Chapter 13

States of Matter

13.1 Gases

The Kinetic-Molecular Theory

Describes the behavior of gases in terms of particles in motion

Size Motion Energy

Gas Behavior

Particle Size

Particle size is small relative to the space that surrounds them

The distance between particles is so large that no attractive or repulsive forces exist between gas particles

Particle Motion

Gas particles are in constant, random motion Particles move in a straight line until they

collide with another particle or the walls of a container

Gas particle collisions are elastic. Although particles in collision can transfer kinetic to eachother, the total kinetic energy of the colliding particles remains constant.

Particle Energy

All particles in a gas have the same mass, but not the same kinetic energy.

Temperature is the measure of the average kinetic energy of particles in a sample of matter

2

2

1mvKE

Mass

Velocity

Explaining the Behavior of GasesLow DensityParticle size is small relative to the space that surrounds them As compared to liquids or solids, gases have much smaller densities due to the fact that fewer particles occupy the same volume

Compression and ExpansionParticle size is small relative to the space that surrounds them

The large amount of space surrounding the particles allows particles the room to move closer together as they are compressed

Diffusion and Effusion

Gas particles are in constant, random motionGas particles tend to move from a high area of concentration to a low area of concentration. The rate at which they diffuse is dependent on their mass.

B

A

MolarMass

MolarMass

RateB

RateAComparison of the diffusion rates of

two gases

Graham’s Law of Effusion

Gas escaping through a small hole

MolarMassteEffusionRa

1

Practice Problem

What is the ratio of effusion rates for Nitrogen (N2) and Neon (Ne)?

By: Jordan Gaffin, Sam Bear, and Keith Zubrow

What is Pressure?

Pressure is defined as force applied

per unit area. We measure air and

atmospheric pressure with a

barometer.

Measuring Air Pressure

Italian physicist Evangelista Torricelli was the first to demonstrate that air exerted pressure.

He invented a device called the barometer that assisted him in measuring pressure.

Gas Pressure

A manometer is an instrument used to measure gas pressure in a closed container.

It is a flask that is connected to a U-tube that contains mercury.

Unit Compared with 1 atm

Compared with 1 kPa

Kilopascal (kPa) 1 atm = 101.3 kPa

Millimeters of mercury (mm Hg)

1 atm = 760 mm Hg I kPa = 7.501 mm Hg

Torr 1 atm = 760 torr 1 kPa = 7.501 torr

Pound per square inch (psi or lb/in)

1 atm = 14.7 psi 1 kPa = .145 psi

Atmosphere (atm) I kPa = .009869 atm

Units Of Pressure

Units of Pressure

The SI unit of pressure is the pascal (Pa). 1 pascal = the force of 1 Newton per square

meter. The pressures measured by barometers and

manometers can be reported in millimeters of mercury (mm Hg).

There is also a unit called the torr which is equal to 1mm Hg.

Often air pressure is reported in a unit called atmosphere (atm).

1 atm = 760mm Hg

Conversions

To convert the different measures of pressure, we use factor label.

Use the given measurements: 1 atm= 101.3 kPa = 14.7 psi = 760 torr

Dalton’s Law of Partial Pressures

This law states that the total pressure of a mixture of gases is equal to the sum of the pressures of all the gases in the mixture.

Dalton’s law of partial pressures can be summarized as:

Ptotal = P1+P2+P3+... Pn

The P’s in Daltons Theory

Ptotal is the total pressure of a mixture of gas.

All the other P’s (P1, P2, P3, Pn…) are the partial pressures of each gas in the mixture.

Example

A mixture of Oxygen, Carbon Dioxide, and Nitrogen has a total pressure of .97 atm. What is the partial pressure of O2 if the partial pressure of CO2 is .70 and the partial pressure of N2 is .12?

P.97atm = P N2 .12 atm + P CO2 .7atm + P O2 x

P O2 = .97 atm - .70 atm - .12 atm

P O2 = .15 atm

Section 13.2Forces of Attraction

By Dorothy Raginsky, Rima Naseer, Luke Morreale, Emil George, Greg Klein

Table of Contents

Differences between intramolecular forces and intermolecular forces

Dispersion Forces with examplesDipole-dipole Forces with examplesHydrogen Bonds with examples

Intramolecular Forces

The attractive forces that hold particles together in ionic, covalent, and metallic bonds.

Intra- means “within”.

Intermolecular Forces

They hold together identical particles or two different types of particles.

There are three types of intermolecular forces: Dispersion forces Dipole-dipole forces Hydrogen bonds

Some intermolecular forces are stronger than others, but all are weaker than intramolecular forces.

Dispersion Forces

They are weak forces that result from temporary shifts in the density of electrons in electron clouds.

Dispersion forces are sometimes called London forces after the German American physicist who first described them, Fritz London.

They are the weakest intermolecular force. They are the dominant force of attraction

between identical, nonpolar molecules.

Examples

Fluorine, chlorine, bromine, and iodine exist as diatomic molecules.

Recall that the number of nonvalence electrons increases from fluorine to chlorine to bromine to iodine. Because the larger halogen molecules have more electrons, there can be a greater difference between the positive and negative regions of their temporary dipoles, and, thus, the stronger dispersion forces. This difference in dispersion forces explains why chlorine and fluorine are gases, bromine is a liquid, and iodine is a solid at room temperature.

Dispersion Forces

Dipole-dipole Forces

The attraction between oppositely charged regions of polar molecules are called dipole-dipole forces.

