Part 01 - Notes: Intermolecular Forces Honors Chemistry Unit 10 - IMF & Cond States Learning Objectives: Propose reasoning for the relationship between attractive forces and the boiling point and melting point of sub- stances, among other properties. Develop the link between boiling points, attractive forces and (a) mass and (b) molecular polarity. Text Reference: Chang and Goldsby (Chemistry: The Essential Concepts, McGraw-Hill, 2014) pp. 403−409. Intramolecular Versus Intermolecular Forces Two different types of forces the occur between molecules: INTERMOLECULAR and INTRAMOLECULAR What is an INTRAMOLECULAR FORCE? Where do these occur? What are these? What is an INTERMOLECULAR FORCE? Where do these occur? Why are these significant? Discuss the intermolecular forces involved in solids, liquids, and gases. Condensation is the process by which a gas becomes a liquid. When two molecules in the gaseous state come close to each other, a momentary attraction forms between them. Whether the two particles condense depends upon their kinetic energy (temperature) and the strength of their intermolecular forces. Increasing the pressure can also cause condensation by forcing the particles of a gas close enough to allow the intermolecular forces to hold them together. van der Waals Forces van der Waals forces are a combination of three different kinds of intermolecular forces. Those forces are dispersion forces, dipole-dipole forces, and hydrogen bonding. The ability to determine the polarity of the molecule is the key to determining how many of these forces are acting between molecules in a substance. page of Revised - 2017-2018 - LCA 1 2 KEY: If intermolecular forces are strong, condensation will occur at higher temperatures. In other words, a gas becomes a liquid at a high temperature and it is no longer a gas at that high temperature. If intermolecular forces are weak, condensation will occur at lower temperatures. In other words, a gas becomes a liquid at a lower temperature and it is a gas until it hits this lower temperature.
Transcript
Part 01 - Notes: Intermolecular Forces Honors Chemistry Unit 10 - IMF & Cond States
Learning Objectives: Propose reasoning for the relationship between attractive forces and the boiling point and melting point of sub-stances, among other properties. Develop the link between boiling points, attractive forces and (a) mass and (b) molecular polarity.
Text Reference: Chang and Goldsby (Chemistry: The Essential Concepts, McGraw-Hill, 2014) pp. 403−409.
Intramolecular Versus Intermolecular Forces Two different types of forces the occur between molecules: INTERMOLECULAR and INTRAMOLECULAR
What is an INTRAMOLECULAR FORCE?
Where do these occur? What are these?
What is an INTERMOLECULAR FORCE?
Where do these occur? Why are these significant?
Discuss the intermolecular forces involved in solids, liquids, and gases.
Condensation is the process by which a gas becomes a liquid.
When two molecules in the gaseous state come close to each other, a momentary attraction forms between them. Whether the two particles condense depends upon their kinetic energy (temperature) and the strength of their intermolecular forces. Increasing the pressure can also cause condensation by forcing the particles of a gas close enough to allow the intermolecular forces to hold them together.
van der Waals Forces van der Waals forces are a combination of three different kinds of intermolecular forces. Those forces are dispersion forces, dipole-dipole forces, and hydrogen bonding. The ability to determine the polarity of the molecule is the key to determining how many of these forces are acting between molecules in a substance.
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KEY: If intermolecular forces are strong, condensation will occur at higher temperatures. In other words, a gas becomes a liquid at a high temperature and it is no longer a gas at that high temperature.
If intermolecular forces are weak, condensation will occur at lower temperatures. In other words, a gas becomes a liquid at a lower temperature and it is a gas until it hits this lower temperature.
Part 01 - Notes: Intermolecular Forces Honors Chemistry Unit 10 - IMF & Cond States
Dispersion Forces Dispersion forces are the ONLY forces in nonpolar molecules.
Within all molecules, momentary dipoles arise because electron distributions are constantly changing. When the elec-tron distribution becomes unsymmetrical, it forms momentary dipoles. Dispersion forces occur when the partially nega-tive end of a molecule’s momentary dipole attracts the partially positive end of another molecule’s momentary dipole.
An important fact to note is that the larger the molar mass of a nonpolar compound, the greater the dispersion forces between the molecules of that substance. The greater the intermolecular forces in a substance, the higher the substance’s boiling temperature will be. This is be-cause, the stringer the intermolecular forces, the stronger the attraction of the molecules to one another and this means it will take more energy to separate these molecules. More energy to separate these molecules means that it will take a higher temperature before the molecules separate.
