chapter 5 lecture notes: solids, liquids, and gases 108 lecture notes chapter 5: solids, liquids,...

Chemistry 108 lecture notes Chapter 5: Solids, Liquids, Gases

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Chapter 5 Lecture Notes: Solids, Liquids, and Gases Chapter 5 Educational Goals

1. Define, compare, and contrast the terms specific heat, heat of fusion, and heat of vaporization. Know the equations that involve these concepts and be able to use them in calculations.

2. Understand the concepts of energy change and free energy change. Know if a process is spontaneous or not based on the free energy change.

3. Know the definition of pressure, vapor pressure, and atmospheric pressure and be able to convert between pressure units of atm, torr, and psi.

4. List the variables that describe a gas (P,V,n, and T) and be able to write and use the equations for the various gas laws.

5. Explain Dalton’s Law of Partial Pressures and define partial pressure. 6. Understand the definitions of density and viscosity. Given the density, and either

the mass or volume of a substance, be able to determine the volume or mass (respectively).

7. Know that a liquid in an open container will boil when its vapor pressure is equal to the atmospheric pressure.

8. Describe, compare, and contrast amorphous solids and crystalline solids. 9. Describe the makeup of the four classes of crystalline solids.

Why are some molecular compounds solid while others are gaseous and others are liquid at room temperature? Competing Powers

• ___________________________forces working to hold particles together as liquids or solids

• _____________________ _______________= Motion = Temperature, work to separate particles

One major factor that is responsible for the varied behavior of solids, liquids, and gases is the nature of the interaction that attracts one particle (atom, ion, or molecule) to another. What forces hold matter together to make liquids and solids? The attractive forces that hold molecules together are called intermolecular forces. 3 Types of Intermolecular Forces

• 1) Dipole-Dipole • 2) Hydrogen Bonding • 3) London Forces (Induced Dipole Forces)


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Other Noncovalent Interactions Noncovalent interactions are interactions that do not involve the sharing of valence electrons (covalent bonding). Other noncovalent interactions due to the attraction of permanent charges.

• 1) Salt bridges • 2) Ion-dipole interactions

• A salt bridge is another name for ionic bond. • Ion-dipole interactions occur between ions with a full charge and atoms with a

partial charge.

Energy meets Matter Phase Changes

Adding energy to liquids will overcome the forces holding the molecules together– boiling Adding energy to solids will overcome the forces holding the molecules together– melting

Calculations Involving Heat Energy Units of Energy

• One __________________________ is the amount of energy needed to raise the temperature of one gram of water by 1°C

• joule – 4.184 J = 1 cal

• In nutrition, calories are capitalized – 1 Cal = 1 kcal

Converting Calories to Joules Example: Convert 60.1 cal to joules Equivalence statement: 1 cal = 4.184 J


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Calculations Involving Heat Energy • One of two things will happen if energy is added or removed from matter (assuming

no chemical change takes place).

– 1) Change the ______________of the substance – Example: melt, freeze, vaporize (boil)

– 2) Change the _______________of the substance

• You can only do ______________ of these at a time!!!

1) Phase Change Calculations Energy calculations for phase changes may be carried out using the tabulated values for:

• ____________ ___ ____________ (symbol = Hfus) for a substance (Table 5.2). • Energy required to melt one gram of a solid • Change sign to negative for freezing (liquid to solid)

• ____________ ___ ____________ (symbol Hvap) of a substance

• Energy required to vaporize one gram of a liquid • Change sign to negative for gas going to liquid

Energy Change = (mass) x (heat of fusion or vaporization)

ΔE = (mass) x (Hfus or vap)

Example: Determine the amount of heat needed to melt 155 g of ice at 0°C, we use the heat of fusion of water (79.7 cal/g) as a conversion factor. Note: No Temperature Change! Ice (0oC) → Water (0oC)


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Group Work: A patient with a fever is sponged with 50.0 g of 2-propanol. How much heat energy is drawn from the patient when 2-propanol vaporizes? (heat of vaporization for 2-propanol is 159 cal/g)

2) Calculations for Changing the Temperature of Matter The amount the temperature of an object increases depends on the amount of _________ added (Q).

