1065: Chemistry Chapter 5

Atoms

Interact electrostatically, when they form stable systems, the forces are balanced (attractive/repulsive).

• System = interacting atoms

• Energy of the system decreases if it is transferred to the surroundings (collisions)

Interactions

All require energy to overcome

• Typically derived from collisions or photons

Properties of Compounds

Emergent, as they are different from the isolated component’s properties

• Influenced by types of bonds present and spatial organization.

Temperature

A characteristic of a system that details the direction in which heat will flow

• Independent of the size of an object

Thermal Energy

The sum of kinetic and other potential energies of particles in a system

• Depends on the size of the system; systems with more matter have more energy

• Also depends on the type and composition

• Joules (J)

Temperature + Movement

Faster-moving particles have a higher temperature

• It may take more energy to get particles to move at the same average kinetic energy

Translational Motion

Atoms can only move from one place to another in a stagnant line until they bump into something else

Kinetic Energy Formula

KE = ½ mv = 3/2 KT (relationship to kinetic energy)

• m = mass

• v = average velocity

• K = Boltzmann constant

• T = temperature

Kinetic Energy

The energy of motion, directly related to temperature, but a characteristic that a single atom can have

• Molecules in a system can have different kinetic energies

• They change constantly

• Individual kinetic energies are what are critical

Velocities

Speeds and directions, influenced by collisions

Large # of Particles

When involved in a phenomenon, their individual actions aren’t important

• Ex: temperature of pressure

Chemical Reactions

Influenced by the individual collisions of molecules in a population

Populations

Characterized by a distribution with predictable behavior, thus statistical methods are applicable

Distribution of Kinetic Energies

Depends on the temperature of the system

Maxwell-Boltzmann Distribution

Details the distribution of speeds of a certain population at a specific temperature

• Allows us to calculate the probability of a particular molecule moving at a particular speed

• Shape as a function of temperature

• Peaks quickly and falls at higher velocities/kinetic energies

• Probability vs. KE (area = 1)

• Most probable speed (peak) + average speed (higher) increase with temperature

Average Kinetic Energies

…of any gas at the same temperature are equal

• Average velocities depend on the gases mass (lighter = faster)

Uncertainty Principle

There is some uncertainty in a molecule’s energy

• Zero point energy instead of 0K

Monoatomic Gases + Temperature

A measure of the average kinetic energy stored in translation (moving through space)

• Complex System: energy added can increase speed, but also cause vibration, bending, or rotation

  • Distinct and depends on the shape or composition

  • Quantized: only certain energy packets create specific reactions

Determining a Molecule Through Added Thermal Energy

Utilizing molecule-specific energy states to determine identity and structure

• As a molecule moves between vibrational and rotational states

Raising Temperature

More complex structures require more heat because there are more ways for it to vibrate, bend, and rotate

• Thermal energy used for vibrations, rotations, translations, and bending

Heat Capacity

Measures how much energy it takes to change the temperature of a substance by 1 degree, and how much energy a substance can store at a given temperature

• Specific (J/gC) or molar (J/molC)

• Depends on molecular structure and intermolecular forces (IMFs)

  • The more complex of stronger the forces, the more energy that is required

Specific Heat of Water

4.18 J/gC

• Higher than other alcohols because water molecules are smaller, so there are more in a gram, and because water molecules can form more hydrogen bonds, which must be overcome

• On earth water absorbs and releases energy, moderating what would otherwise be drastic temperature changes

Removing Energy From A Gas

Done by lowering the temperature, decreasing the average kinetic energy, causing the molecules to fall out/condense into the liquid state

• Condensation

• Longer interaction times between molecules, where IDF, dipole-dipole, and hydrogen bonds cause them to merge

• Smaller volume, higher density

Boiling/Condensation Point

Temperature at which the gas/liquid phase first appears

• Temperature of the system doesn’t change until almost all the water vapor has condensed into liquid

Evaporative Cooling

Where some molecules in a liquid are in the gaseous state as they have enough momentum to break free, escaping with a little kinetic energy, removed from the liquid system

Liquid → Solid

Removing energy means molecules can no longer sometimes overcome IMFs, causing them to form structures determined by molecular shape and geometry

• More stable

Melting/Freezing Point

Temperature at which a substance becomes a liquid/solid

• Melting: energy is used to overcome intermolecular interactions, not really increase speed

System

The part of the universe we are observing

• Separated from its surroundings

• Important to determine if it is isolate, closed, or open

Isolated Systems

Niether energy or matter moves between the system and its surroundings

Open System

Both energy and matter can enter or leave the system

• All biological systems

Closed System

The amount of matter is constant and only energy can enter or leave

Thermochemistry - 1st Law

Energy cannot be created or destroyed, but it can be transferred from a system to its surroundings and vice versa

