Chemistry Study Guide Test 10

Energy

the ability to do work or transfer heat

• Energy used to cause an object that has mass to move is called work.

• Energy used to cause the temperature of an object to rise is called heat.

Law of Conservation of Energy

 Energy cannot be created or destroyed.

Kinetic energy

energy objects (including atoms and molecules) possess by virtue of their motion

• Temperature is average kinetic energy

Potential energy

energy an object possesses by virtue of its position or chemical composition.

• The potential energy in molecules is stored in chemical bonds.

Units of Energy

SI unit of energy: joule (J)

• An older, non-SI unit is still in widespread use: the calorie (cal): 1 cal = 4.184 J

Other Energy Units

• Dietary calories are actually “kilocalories,” kcal.

• Large energy changes are often expressed in kilojoules, kJ.

Work

Energy used to move an object over some distance

• w = F × d, where w is work, F is the force, and d is the distance over which the force is exerted.

Conversion of Energy

 Energy can be converted from one type to another, such as potential to kinetic.

• Heat transfer is determined through a process called calorimetry.

Calorimeter

instrument used to measure changes in heat 

Calorimetry

• Initial temperature of the system

• Final temperature of the system

• Mass of the material changing temperature.

• Specific heat is usually given, unless that is what you are solving for.

Heat

• Energy can be transferred as heat. 

• Heat flows from warmer objects to cooler objects.

• Heat can not be measured directly, but it can be calculated.

• The heat associated with chemical reactions is called enthalpy.

Q = mcΔT

• Q is heat, in calories or joule  (make sure the energy terms in Q and c match)

• m is mass, in grams

• c is specific heat, an intensive property

• ΔT is temperature change, in Celsius or Kelvin

Exothermic

heat is released by the system into the surroundings

• ΔH is negative

Endothermic

 heat is absorbed by the system from the surroundings

• ΔH is positive

Enthalpy of Reaction

ΔH = Hproducts − Hreactants

• The change in enthalpy, ΔH, is the enthalpy of the products minus the enthalpy of the reactants

• This quantity, ΔH, is the enthalpy of reaction, or the heat of reaction.

Enthalpy

1. Enthalpy is an extensive property.

2. ΔH for a reaction in the forward direction is equal in size, but opposite in sign, to ΔH for the reverse reaction.

3. ΔH for a reaction depends on the state of the products and the state of the reactants.

First Law of Thermodynamics

1. Energy cannot be created or destroyed.

• Energy can, however, be converted from one form to another or transferred from a system to the surroundings or vice versa.

Thermochemical Equations

Heat is always added to an equation, NEVER subtracted.

• In exothermic reactions, heat is a product.

• In endothermic reactions, heat is a reactant.

Energy Changes Associated with Changes of State

• The heat added to the system at the melting and boiling points goes into breaking intermolecular bonds.

•  Temperature of substance does not rise during a phase change.

Calculating Enthalpy Change

• ΔH is well known for many reactions

• Inconvenient to measure ΔH for every reaction 

• Estimate ΔH for the ΔH published values and properties of enthalpy change

Hess’ Law

“If a reaction is carried out in a series of steps, ΔH for the overall reaction will be the sum of the enthalpy changes for the individual steps.”

• The total amount of energy (heat) change in a chemical reaction is the same, no matter how many steps it takes to get there

Collision Theory

1. Reactions occur when particles collide

2. Most collisions occur 2 particles at a time

3. Not ALL collisions are effective

Reaction Mechanism

series of small reactions that result in a net reaction

Net equation

final equation/answer after adding, subtracting, and cancelling

Intermediate product

chemical formed in one step and used up in another step

• Appears a PRODUCT

• Then, appears as a REACTANT

Rate determining step

slowest step in a reaction mechanism

Catalyst

 makes reaction go faster 

• Appears as a REACTANT

• Then, appears as a PRODUCT

Standard Enthalpy of Formation

 energy change when 1 mole products is made from free elements at 250 degrees Celsius and 1.0 atm

ΔHf

the enthalpy change for the reaction in which a compound is made from its constituent elements (building blocks, like atoms) in their elemental forms.

• ΔH = ΣnΔHf,products – ΣmΔHf°,reactants: Summation Equation

Spontaneous reaction

can proceed without input of energy

• Processes that are spontaneous in one direction are nonspontaneous in the reverse direction.

• Enthalpy, temperature and entropy favor spontaneous reactions

What is always and never spontaneous?

• +ΔH and -ΔS = NEVER spontaneous

• -ΔH and +ΔS = ALWAYS spontaneous

Entropy

thought of as a measure of the randomness of a system.

• Related to the various modes of motion in molecules

• Example: A hot cup of coffee has energy concentrated in one spot; as it cools down, that energy spreads into the room.

• ΔS = Sfinal − Sinitial

• Increases in temperature, amount of material, number of substances present, and atoms per molecule

• Larger and more complex molecules have greater entropies.

Entropy Changes

Entropy changes for a reaction can be estimated in a manner similar to how ΔH is estimated: 

ΔS° = ΣnΔS°(products) — ΣmΔS°(reactants)

Second Law of Thermodynamics

• The second law of thermodynamics states that the entropy of the universe increases for spontaneous processes.

Third Law of Thermodynamics

• The entropy of a pure crystalline substance at absolute zero is 0.

Gibbs Free Energy

• Free energy is energy available to do work (useful energy).

• Negative ΔG: The reaction is spontaneous. It happens naturally (like a ball rolling downhill).

• Positive ΔG: The reaction is non-spontaneous. It needs an outside energy boost to happen (like pushing a ball uphill).

• Zero ΔG: The system is at equilibrium. Nothing changes.

• ΔG° = ΔH° − TΔS°