answer is 11.3kJ.How to determine ∆H of reactions?1. You could use calorimetry.2. You can use Hess' Law3. New way: Enthalpies of formation (Heat of formation)Heat of Formation- table on pg. 1059The enthalpy change associated with the formation of 1 mole of a compound or polyatomic ion from its elements (in their most standard states).Heat of formation of elements in their standard state is 0.

∆H º of O