Polar molecules contain permanent dipoles; that is, some polar regions of a polar molecule are always partially negative and some regions are always partially positive.

Neighboring polar molecules orient themselves so that oppositely charged regions line up.

Examples

When hydrogen chloride gas molecules approach, the partially positive hydrogen atom in one molecule is attracted to the partially negative chlorine atom in another molecule.

Dipole-dipole Forces

Hydrogen Bonds

A hydrogen bond is one special type of dipole-dipole attraction that occurs between molecules containing a hydrogen atom bonded to a small highly electronegative atom with at least one lone electron pair.

For a hydrogen bond to form, hydrogen must be bonded to either a fluorine, oxygen, or nitrogen atom. These atoms are electronegative.

Examples

In a water molecule, the hydrogen atoms have a large partial positive charge and the oxygen atom has a large partial negative charge. When the water molecules approach, a hydrogen atom in one molecule is attracted to the oxygen atom on the other molecule.

Hydrogen Bonding

Section 13.3 Liquids

Ryan Desch, Dave Derr, Colin Drummond,Kyle Giordano, Jake

Long, Pratt Templeton

Density and Compression

At one atmosphere of air pressure liquids are more denser than gases.

Change in volume for liquids is much smaller than gas because liquid particles are already tightly packed.

An enormous amount of pressure must be applied to reduce the volume of liquid by even a few percent.

Fluidity

Fluidity- The ability to flow

A liquid is considered as fluid because it can flow.

Liquids can defuse through other liquids due to their fluidity.

Liquids are less fluid than gases.

Viscosity Viscosity- A measure of the

resistance of a liquid to flow The particles in liquid are

close enough for attractive forces to slow their movement.

Viscosity decreases with temperature (Increase in temperature = increases the average kinetic energy of molecules).

For Example- When you pour a tablespoon of cooking oil into a frying pan the oil tends not to spread across the bottom of the pan until you heat the oil.

Surface Tension Surface Tension- The

energy required to increase the surface area of a liquid by a given amount.

The stronger the attractions between particles, the greater the surface tension.

Surfactants- Compounds that lower the surface tension of water.

Capillary Action When liquid is placed in a narrow container the surface forms a concave meniscus (The surface dips in the center).

This movement of the liquid, such as water, is called capillary action or capillarity.

Cohesion- The force of attraction between identical molecules

Adhesion- The force of attraction between molecules that are different.

Solids

SolidsJoe BruchBria CollinsMelanie SpivackSiobhan BradyLaura Betterly

Density

Solids are more dense than liquids and gases All the solids remain a compact phase, where

the arrangement of molecules is very ordered and well defined. With this in mind it is surprising that there exists some solids (few) less dense than liquids.

P=m( mass)/v (volume). This is the equation for density

Crystalline Solids Are arranged in fixed geometric patterns or lattices. Ex. ice methanol and sodium chloride They have an orderly arranged units and are practically

incompressible Crystalline solids, also, show a definite melting point They passed rather sharply from solid to liquid state There are various crystalline forms which are divided into

six crystal systems or shapes: cubic, tetragonal, hexagonal, rhombic, monoclinic, and triclinic

Molecular Solids

a solid composed of molecules held together by relatively weak intermolecular forces

are soft, and generally have low melting and boiling temperatures

Most solids are nonconducting when pure and are insoluble in water, but soluble in non-polar solvents

Ex. Sulfur, ice, sucrose, and carbon dioxide solid

Metallic Solids

More than 80 elements in the periodic table are metals

High densities Large atomic radii Low ionization energies Low electronegativities Usually, high deformation Malleable Ductile Thermal conductors

Covalent Network of Solids

-Have no discrete molecular units

-Held together by conventional covalent bonds

-Continuous network of bonded atoms

examples – diamond and quartz

Amorphous solids

is a solid in which there is no long-range order in position- molecules are arranged in a random manner

common window glass and cotton candy and plastic are examples

Amorphous solids also unlike crystalline solids. They do not have definite melting points

Phase ChangesBy, Taylor Brink, Lindsay

Coloracci, Amanda Couch, Emily Rorer, Callie Wendell, and Jeff

Wright

Endothermic Changes (within)

Melting

Hot flows into cold Heat disrupts the hydrogen bonds holding

the ice together These bonds are strong, therefore, lots of

energy is required Melting Point- temperature at which a

solid becomes a liquid- hard to determine exact point

Vaporization vs. Evaporation

Vaporization liquid changes into a gas (vapor) the gradual input of energy equals the amount of

molecules escaping from the surface Vapor pressure- pressure exerted by a vapor over a

liquid Surface molecules are less attracted to other molecules

Evaporation molecules escaping from the surface better in warmer temperatures because there is more

energy all evaporates over time depending on the amount of

water and energy

Boiling Process

Boiling point- temperature where vapor pressure equals the atmospheric pressure

At this temperature, all molecules have enough energy to vaporize, where as only the surface molecules could before

Sublimation

Transition from a solid to a gas, skipping a liquid examples:

• Carbon Dioxide (dry ice)• Air fresheners• Ice in a freezer

Happens when the pressure is low

Exothermic Changes(outside)

Condensation

Transition from gas to liquid Hydrogen bonds begin to form, and

the vapor looses energy examples

• Sides of a cold glass• Dew• Fog• Rain

Deposition

Transition from gas to solid Vapor settles on to a solid

Examples • Frost• Snowflakes

Freezing

Transition from liquid to solid Water looses energy (heat) and

hydrogen bonds form, stopping the motion of water molecules

Freezing point- temperature where liquid turns to solid