Dispersion forces can also occur between atoms. Dispersion forces are the only forces of attraction between the atoms of a noble gas. Noble gases have very low boiling points.
Dipole-Dipole Forces Polar covalent molecules are subject to both dispersion forces and dipole-dipole forces. Dipole-Dipole forces arise from the electrostatic attraction between the opposite charges of (permanent) dipoles in different molecules. This signifies that these forces occur only in molecules that have molecular dipoles, in other words, molecules that are polar. Ionic and covalent bonds are about 100 times stronger than dipole-dipole forces.
Comparing two compounds of equal molar mass, one composed of nonpolar molecules and the other composed of po-lar molecules, the intermolecular forces are not the same, despite the molar mass. The intermolecular forces would be stronger in the polar compound because both dispersion and dipole-dipole forces are present in the polar compound as opposed to only dispersion forces in the nonpolar compound.
Hydrogen Bonding Hydrogen bonding is an especially strong intermolecular force. It is a special type of dipole-dipole force. Hydrogen bonding occurs between molecules that have N-H, O-H, or F-H attractions. Nitrogen, oxygen, and fluorine are very electronegative and draw electrons to itself more strongly than most other atoms in a molecule. This creates an especial-ly strong molecular dipole. This strong molecular dipole gives rise to an especially strong dipole-dipole force between two molecules. Hydrogen bonding is about one-tenth the strength of covalent bonding.
Note: Hydrogen bonding occurs BETWEEN molecules, NOT WITHIN molecules.
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Part 02 - Notes: The Nature of Liquids & Vapor Pressure Honors Chemistry Unit 10 - IMF & Cond States
Objectives: Describe the nature of a liquid in terms of the attractive forces between the representative particles. Differentiate between evaporation and boiling of a liquid using the kinetic molecular theory. Describe how the degree of organization of particles distinguishes solids form liquids and gases. Identify, define, and explain: vaporization, evaporation, vapor pressure, boiling point, normal boiling point, and melting point.
Text Reference: Chang and Goldsby (Chemistry: The Essential Concepts, McGraw-Hill, 2014) pp. 410−413.
Condensed States Solids and liquids are considered the condensed states of matter because their molecules are more closely arranged than those of substances in the gaseous state.
Initial Question: What makes a liquid different from a solid?
What is the change of state involved in VAPORIZATION?
What is the change of state involved in EVAPORATION?
What is the difference between vaporization and evaporation?
Evaporation is a cooling process. What does this statement mean?
Why do things evaporate more quickly when heated?
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Part 02 - Notes: The Nature of Liquids & Vapor Pressure Honors Chemistry Unit 10 - IMF & Cond States
What do you know about the kinetic energy of the molecules of a sample of a solid or a liquid? Do all the molecules have exactly the same kinetic energy? What does this mean for the system?
Equilibrium: the condition, in any reversible process, where the forward and reverse processes occur at the same rate Example of reversible conditions. . .
Vapor Pressure: the pressure of the vapor over a solid or liquid when the two states are in equilib-rium
Remember absolute temperature is proportional to average kinetic energy. T ∝ KEavg
Q1: How do you increase the vapor pressure of a liquid or solid? Why?
Q2: What is the difference between the vapor (from vapor pressure) and a gas?
Vapor Pressure Recall from above . . .
Equilibrium: the condition, in any reversible process, where the forward and reverse processes occur at the same rate
Vapor Pressure: the pressure of the vapor over a solid or liquid when the two states are in equilibrium
Now, you know that when a substance boils, it changes from a liquid to a gas, and the two states are in equilibrium. Re-lating this back to equilibrium and vapor pressure, what exactly is boiling point.
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Part 02 - Notes: The Nature of Liquids & Vapor Pressure Honors Chemistry Unit 10 - IMF & Cond States
Boiling Point:
Graph and Questions: The vapor pressure of a gas above its own (solid or) liquid depends on temperature. The boiling point or temperature at which bubbles of vapor form within a liquid, depends on both vapor pressure and at-mospheric pressure. The following table and graph shows the vapor pressure of a certain liquid at various temperatures.
Q3: What effect does increasing the temperature have on the vapor pressure? Explain in terms of energy and forces.
Q4: If the atmospheric pressure during this experiment is 96 kPa, what would be the boiling point of the graphed substances? Mark this point on the graph.
Q5: What would happen to the boiling point if the atmospheric pressure were to rise? WHY???
