Thermochemistry and Chemical Kinetics

Thermochemistry and Chemical Kinetics

Thermochemistry is the study of heat changesthat occur during chemical reactions and changes in the state of matter. (Example liquid to a solid)

Heat energy is measured in a quantity called enthalpy
The change in heat energy that accompanies a chemical reaction is represented as ΔH
Hess’s Law provides a method to calculate ΔH: The overall enthalpy change in the equation is equal to the sum of the enthalpy changes for the individual steps in the process

Energy is the ability to do work.

Two types:

Potential energy–Stored energy
Kinetic energy-Energy in motion

Law of Conservation of Energy: In any chemical reaction or physical process, energy can be converted from one form to another, but neither can be created or destroyed

Temperature-a measure of how hot (or cold) something is, specifically it is a measure of the average kinetic energy in the particles of an object

It is an intensive property, which means that the temperature of a sample does not depend on the amount of the sample
KE = ½mv2
Measured in Kelvin (K) (Celsius +273=K)

Heat- (q) The energy transferred between objects that are at different temperatures

It is an extensive property, which means that the amount of the energy transferred as heat by a sample depends on the amount of the sample
Measured in Joules (J) or calories (cal)

Measuring Heat

Calories- The amount of heat required to raise the temperature of one gram of pure water by one degree Celsius.

kcal = 1000 calories or One Calorie
SI units of heat and energy is joules (J)
One J = 0.2390 cal
One cal = 4.184 J
KJ=1000 Joules

Energy of Food

Sources of energy – fat, protein, carbohydrates (Macromolecules)
Body efficiency ≠ calorimeter efficiency
Expressed in Calories

Known balanced equation can determine heat produced from amounts of reactants consumed or products formed.

Specific Heat is the amount of heat required to raise the temperature of one gram of that substance by one degree C. (Cp) (Units: J, kJ, or cal)-all substances have a different Specific Heat (also known as Molar heat capacity)

Basic Equation q=Cp x m x ΔT
q= the heat absorbed or released
C (sometimes seen as Cp) =the specific heat
m=mass in g
ΔT = change in temperature in oC ΔT (T final –T initial)
Lots of variations
n (number of moles)= q ÷ Cp ΔT
n= mass (m) ÷ Molar mass (M) n=m ÷ M
c (calorie) =Cp ÷ M
ΔT =Tf (Final temp) – Ti (Initial Temp)
ΔH (Change in Heat) = C x ΔT
q= n x C x ΔT
C = q ÷ n x ΔT
ΔT = q ÷ n x C

Heat of vaporization Hvapor

Heat of Vaporization: (liquid-gas, gas-liquid)

energy that must be put into a liquid in order to vaporize it into a

gas.

energy must be put in, when a liquid vaporizes, and energy is

released when a gas condenses.

∆Hvap x moles = q (boiling)

Heat of fusion Hfusion (or Hfus)

Heat of Fusion: (solid-liquid, liquid-solid)

a. energy that must be put into a solid in order to melt it

b. needed to overcome the forces holding the solid together

c. the heat given off by a substance when it freezes

∆Hfus x moles = q (melting)

Calorimetry-the measurement of heat related constants, such as specific heat

Calorimeter-a device used to measure the heat absorbed or released in a chemical or physical change

Observes the change in water temperature when chemical reaction occurs.
Temperature change is a result of a phase change or reaction taking place.
Examples:
a. coffee cup – rxn which do not produce pressure (gas)
b. bomb calorimeter – rxn where pressure is produced (gas produced like combustion).

Phase Diagram

Phase diagram– a graph of the relationship between the physical state of a substance and the temperature and pressure of the substance.

Triple point– the temperature and pressure conditions at which the solid, liquid, and gaseous phases of a substance co-exist at equilibrium.
Critical point – the temperature and pressure at which the gas and liquid states of a substance become identical and form one phase.
Sublimation-The change of a solid substance to a gas (vapor) without passing through the liquid phase.
Deposition-The change of a gas (vapor) to a solid without passing through the liquid phase.
Condensation-Gas to a liquid phase
Evaporation-Liquid phase to a gas (vapor) phase

During this process only those molecules with a certain Kinetic Energy can escape from the surface of a liquid

Vaporization-Process where a Liquid is converted to a gas phase

Heating Curve- shows the change in temperature of a substance as heat is added and the substance undergoes phase changes.

