Chapters 11 / 12

Causes of Chemical Change

And

Gas Behavior

Learning Objectives:1. Define heat in terms of particle motion.2. Explain the concept of molar heat capacity.3. Describe enthalpy and entropy for chemical changes and

explain the signs of ΔH and ΔS for exothermic and endothermic reactions.

4. Use Gibbs free energy signs to predict spontaneity of reactions.5. Discuss the Kinetic molecular theory and the concept of an

ideal gas.6. Interpret phase diagrams. Identify normal boiling and freezing

points, triple point and critical point.7. Convert between common units for pressure and temperature.8.Perform calculations using:

a. Boyle’s lawb. Charles’ lawc. Combined gas lawd. Ideal gas lawe. Law of partial pressures

9. Differentiate between diffusion and effusion.

A quick review:

• Heat – The total kinetic energy of the particles of a sample of matter.

A bucket of water at 50oC contains more heat than a cup of water at 50oC.

• Temperature – A measure of the average KE of the particles in a sample of matter.

The water molecules in the cup have the same average KE as those in the bucket.

Demo – states of matter - JAVA

A closer look at heat!

KE = ½ mv2 m=mass, v=velocity

Units = J (joules) 1J = 0.239 calories

This is about the amount of heat needed to raise the temperature of 1 gram of water by 1K.

(note: the Calorie used in nutrition is actually 1000 calories!!!)

So, if all the particles are moving with some KE and therefore some velocity, why isn’t the sample of matter moving?

• A more precise definition of heat is the random motion of the particles of a sample of matter.

random motion uniform motion

Enthalpy – The heat, or random motion, of the particles of a substance. Symbol = H

Does heat affect all substances to the same extent?

• See demo!

• Molar Heat Capacity is defined as the amount of heat required to raise the temperature of 1 mole of a substance by 1K. Symbol: C

• q = nCΔT

q=heat, n=moles, C=molar heat capacity, ΔT=change in temperature.

C units = J/K mol

• Essentially, C describes the amount of heat needed to raise the average KE of the particles so that T goes up by 1K.

• The amount of heat depends on the number of particles and also on the types of motion possible – vibration, rotation, translation

• Diagrams:

It is very difficult to measure enthalpy in a substance. Scientists can much more easily measure the change in enthalpy when a substance undergoes a change.

ΔH is the enthalpy change when a substance undergoes a change:

examples: ΔHreac = is the enthalpy change during a chemical reaction

ΔHvap = enthalpy change when a substance evaporates

ΔH for exothermic and endothermic reactions:

Heat is lost to the reaction system so ΔH is negative.

Heat is gained by the reaction

system so ΔH is positive.

There is a second driving force behind any chemical reaction:

EntropyThere are a number of ways to think about entropy.

1. The dissipation or spreading out of heat or energy.

2. The randomness or disorder of a system.

Nature always prefers an increase in entropy!

• S = symbol for entropy

• ΔS = change in entropy for a given system.

ΔS is negative for endothermic reactions

ΔS is positive for exothermic reactions

Why????

Whether a reaction will occur spontaneously or not can be predicted by a quantity called “Gibbs Free Energy” or ΔG. (ΔG= ΔH- TΔS)

ΔG is negative for spontaneous reactions.

If ΔG is positive the reaction will not occur.

If ΔG is 0, then the reaction is at equilibrium.

Chapter 12 -Gases

We live at the

Bottom of an

Ocean!

ISS-260mi

HST-350mi

Kinetic-Molecular Theory

• The KMT explains the behavior of gases under most ordinary conditions.

• Pictures gas particles as tiny billiard balls zooming around, bouncing off of each other and everything else.

• Gas particles all behave the same. The chemical identity does not matter.

The 3 assumptions of KM theory:

1. Particles have no attraction to each other.

2. All collisions are elastic. Energy is transferred between particles-none is lost.

3. Particles take up no space. Volume of individual particles is zero!!!!

An Ideal Gas is a theoretical gas that behaves according to the assumptions of KM theory

Properties of Gases

• Temperature (T) – average KE of particles. Always use Kelvin scale!

• Volume – amount of space the gas takes up. Usually liters or mL.

• Amount – number of particles. Use moles.

• Pressure – A measure of the number of collisions between the particles and surroundings. Force per unit of surface area. Units: atm, torr, psi, Pa

•Units for P:

•1 torr = 1 mmHg (millimeters of mercury)

•1 atm=760 torr (atm=atmospheres)

=14.7 psi (pounds per square inch)

=101,000 Pa (Pascals)

=29.92 inHg (inches of mercury)

All of these can be used as conversion factors!Ex: 14.7psi=101,000 Pa

Practice:

During hurricanes, atmospheric pressure can drop as low as 27 inHg.

Convert this value to torr and to atm.

Car tires are filled to 32 psi. Convert this to torr and atm.

• Convert 150 torr to:

atm

psi

Pa

mmHg

Phase Diagrams

A phase diagram is a graphical plot of P (y-axis) versus T (x-axis) for any substance over a wide range of P and T. The diagram will show under what range of conditions the substance will be in the solid, liquid and gas phases.