Q6: What would happen to the boiling point if the liquid were tested at a higher altitude? WHY???
Q7: Why do we say that water boils at 100oC? What is 100oC in terms of water?
Temperature(oC) Pressure(kPa)
0 1
10 2
20 4
30 7
40 11
50 16
60 22
70 29
80 37
90 46
100 56
110 67
120 79
130 92
140 106
150 121
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0
20
40
60
80
100
120
140
0 20 40 60 80 100 120 140 160
Vapo
rpressure(k
Pa)
Temperature(oC)
VaporPressurevsTemperature
Part 02 - Notes: The Nature of Liquids & Vapor Pressure Honors Chemistry Unit 10 - IMF & Cond States
Normal boiling point:
There is also a normal freezing point. Normal freezing point:
Volatility Volatile:
Volatile substances have high vapor pressure because they turn easily from a liquid to a gas.
Rubbing alcohol is an example of a highly volatile substance.
What can you say about the intermolecular forces of volatile substance?
Q8: Which is more volatile: olive oil or rubbing alcohol?
Q9: Which is more volatile: acetone (nail polish remover) or motor oil?
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Part 03 - Notes: The Nature of SolidsHonors Chemistry U10 - IMF & Cond States
Objectives: Identify, define, and explain: crystal, unit cell, allotropes, molecular solids, metallic solids, ionic solids, network (covalent) solids, and amorphous solids. Differentiate between the various types of solids using their intermolecular forces and other properties.
Text Reference: Chang and Goldsby (Chemistry: The Essential Concepts, McGraw-Hill, 2014) pp. 419−422.
There are various types of solids: amorphous, molecular, metallic, ionic, and network (covalent). What is a key difference between these types of solids? Thetypesofintermolecularforcesbetweenthemandthearrangement.
Part 03 - Notes: The Nature of SolidsHonors Chemistry U10 - IMF & Cond States
Question: There are three different molecular forms of carbon. What are these three forms? Graphite,diamond,andBuckminsterfullerenearethethreeallotropesofcarbon.
Part 04 - Notes & ALE: Heating Curves & Phase DiagramsHonors Chemistry U10 - IMF & Cond States
Objectives: Analyze and interpret the phase diagram of water or other substance at any given temperature and pressure. Describe the behavior of solids that change directly to the gaseous state and re-condense to solids without passing through the liquid state. Identify, define, and describe: phase diagram, triple point, sublimation, deposition, boiling, evaporation, melting, phase change, heat of fusion, heat of vaporization, heat capacity, heating curve, cooling curve, heat, calorie, and joule.. Classify, by type, the heat changes that occur during melting, freezing, boiling, and condensation. Graph and interpret a heating curve or cooling curve.
Text Reference: Chang and Goldsby (Chemistry: The Essential Concepts, McGraw-Hill, 2014) pp. 423−431.
Heating Curves Q1: Is evaporation endothermic or exothermic? How do you know?
Q2: Because of the attraction between molecules in a liquid, what must happen for a liquid to change to a gas?
Heat supplied at a constant rate to a substance will be absorbed by that substance. Temperature measurements reveal energy changes and changes of state of the substance.
Model I: Heating Curve of Water Let’s examine temperature data recorded and the graph created when a sample of ice was heated from a temperature of -10oC to a temperature of 110oC. Consider this information as you answer the following questions.
Note, there are 5 segments in this graph. It shows the heating of solid H2O (ice) to gaseous H2O (water vapor). What is happening in each segment?
Part 04 - Notes & ALE: Heating Curves & Phase DiagramsHonors Chemistry U10 - IMF & Cond States
Segment 1:
Segment 2:
Segment 3:
Segment 4:
Segment 5:
The above graph gives a qualitative picture of the heating of water from solid to gaseous states. But, in chemistry we need to be quantitative. Let’s define some key terms.