Heating Curve of Water

At the melting point (mp): When the ice melts, the temperature stays at 0 oC until all the ice has melted.

At the boiling point(bp): The temperature stays at 100oC until all the water has boiled.

At freezing point (fp): The temperature which a liquid changes to a solid.

Amorphous solid: A solid lacking a crystal unit (Glass)

Chemical Kinetics-The Study of Reaction Rate

Activation Energy-The minimum amount of energy required to start a chemical reaction

Reaction Pathways

1.)Exothermic reactions- the products are lower energy level than the reactants (makes chemical reactions rise in temperature) so the ΔH is negative

Energy is released for exothermic reaction making the product/s hotter than the reactant/s.

2.)Endothermic reactions- The energy of the products is greater than the reactants (chemical reaction lowers in temperature) so the ΔH is positive

Energy is absorbed for endothermic reaction making the product/s colder than the reactant/s.

Catalyst –speeds up a reaction by providing the reactants with an alternate pathway that lowers the activation energy-not consumed in a reaction
Enzymes: Are proteins that act as a biological catalyst-increases speed of chemical.

Inhibitor-slows down and can stop a reaction

Driving Force of Chemical Reactions

Enthalpy- represented as H, is the total energy content of a sample.

If pressure remains constant the enthalpy increases in a sample of matter equal to the energy as heat that is received.
Heat of Reaction, enthalpy change - ∆H
Large, negative enthalpy change results in reaction being driven from reactants to products.
∆Hrxn = ∆Hproducts - ∆Hreactants

Entropy (S) is a measure of randomness or disorder in a system and is a thermodynamic property

Units J/K
Used to determine some endothermic reactions
Diffusion causes a positive change in entropy (ΔS)
Hess’s Law applies to an increase in entropy such as in mechanical shock
The Law of disorder is the natural tendency for a system or systems to move in a direction of increasing disorder or randomness

Gibbs energy (G) G = H – T x ΔS

Also known as free energy
It is the energy in a system that is available to work
ΔG is the change in Gibbs energy
May or may not be efficient (spontaneous or nonspontaneous)

See different types of Enthalpy and Entropy on pages 632-633

The size and direction of enthalpy changes and entropy changes together determine whether or not a chemical reaction is Spontaneous
Energy can only be obtained if the chemical reaction takes place.
Spontaneous reactions naturally occur and form large quantities of a product or products and release free energy.
Reactions that have a low activation energy, reactions involving active elements (like in alkali metals), high temperature reactions, and reactions that produce a precipitate often spontaneous combustion
Photosynthesis is an example: 6CO2 + 6H2O → C6H12O6 + 6O2
Nonspontaneousis the opposite (often do not produce a product or products)
Reversible Reactions↔One side of the reactions is favored over the other. (equilibrium)
Le Chatelier’s principle - Stresses that change the equilibrium include changes in pressure, temperature, and concentration causing the reaction to shift.

Molar Heat Capacity- (C) in a pure substance it is the energy as heat is needed to increase the temperature of one mole of a substance by 1 Kelvin.

Units are J/K x mol
q is the heat needed to increase the temperature of a number of moles of a substance by change in temperate (ΔT)

q = nC∆T

n = moles

C = molar heat capacity

∆T = temperature change (K)

Vapor Pressure

Vapor pressure – is the pressure of the vapor found above a liquid in a closed container.

When the vapor pressure equals the atmospheric pressure, a liquid boils.

Volatile liquids have high vapor pressures, weak intermolecular forces, and low boiling points

Nonvolatile liquids have low vapor pressure, strong intermolecular forces, and high boiling points.

Vapor Pressure

Factors that affect the reaction rate:

Increase with increasing concentration of reactants
Increase with increasing temperature
Catalyst and enzymes increase the reaction and lower the activation energy
Increase pressure will compress the gas in the reaction

Collision theory-is used to describe how chemical bonds are broken and formed in chemical reactions

The reactants are continuously colliding and moving around each other.