– If you double the added heat energy the temperature will increase twice as much.

The amount the temperature of an object increases depends on its _______________. – If you double the mass it will take twice as much heat energy to raise the

temperature the same amount. • Energy calculations may be carried

out using the values for the specific heat of a substance.

• Specific heat is the amount of energy required to raise the temperature of one gram of a substance by one Celsius degree.

Energy required = Specific Heat x Mass x Temperature Change Q = S x m x ΔT ∆ is always: (final) – (initial) (∆T) = Tfinal-Tinitial The table above gives the specific heats of various substances with units of cal/g oC. Specific heats can also be tabulated with units of J/g oC.

• For example, since 1 cal = 4.184 Joules, the specific heat of water is 4.184 J/g oC • If you use cal/g oC in your calculation, the energy (Q) will be in calories. • If you use J/g oC in your calculation, the energy (Q) will be in joules.

Substance Specific Heat (cal/g oC) Water 1.000 Ice 0.500 Steam 0.480 Ethanol (l) 0.586 Copper (s) 0.0924 Aluminum (s) 0.0215 Gold (s) 0.0310


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Example: Calculate the amount of heat energy (in joules) needed to raise the temperature of 7.40 g of water from 29.0°C to 46.0°C

Group Work How much energy needs to be removed from 175 g of water to lower the temperature from 23.0oC to 15.0oC ?

Do this problem, you should get 3.4 x 103 cal for the answer.


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New Topic: Will a Change Occur? Spontaneous vs. Nonspontaneous Changes

• An important question to ask is why some changes are:

– ____________________________(continue to occur once they are started) OR

– ____________ ____________________________(will not run by themselves unless something keeps them going).

– • Energy is the key factor in determining this.

Energy vs. Free Energy The energy (E) of a sample matter depends on the position (potential energy) and velocity (kinetic energy) of every molecule in the sample.

E = Epotential + Ekinetic This is not practical to measure in the lab or to model in calculations! When working at constant temperature and pressure, it is mathematically convenient and experimentally practical to look at the: _________________ ___________________ ____ Just like the energy (E), in nature, given the chance, everything proceeds to the lowest possible free energy (G)!

• The “free energy” (ΔG) of a process can be thought of as the potential for change….

∆G = Gf - Gi A spontaneous process has a negative ∆G and a nonspontaneous process has a positive ∆G.


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Gases and Pressure Properties of Gases Gas molecules or atoms are very ______________ apart from one another. -different from liquids and solids!!

• Gas particles move in a straight line until they collide with another particle or the container wall. .

Gases Have _____________________ Density Because of the relatively large distances between gas particles, most of the volume occupied by a gas is empty space. Gases completely ______________their container.

• Except for a few very heavy gases, most gasses will completely fill their container. Gases Are Highly _________________________________.

Compressibility is the ability to make the space a substance takes up become smaller.

Gases can____________________________.

• Gaseous molecules travel at high speeds in all directions and mix quickly with molecules of gases in the air in a process called diffusion.

• ______________________ is the movement of one substance within another substance until it is evenly distributed.

Gas Pressure Pressure = total _______________________applied to a certain area

larger force = larger pressure

Gas pressure is caused by gas molecules __________________with container walls or surfaces.


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Air Pressure • Constantly present when air present • Decreases with altitude

– less air • Measured using a __________________________

– Column of mercury supported by air pressure – Force of the air on the surface of the mercury balanced by the pull of

gravity on the column of mercury Various Units for Gas Pressure

• 1) atmosphere (atm) • 2) height of a column of mercury (mm Hg, in Hg) • 3) Torr • 4) Pascal (Pa) • 6) pounds per square inch (psi, lbs./in2)

Units we will use for pressure:

• Atmospheres (atm) • Pounds per square inch (psi) • Millimeters of mercury (mm Hg)

– also called torr (1mm Hg = 1 Torr) Relationships: 1 atm = 760. mmHg 1 atm = 760. Torr 1 atm = 14.7 psi Pressure Unit Conversions Example A pressure of 690. Torr is how many atmospheres? (1 atm = 760. Torr) A pressure of 35.0 psi is how many atm? (1 atm = 14.7 psi) A pressure of 812 mm Hg is how many atmospheres? (1 atm = 760. mm Hg)


