Chem Chapters 13, 14 Williams

States of Matter and the Behavior of Gasses (Gas Laws)

kinetic energy

kinetic energy

energy of motion

KE formula

• KE = ½ mv^2 = 3/2 kT

  • k = boltzmann’s constant (maxwell - boltzmann’s distribution)

  • T = temperature in kelvin

  • m = mass (kg)

  • v = velocity (m/s)

  • v 𝜶 T

“proportional to”

𝜶

sublimation

solid to gas

deposition

gas to solid

melting

solid to liquid

freezing

liquid to solid

evaporation

liquid to gas

condensation

gas to liquid

vaporization

solid or liquid to gas when boiling

gas assumptions

• Particles are small, hard spheres with insignificant volume

• Particles move in a straight line and random motion

• Particles colliding with other particles or the container walls are elastic

  • In elastic collisions, particles bounce off each other like billiard balls

  • When inelastic collisions take place, particles can stick together (bond) or deform from the impact (cars)

gas pressure

The result of gas particles striking the surface of the container

what is pressure measured in?

• Pressure (kPa) = force/area = newtons/m2

• Measurement unit for pressure: SI unit for pressure: Pascal (Pa) and NY State also uses atmospheres (atm)

  • Units for pressure: in Hg, mm Hg, PSI, torr, bar

  • 1 mmHg = 1 torr

  • 30 in Hg = 760 mmHg = 760 torr = 14.7 psi = 1 atm

devices used to measure gas pressure

barometer, manometer, gauge

barometer

Pressure of atmosphere pushes down on mercury + based on the air pressure the mercury will rise or fall certain levels

manometer

Fill with tested gas + high density liquid (aka mercury) to determine pressure - if pressure were the same, the blue levels would be parallel to each other

gauge

Pressure gauges work through a Bourdon tube, a hollow piece of metal. When a gas or liquid enters the tube, it expands and pushes a lever. The distance the lever moves is proportional to the gas or liquid pressure inside the hollow tube.

nature of liquids

Intermolecular forces dictate the properties of liquids

• Van der Waal forces

  • Hydrogen bonding

• London Dispersion forces (LDF)

  • Dipole - dipole

  • Dipole - induced dipole

  • Induced dipole - induced dipole

dipole

the electrons spending more time at one end of the molecule because of the charges attracting

evaporation

• Particles at the surface of a liquid have sufficient kinetic energy to “escape” the intermolecular forces

  • Liquid temperature is well below the boiling temperature

  • Temperature of liquid particles follow a “normal distribution”

    • Particles on the right have enough energy (temperature) to overcome the intermolecular forces holding them in the liquid

boiling point

temperature when the atmospheric pressure is equal to the vapor pressure

• At the boiling point, there is gas inside the liquid (bubbles)

  • The boiling point of a liquid is not a specific value

    • The boiling point changes based upon the atmospheric pressure

    • The “Normal Boiling Point” is the temperature that a liquid boils when the pressure is 101.3 kPa or 1.00 atm

Vapor pressure

• When particles “escape” the liquid, they become gas particles

  • These gas particles generate a pressure

  • As temperature increases, the number of particles escaping the liquid increases

    • More particles mean that they will strike the surface of the pressure sensor

    • Measurement of pressure increases

    • Manometer commonly used to measure vapor pressure

• In an open container, the gas particles can drift away which allows more liquid particles to evaporate

• In a closed container, the gas particles are trapped. If the gas particles lose enough energy through collisions, they can be “recaptured” by the liquid

Establishes a dynamic equilibrium

model for solids

Solids reflect an organized arrangement of the particles in fixed locations

• It is important to remember that the particles in the solid are in constant vibrational, rotational, and very limited translational motion

  • The lower the temperature goes, the slower these movements become

  • At 0 K, all movement of the particles theoretically “stops”.

How does 0 K violate the conservation of energy law?

law says energy can’t be created nor destroyed - 0K is saying that there is essentially no energy left which can never be true

unit cell

All solids are build on one of these structures

• Each unit cell has unique properties

• When exposed to heat, solids can change the unit cell

Simple Cubic

tends to be fragile

Face centered cubic (FCC)

tend to be strong but brittle

Body centered cubic (BCC)

tends to be strong and elastic

crystal systems

All solids are build around one of these seven crystal systems - cubic, tetragonal, monoclinic, orthorhombic, triclinic, hexagonal, trigonal

allotrope

substance that is made from the same elements but can form two or more different structures

amorphous solid

a substance that lacks an ordered internal structure

• Glass is an example of an amorphous solid

Phase Diagrams

Shows the relationship between phase changes as a function of temperature and pressure

• Every substance has a phase diagram  describing:

  • melting/freezing point

  • boiling /condensation point

  • Triple point: point where all three phases coexist

Sublimation range: range when change from gas to solid

critical point

where substance enters plasma phase

property of gas: compressibility

Since gas particles are “far” apart, they can be forced closer together without dramatically altering the properties of the gas

• Pressure and density increase when a gas is compressed, but it remains a gas

What factors affect gas pressure?

Kinetic Theory of Gasses can be used to understand these factors

• Amount of gas

  • The quantity of gas inside the container

• Volume of the container

  • The volume of the container

• Temperature of the gas

  • Temperature is a measurement of the average kinetic energy

Boyle’s Law

as volume increases, pressure decreases

• Pressure (kPa/atm) vs Volume (L) - TEMPERATURE AND N ARE CONSTANT

  • P1 x V1 = P2 x V2

  • Indirect relationship

Charles’ Law

as temperature increases, volume increases - PRESSURE AND N ARE CONSTANT

• Volume (L) vs Temperature (K)

• V1/T1 = V2/T2

• Direct relationship

• Must use Kelvin scale because does not have negative values

Gay-Lussac’s Law

as temperature increases, pressure increases

• Pressure (kPa/atm) vs Temperature (K)

  • P1/T1 = P2/T2

  • Direct relationship

• Must use Kelvin scale

Combined Gas Law

When all three responses are combined with a fixed amount of gas

• Moles of gas are constant

• Allows for each variable to be manipulated independently

• PV/T = K

The Ideal Gas Law

expands the Combined Gas Law to include the variable of matter - moles

• PV/T = constant

• PV/nT = R

  • n is the mole value of the gas

  • R is the universal gas constant

  • R = 8.31L•kPa/mol•K = 0.0821 L•atm/mol•K

Dalton’s Law

The total pressure is the sum of the partial pressures of each gas in the mixture

• PT = P1 + P2 + P3 + …

Graham’s Law

Describes the rate one gas will move through another gas

• RateA/RateB = √ (molar massB/molar massA)

Effusion

particles having a barrier, when removed they randomly mix - or particles escaping through a hole in barrier (pinhole) due to differences in pressure inside and outside

Diffusion

gas particles move from higher concentration to lower concentration to reach equilibrium