Label on the diagram:

Normal boiling point

Normal melting point

Triple point

Critical point

Freezing

Melting

Boiling

Condensing

Sublimating

Depositing

Normal boiling point – Boiling point at 1 atm pressure.

Normal melting point – Melting point at 1 atm pressure

Triple point – The point where the three lines meet. All three phases are in equilibrium at this point.

Critical point- The properties of gas and liquid phases become indistinguishable above this point.

Practice: Label solid, liquid, gas areas. Label the triple point and critical point.Label the T range from 0 to 400. If 1 atm pressure is next to the “u” in the word pressure, what are the normal bp and normal fp?

The Gas Laws!!

But first, a reminder:Direct proportion – two quantities increase

or decrease together. The more you eat, the heavier you get. Amount eaten is directly proportional to weight.

Inverse proportion – two quantities change oppositely to each other. As gas prices go up people drive less. Gas prices are inversely proportional to number of miles driven.

1. Avogadro’s law: The volume of a gas increases with the number of particles of gas (moles). The more gas you have the more space it takes up. (Well, duhhh. And he got famous for this??????)

n is directly proportional to V

• Avogadro’s law also says that the identity of the gas does not matter. At a given T and P, 1 mole of H2 gas will occupy the same volume as 1 mole of any other gas!!!! (CO2, N2…). Remember KM theory says that all gases are the same in their behavior!!!!

2. Boyle’s Law – Pressure is inversely proportional to volume (at constant T).

Another way of saying

This is: PV=k where

K is a constant.

Since the product of P x V is a constant for a given system (at constant T), if we change P or V, then the other factor will also change so that the product is the same!

OR

P1V1 = P2V2

Where P1 and V1 are initial values and P2 and V2 are new values.

This is the form of Boyle’s law that we will use

P and V may be in any units, as long as they are consistent.

Practice:

1. A sample of gas has a volume of 500 mL and pressure of 650 torr. If the pressure is increased to 900 torr what is the new volume?

2. 5.0 L of a gas at 8.5 atm pressure is allowed to expand to 22.0 L. What is the new pressure of the gas?

3. Charles’ Law: The volume of a gas is directly proportional to the temperature of the gas (at constant P).

Or:

V/T =k

Another way of writing Charles’s Law is to consider changing a set of conditions:

=

Where V1 and T1 are initial conditions and V2 and T2 are new conditions.

V can be any consistent unit, however T must always be in K.

V1

T1

V2

T2

Practice:

1. A gas has a volume of 250 mL at 298K. What volume will the gas occupy at 500K?

2. A cylinder in your car engine has a volume of 500 mL at 27oC. When the piston decreases the volume to 70 mL, what is the new temperature?

4. The combined gas law: (Boyle and Charles get together!). This gas law is again for changing conditions but P, V and T are all variables:

=

This law can be used for all changing conditions problems: If a factor is held constant it drops out of the equation:

Ex:

P1 V1

T1

P2 V2

T2

Notes on solving the algebra!

Isolate the unknown on one side on top.

What ever is on top with the unknown goes to the bottom on the other side.

What ever is on the bottom on the unknown side goes to the top on the other side:

Ex:

=

Solve for a, then x, then z

a b

c

x y

z

Practice:

1. A 200 mL sample of gas has a pressure of 0.8 atm at 298K. If the temperature is changed to 400K and pressure to 2.0 atm, what is the new volume?

2. A 2.0 L gas sample is at a pressure of 800 torr and temperature of 50 oC. If the pressure is changed to 500 torr and the volume to 5.0 L, what is the new temperature?

5. The ideal gas law: Includes the amount of gas in moles!

PV = nRT

n= number of moles

R= ideal gas constant 0.0821 L atm/ mol K

To use this law, the units must match those in R!!!!!

V – Liters

P – atm

T – K

n - moles

Practice:

1. A gas sample at 2.0 atm and 400K contains 2.5 moles of gas. What is the volume of the container?

2. If 2.0 moles of a gas occupies a volume of 25 L at a pressure of 800 torr, what temperature is the gas?

3. How many grams of CH4 will occupy a volume of 10 L at a pressure of 0.8 atm and temperature of 300K?

4. 15g of an unknown gas occupies a volume of 12L at a temperature of 250K and pressure of 900 torr. What is the molar mass of the gas?

And Finally!!!!!

The Law of Partial Pressures!!!

For a mixture of gases, the total pressure of the mixture equals the sum of the pressures of the individual gases.

If you have a mixture of gases (gas A, B, and C) in a box. Then the total P in the box is the sum of the P’s of the 3 gases!

PT = PA + PB + PC

This makes sense if you remember that all gases have the same behavior!

Practice:

The total pressure in a container of mixed gases (N2 , O2 , CO2) is 800 torr. If the partial pressures of N2 is 720 torr and O2 is 40 torr what is the pressure of CO2?

Diffusion and Effusion

Diffusion is the spreading out of a gas to fill the available volume.

Effusion is the spreading of a gas through a small opening (think air leaking out of a tire)

The rates of both diffusion and effusion depend only on the mass of the particles.

The heavier the particles the slower the rates. If two particles of different masses have the same KE then the heavier one will have a lower velocity.

KE = ½ mv2