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Gas Laws

Gas Laws Instructional Goals Understand and be able to use the following gas laws in calculations:

• Boyle’s Law (relationship between pressure and volume) • Charles’ Law (relationship between volume and temperature) • Gay-Lussac’s Law (relationship between pressure and temperature) • Avogadro’s Law (relationship between moles and volume) • Combined Gas Law (relationship between pressure, volume and temperature) • Ideal Gas Law (relationship between pressure, volume, number of moles, and

temperature) The gas laws are the mathematical equations that show the ___________________ between volume, temperature, pressure, and amount of gas. 1) Boyle’s Law

• Boyle studied the relationship between volume and pressure.

• The inverse relationship between pressure and volume is known as _____________ ___________. When the volume decreases, the pressure increases When the volume increases, the pressure decreases • Boyle also noticed that when the pressure and/or volume of a gas is changed the __________________ of the pressure and volume remains the same.

• PxV = Constant Boyle’s Law:

• Remember that when using Boyle’s Law, that the _______________is never

changing. • Only the pressure and volume change.


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Example: The initial volume of the gas in the piston below is 3.00 liters and the initial pressure is 1.00 atm. The piston compressed (at constant temperature) to a new final volume of 1.00 L. What is the final pressure?

Group Work: If the syringe shown has an initial volume of 0.50 mL and the gas in the syringe is at a pressure of 1.0 atm, what is the pressure inside the syringe if your finger is placed over the opening and the plunger is pulled back to give a final volume of 3.0 mL?

P1= V1= P2=???? V2=


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2) Charles’ Law • Charles observed that as the temperature increases, the volume increases and vice

versa. • The direct relationship between temperature and volume is known as _________________ ___________. • Charles also noticed that ratio of volume to temperature of a gas is always the

same. Charles’ Law:

• Remember that when using Charles’ Law, that the ____________________is never changing.

• Only the temperature and volume change. • Temperature must be Kelvin (K).

• Kelvin temperature scale is always positive • K = oC + 273.15

Example: The initial volume of the gas in the piston below is 1.35 liters. The temperature is lowered from 373 K to 250. K (at constant pressure). What is the final volume? Group Work A balloon is inflated to 665 mL volume at 27°C. It is immersed in a dry-ice bath at −79°C. What is its volume, assuming the pressure remains constant?


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3) Gay-Lussac’s Law • Gay-Lussac’s observed that as the temperature increases, the pressure increases

and vice versa. The direct relationship between temperature and pressure is known as ____________ _____________. Gay-Lussac also noticed that ratio of pressure to temperature of a gas is always the same. • Remember that when using Gay-Lussac’s Law, that the __________________ is

never changing. – Only the temperature and pressure change.

• Temperature must be Kelvin (K). Example: The initial pressure of the gas in the container below is .870 torr and the initial temperature is. 300. K. The temperature is raised from 300. K to 1250 K (at constant volume). What is the final pressure?

Group Work An aerosol can containing gas at 25 atm and 22°C is heated to 55°C. Calculate the pressure in the heated can.

K = oC + 273.15


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4) The Combined Gas Law Boyle’s, Charles’s, and Gay-Lussac’s Laws can be _________________mathematically. The relationship between temperature, volume, and pressure is known as the ______________ ____________ ______ . Example: At an ocean depth of 33 ft, where the pressure is 2.0 atm and the temperature is 285K, a scuba diver releases a bubble of air with a volume of 6.0 mL. What is the volume of the air bubble when it reaches the surface, where the pressure is 1.0 atm and the temperature is 298K ? P1= _______ T1= _______ V1= _______ P2= _______ T2= _______ V2= ?????

Avogadro’s Law • Avogadro’s observed that the volume of a gas is directly proportional to the

number of gas molecules. • The direct relationship between moles of gas molecules and volume is known as

__________________ ____________.

• When the number of moles of gas increases, the

volume increases. • Avogadro noticed that ratio of volume to the number of moles

of a gas is always the same.

Remember that when using Avogadro’s Law, that the _______________and ___________________ are never changing.