Internal Energy (E)

Can change by altering the total amount of thermal energy in a system (q), or the system can do work/have work done on it (w)

• Can only really measure change (∆E), not the exact quantity

• ∆E = q + w

Work (w)

Force over a distance, accompanied by expansion or compression of a system that it usually approximated to 0

• ∆E = q

Enthalpy (H)

A state function of heat change at constant pressure

• Describes thermal energy change (∆H) that doesn’t depend of the path, unlike q

• Applies to systems with a constant pressure with no volume change

  • ∆H = ∆E

State Funciton

A property of a system that does not depend in the path taken to get to that particular state

• Denoted by uppercase letter symbols

Calculating Change in Enthalpy (∆H)

∆H(J) = mass(g) x specific heat (J/gC) x ∆T(K)

• -∆H = exothermic, the surroundings gain energy

• +∆H = endothermic, the system gains energy

• Only applies when the thermal energy is used to increase temperature

2nd Law of Thermodynamics

Any change in the system results in an increase in the entropy (S) of the universe

• Some energy will be changed into a form that is no longer useful

• The system may become more ordered, but that energy is transferred to the surroundings as ∆S > 0

• “Times Arrow” = ∆S

Entropy (S)

A measure of disorder, but also of probability, as to which outcome is most favorable- like heat moving from hot to cold

• The more probable arrangements will occur more frequently, but ordered arrangements can happen, just rarely

• Ex: gases > liquids > solids

• Ex: mixed solutions > unmixed components

• Ex: uniform temperature > hot and cold sites

Calculating Entropy

S = K x lnw

• K = Boltzmann constant

• w = # of possible arrangements of the system (the greater w, the greater entropy)

Gibbs Free Energy Equation

∆G = ∆H - T∆S

Gibbs Free Energy

Used to determine whether reactions are thermodynamically favorable

• +∆G = unfavorable, universal entropy decreases; possible only if paired with a larger -∆G

• -∆G = favorable, universal entropy is increasing

Change in Gibbs Free Energy

Describes how much energy is available to do work, by differentiating energy produced from change from energy lost to the universe as entropy

• -∆G = the direction that the reaction is most likely to occur in

-∆G

Thermodynamically favorable at all temperatures (unaffected)

• -∆H = exothermic

• +∆S = system increases entropy

+∆G

Thermodynamically unfavorable at all temperatures and usually coupled

• +∆H = endothermic

• -∆S = system entropy decreases

Temperature + ∆G

• -∆H (exothermic) + -∆S (entropy decreases): temperature-dependent; as temperature increases, ∆G becomes more positive and less favorable. Favorable at low temperatures.

• +∆H (endothermic) + +∆S (entropy increases): temperature-dependent; as temperature increases, ∆G becomes more negative and more favorable. Favorable at high temperatures.

Standard State

298K, 1 atm, and 1M concentration

• Denoted with a small circle (nought)

System at Equilibrium

∆G = 0

• Ex: A mixture of ice and water is at 0C (273K)

Thermal Energy Notes

• Higher T + same amount of matter → more TE

• Same T + more amount of matter → more TE

Heat

Measure of thermal energy

• Hot → cold (2nd law)

Phase Transitions

• Melting/boiling: IMFs are overcome, PE increases, +∆H, and +∆S (entropy increases)

• Solidification/Condensation: IMFs take over, PE decreases, -∆H, and -∆S (entropy decreases)

Temperature - Phase Transitions

The average kinetic energy remains constant during the transition

• ∆H > 0 = energy spent to overcome the IMFs (PE increases)

• ∆H < 0 = energy released from forming IMFs (PE decreases)

Entropy - Enthalpy + Temperature

∆S = ∆H/T

Second Law Scenarios

The total entropy of the universe must increase, but the signs of the change in entropy for the system and surroundings don’t always have to be opposite

• -∆S system + +∆S surroundings = +∆S total

• +∆S system + -∆S surroundings = +∆S total

• +∆S system + +∆S surroundings = +∆S total

Gibbs Free Energy - Notes

Measures formal entropy change, only using system terms, which are the easiest to measure.

• -T∆S universe → ∆G must be negative (thermodynamically favorable)

• Always when -∆H and +∆S

• Never when +∆H and -∆S

• Sometimes when +∆H and +∆S, or -∆H and -∆S

• Kelvin is always positive, it is only the magnitude that has an impact

Endothermic Reactions

More likely to happen with increased temperature

Longer Reaction - Conditions

Make a reaction go farther by cooling it down (making ∆G more negative)

Equilibrium - Notes

∆G = 0, and the reaction stops because it is happening in both directions at the same rate

• -T∆S = ∆H

• Heating further may reverse the reaction in some instances

• Ex: Temp of phase change