Chapter 5

Homework:5.17, 5.18, 5.20, 5.24, 5.28, 5.29, 5.30, 5.37, 5.39, 5.51,

5.66, 5.74, 5.84

Gases, Liquids, Solids

• Various forces hold matter together

• The strength of these forces decide the state of matter

• There are attractive forces that hold molecules together

• They are weaker than the forces that hold ions together

• These attractive forces counter act Kinetic Energy in molecules

• With out these forces, the kinetic energy that particles posses would keep them moving mostly in random, disorganized ways

• The kinetic energy increases with temperature

• Therefore, the higher the temperature, the greater the tendency of particles to fly around randomly

• The physical state of matter thus depends on a balance between the kinetic energy and the attractive forces

• At high temperatures, molecules posses a high K.E. and move so fast that the attractive forces are too weak to hold them together

• This is the gas state

• At lower temperatures, a gas condenses to form the liquid state

• Molecules are still moving past each other in the liquid state, but do so much slower

• At even lower temperatures, molecules no longer posses the velocity to move past each other

• This is the solid state

• Molecules have a certain number of neighbors and these neighbors do not change

• The strength of the attractive forces does not change in the 3 states!!!

• Only the Kinetic Energy changes!!

• Most substances can exist in all three states

• A solid, when heated to a sufficient temperature usually melts and forms a liquid

• That temperature is called the Melting Point.

• Further heating causes the temperature to rise to the point at which the liquid boils and becomes a gas

• That temperature is the Boiling Point.

• Not all substances exists in 3 states.– Examples:

Gas Pressure

• We live under a blanket of air that presses down on us

• The amount it presses down on us changes

• We measure the air pressure with a barometer.

• Pressure is most commonly measured in millimeters of mercury (mmHg)

• In can also be measured in torr

• At sea level, the average pressure of the atmosphere is 760 mmHg.

• This is commonly called 1 atm

• 1 atm = 760 mmHg

= 760 torr

= 101,325 pascals

= 28.96 in Hg

- A Manometer is used to measure the pressure of a gas in a container.

Gas Laws• Gas Laws- relationships observed under

different temperatures, volumes, and pressures

• Boyle’s Law- define the relationship between pressure and volume by stating, “At constant mass and temperature, the volume of a gas is inversely proportional to the pressure.”

Charles’s Law

• Defines relationship between temperature and volume by stating, “At constant mass and pressure, the volume of a gas is directly proportional to the temperature in Kelvins (K).”

Gay-Lussac’s Law

• Defines relationship between temperature and pressure by stating, “At constant mass and volume, the pressure is directly proportional to temperature in Kelvins.

Combined Gas Law

• Problem: 3.00 L of a gas is at 2.00 atm. What would the volume be in the pressure was increased to 10.15 atm

Avogadro’s Law

• Defines relationship between mass and volume by stating, “Equal volumes of a gas at the same temperature and pressure contain equal numbers of molecules.”

• The temperature and pressure of gases does not matter when we are comparing them

• However, chemist have selected some standards for convenience:

1 atm = standard pressure

273 K (0oC) = standard temperature

• If a sample is at both standard temperature and pressure, we abbreviate by saying it is at STP

Another Standard

• The volume of one mole of gas (6.022 x 1023 molecules) at STP is 22.4 L

• This is called the Molar Volume

Ideal Gas Law• Combining Avogadro’s law with the Combined

Gas law, we can write an equation that is valid for any pressure, volume, temperature, and quantity of gas.

• This is the IDEAL GAS LAW

P V = n R T

P= pressure, V= volume, n= moles, T=temp

R= Ideal Gas Constant = 0.0821 L atm mol-1 K-1

Ideal Gas constant

• The value of R can be calculated by using the standards mentioned earlier, one mole of any gas at STP occupies 22.4 L

• The ideal gas law holds for all ideal gases

• There are no ideal gases!!!

• The gases that we are surrounded by and work with are real gases

• However, under most experimental conditions, real gases behave sufficiently like ideal gases that we can use the ideal gas law

• Thus we can use PV=nRT to calculate any of the variable as long as we know the other 3.

Example

• If there are 5.0 g of CO2 gas in a 10.0 L cylinder at 350 K, what is the gas pressure within the cylinder?

Gas Mixtures• In a mixture of gases, each of the gas

molecules act independently of all others

• Dalton’s Law of Partial Pressures- the total pressure, PT, of a mixture of gases is the sum of the partial pressures of each gas.

• Partial Pressure of a gas in a mixture is the pressure that the gas would exert if it were alone in the container

Example

• 5.0g of O2, 10.0 g of CO2, and 4 g of N2 are placed in a 10.0L vessel at 300K. What is the total pressure of the vessel?

Kinetic Molecular Theory

• This theory explains the relationship between the observed behavior of gases and the behavior of individual gas molecules within a gas

• There are six assumptions that it makes:

1) Gases consist of particles, either atoms or molecules, constantly moving through space in straight lines, in random directions, and with various speeds

2) The average kinetic energy of gas particles is proportional to the temperature in kelvins

3) Molecules collide with each other, like billiard balls, bouncing off each other and changing direction

4) Gas particles have no volume

5) There are no attractive forces between gas particles

6) Gas particles collide with the walls of the container, and these collisions constitute the pressure of the gas.