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– Only the number of particles and volume change. Example: The initial volume of the 3.5 moles of gas in the container below is 1.5 L. Amadeo adds 2.0 moles of gas. (at constant temperature and pressure). What is the final volume?

Group Work A balloon has a volume of 2.4 L and contains 0.12 moles of air. A child blows more air into the balloon until it has a final volume of 3.5 L. How many moles of gas are in the balloon?

Gas Law Summary Combined Gas Law


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The Ideal Gas Law No gas perfectly obeys all four of these laws under all conditions. These assumptions work well for most gases and most conditions. One way to model a gas’s behavior is to assume that the gas is an __________ ________ that perfectly follows these laws. If we combine all these equations, we get the Ideal Gas Law.

The gas constant (R) is a mathematical combination of all the individual gas law constants (Cb, Cc, Cg, Ca)

The Ideal Gas Law is more commonly written as: The previous gas laws we studied involved a ________________in either P, V, T, or n.

• The ideal gas law is used for any gas system, any time. • No changes are involved in the equation

The value of R is:

• When using this equation you must have the following units: • Pressure = atm • Volume = liters • Temperature = K

There are 4 variables in this equation:

In problems, we will always be given 3 of the 4 variables, then solve for the unknown variable.

Example: How many moles of gas are contained in 11.2 liters at 1.00 atm and 0.0°C?


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Group Work: What is the volume of 25.0 grams of oxygen gas (O2) at room temperature (22 oC) and 1.00 atm pressure?

Partial Pressure

Dalton’s law of partial pressure states that the total pressure of a mixture of gases is the sum of the partial pressures of its components.

• The partial pressure of a gas in a mixture is the pressure that the gas would exert if alone

When two gases are present, the total pressure is the sum of the partial pressures of the gases.

P= T= n= V=????


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Liquids Properties of Liquids Viscosity is the resistance to flow.

- It is related to the strength of the non covalent interactions between the molecules that make up the liquid - the stronger the attractions, the thicker the liquid.

- Temperature has an effect on viscosity.

- As temperature rises, the increase in the kinetic energy of the

molecules in the liquid helps the molecules pull away from one another - higher temperature produces lower viscosity.

Vapor Pressure

Due to collisions that take place between particles (atoms or molecules) that make up a liquid, particles at the surface are continually evaporating - being “bounced” off into the gas phase. At the same time gas phase molecules are being trapped and converted to liquid.

The vapor above the liquid causes “vapor pressure”.

• The boiling point of a liquid is the temperature at which the vapor pressure of the liquid equals the atmospheric pressure.

• Liquids boil when their vapor pressure equals the pressure of the air above them.


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Solids

• The atoms, ions, or molecules that make up a solid are held close to one another

and have a limited ability to move around.

• Solids can be classified based on whether or not the arrangement of these particles is ordered (in crystalline solids) or not (in amorphous solids).

Crystalline Solids

• Ionic – consist of oppositely charged ions held to one another by ionic bonds

• Molecular – consist of an ordered arrangement of molecules attracted to one another by

noncovalent interactions • Covalent Networks

– atoms are held to one another by an arrangement of covalent bonds that extends through the solids.

• Metallic – An array of metal cations immersed in a cloud of electrons that spans the

entire crystalline structure.

Metallic Solids Metallic Bonding

• The valence electrons in metals are free to move about the entire crystal of metal nuclei and core electrons.

• We can imagine it like a “sea of electrons” that are bonding the positive nuclei together


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Properties of Metallic Substances Metallic substances are ___________________at room temperature.

• Except for :_______________________

Metallic substances are malleable (they can be hammered or beaten in thin sheets) Metallic compounds are ductile (they can be drawn, pulled, or extruded through a small opening to produce wire. Metallic substances are good conductors of electricity. Covalent Networks A few substances exist as: Covalent Networks

• Atoms are covalently bonded as if it was a huge molecule • Not too many covalent network substances exist

• Examples: Diamond (carbon) and Silicon

Silicon Dioxide (sand)

Amorphous Solids- no regular repeating pattern of ions or molecules. Example: rubber


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chapter 5 lecture notes: solids, liquids, and gases 108 lecture notes chapter 5: solids, liquids,...

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