These 6 assumptions give us an idealized picture of a gas

Real Gases

• In real gases, there are attractive forces and the particles (atoms or molecules) have volumes

• This is why there are no ideal gases

• But at STP, most real gases behave like ideal gases

Forces of Attraction

• In general, the closer the molecules are, the greater effect of intermolecular forces

• At room temperature and 1 atm, molecules are so far apart, we can basically ignore the intermolecular forces

• But as temperature decreases and/or pressure increases, the molecules get closer together so the effect of the intermolecular forces increases

• It is these forces that cause Condensation and Solidification

• Condensation- change from gas to liquid

• Solidification- change from liquid to solid

• There are three types of intermolecular forces:– London Dispersion Forces– Dipole-Dipole Forces– Hydrogen Bonding

London Dispersion Force• All molecules have intermolecular forces

• We know this because all gases can be condensed to a liquid

• These forces that every molecule posses are called London Dispersion Forces

• They occur due to the spontaneous uneven distribution of electrons at any give moment in time.

Dipole-Dipole forces

• These are the attraction between the positive end of one molecule an the negative end of another molecule

Hydrogen Bonding

• These are a lot like Dipole-Dipole forces, only much stronger

• The positive end is a hydrogen bonded to an Oxygen or Nitrogen and the negative end is a lone pair of electrons on an Oxygen or Nitrogen of another molecule

Strengths

• Hydrogen Bonds are the strongest of the three forces

• The strength of a hydrogen bond range from about 2-10 kcal/mol

• Even though they are the strongest, relative to regular bonds, they are very weak.

• The typical covalent bond between an Oxygen and Hydrogen is 119 kcal/mol

• Hydrogen bonding is very important in biological molecules

• They are what hold the double helix of DNA together.

Properties of Liquids

• As the pressure on a gas increases, the intermolecular forces have more of an impact because the molecules pack closer together

• Once the molecules pack so tightly that almost all the molecules touch or almost touch one another, the gas condenses to a liquid

• Liquids have a set volume, regardless of the container

• They do not expand to fill the container like a gas

• Liquids also have very little space between the molecules, so they are much tougher to compress, unlike gases

• It is so tough that liquids are said to be incompressible

• The density of liquids are greater than gas because the same mass occupies a much smaller amount of space, or volume

• The position of the molecules in a liquid is random

• There is some empty spaces that the molecules can slide into

• So the molecules are constantly sliding past one another, changing their position with respect to their neighbors

• This causes the liquids to be fluid, meaning they have definite volume, but not definite shape

• Liquids have surface properties as well

• One of these is called Surface Tension

• The strength of the surface tension is directly proportional to the strength of the intermolecular forces it possess

• Water has a very high surface tension because of the very strong hydrogen bonding among water molecules

Vapor Pressure and Boiling Point

• An important property of liquids is their tendency to evaporate

• Evaporation:

• Equilibrium- A condition in which two opposing physical forces are equal

Vapor Pressure

• Vapor Pressure- the pressure of gas in equilibrium with its liquid form in a closed container

• The vapor pressure is a physical property of a liquid and changes with temperature

• As you increase the temperature of a liquid, you increase the KE of its molecules so more can escape to the vapor phase

• So as you increase the temperature, the vapor pressure increases, until it equals the atmospheric pressure

• At this point, bubbles of vapor begin to form under the liquid surface and force their way upward

• The liquid begins to Boil!!!

• The molecules the evaporate are those with the highest KE

• The ones left behind have lower KE

• The temperature of a sample is proportional to the average KE, so as the average KE drops, the temperature goes down.

• This is the cooling sensation you get when exiting a pool. Water evaporates off your skin!!

Boiling Point

• The boiling point of a liquid is the temperature at which it vapor pressure is equal to the pressure of the atmosphere in contact with its surface

• The boiling point when the atmospheric pressure is 1 amt is the normal boiling point

Factors that affect Boiling Points

1) Intermolecular forces- the tighter the molecules are held together, the more energy is required to overcome this.

2) Molecular Shape- As surface area of the liquid particles decreases, contact between molecules goes down, so the strength of london forces goes down, meaning the boiling point will decrease

Solids• When liquids are cooled, their molecules come

so close together and the intermolecular attractions become so strong, that the random motion stops, and a solid forms.

• Formation of a solid from a liquid is called solidification or crystallization

• All crystals have a regular shape which often reflects the arrangement of the particles within the crystal

• Solids almost always have a higher density than liquids

• Crystals typically have characteristic shapes and symmetries

• Some compounds have more than one type of solid state– Example: Cabon has 5!!

• The packing of the carbon molecules is different in all 5 states and leads to different properties– Diamond- very hard, very dense– Soot- an amorphous solid meaning its atom’s have

no set pattern and are arranged randomly

Types of Solids

• Know Table 6.4 page 182

Type Make up CharacteristicsExample

Ionic

Molecular

Polymeric

Network

Amorphous

Phase Diagrams

• Phase Diagrams- show all the phase changes for any substance.

• Temperature is plotted on the x-axis

• Pressure on the y-axis