calorie: the amount of heat energy required to raise the temperature of 1 g of water at standard pressure by 1oC
Calorie: 1 Calorie = 1000 calorie = 1 kilocalorie
Joule: unit of thermal energy; 1 J = 1 kg m/s2; amount of energy required to move a 1 N force a distance of 1 m
Conversion:1 calorie = 4.18 J
Specific Heat: the name given to the amount of heat required to raise the temperature of matter
Specific Heat Capacity = Cp: the amount of heat required to raise the temperature of 1 g of a sub-stance by 1oC
The amount of heat needed to raise the temperature of 1 g of a substance by 1oC varies from substance to substance. It also varies with the state of the substance. The specific heat for a substance is different for its solid, liquid, and gaseous states. Also, when the specific heat is given, the temperature must be stated, because the specific heat varies from one temperature to another. page of New - 2016-2017 - LCA2 4
Part 04 - Notes & ALE: Heating Curves & Phase DiagramsHonors Chemistry U10 - IMF & Cond States
Examine the specific heat for water. water(s) = 2.060 J/goC at 0oC water(l) = 4.22 J/goC at 0oC water(l) = 4.184 J/goC at 25oC water(g) = 2.070 J/goC at 100oC water(l) = 4.21 J/goC at 100oC
A substance with a high specific heat can be used as a “heat sponge” and may be used to absorb and store extra heat.
If the temperature of a substance is increased, the substance has absorbed energy. If the temperature of a substance is decreased, the substance has released energy.
The specific heat of water, as listed above, is 4.184 J/goC at 25oC. This specific heat uses joules as the unit of energy. Another unit may also be used: calories. Since a calorie is also a unit of energy, the specific heat of water may also be designated as 1.00 cal/goC at 25oC.
Q3: What does a sloped line in a “heating curve” indicate?
What does a plateau in a “heating curve” indicate?
Q4: Which segment is longer, segment 2 or segment 4? Why do you think this segment is longer?
Q5: What would the graph look like from 10 – 21 minutes if the data was collected in the mountains?
Q6: The graph above is a heating curve. What would the graph of a substance being cooled look like?
Model II: Phase Diagrams Phase diagram: gives conditions of temperature and pressure at which a substance exists as a sol-id, liquid, and gas
The Phase Diagram for water is shown to the right ☞
Part 04 - Notes & ALE: Heating Curves & Phase DiagramsHonors Chemistry U10 - IMF & Cond States
In the phase diagram: • Each region represents a phase of water. • A line that separates two regions shows the conditions at which those two phases are in equilibrium. • The curved line between liquid and vapor also shows how water’s vapor pressure varies with temperature.
Triple Point: the only set of pressure and temperature conditions at which all three states of mat-ter exist in equilibrium with one another. The triple point is a temperature/pressure set.
For water, the triple point is: 0.0098oC and 0.0060 atm (0.61 kPa).
Notice, the normal boiling and melting points indicated on the phase diagram.
Q7: What happens if boiling and melting are carried out at pressures less than 1 atm?
Q8: What happens to the boiling point, if the pressure is higher than 1 atm?
Q9: At what temperature and pressure are the liquid and solid phases of water in dynamic equilibrium?
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Part 05 - Notes: Heat Calculations Honors Chemistry Unit 10 - IMF & Cond States
Objectives: Calculate heat changes that occur during heating, melting, freezing, boiling, condensing, and cooling. Identify, define, and explain: specific heat capacity, heat of vaporization and heat of fusion.
Text Reference: Chang and Goldsby (Chemistry: The Essential Concepts, McGraw-Hill, 2014) pp. 193-198.
Calculating Quantities of Heat specific heat capacity (Cp): the amount of heat required to raise the temperature of 1 g of a substance by 1oC.
On what part of a heating curve would specific heat capacity be significant?
Heat quantities depend on the substance’s specific heat, the quantity being heated, and the temperature change.
The heat change, resulting in a temperature change, may be calculated for the following: a solid being heated to its melt-ing point, a liquid being heated to its boiling point, a gas being superheated, and any of the reverse processes for the sub-stance being cooled.
The formula for calculating heat changes resulting in a temperature change:
__________ = _________________________ with a units of ____________________
__________ = _________________________ with a units of ____________________
__________ = _________________________ with a units of ____________________
__________ = _________________________ with a units of ____________________
Example 1: How much heat is absorbed by 126 g of water if its initial temperature is 47.9oC and its final is 75.6oC?
Q1: Can this formula be used for a sample of water with a temperature change from 56.9oC to 123.9oC? Explain.
The specific heat of a substance describes the energy that must be added to increase the temperature of a substance or the amount of energy that must be removed to lower its temperature. However, when you change the temperature of a substance, it may undergo a change of state, thereby requiring the use of a different specific heat. Also, there are tremendous amounts of energy involved in changes of state.
Heat of fusion:
Heat of vaporization:
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Part 05 - Notes: Heat Calculations Honors Chemistry Unit 10 - IMF & Cond States
So if there is no temperature change and the change occurring is a change of state, you need to calculate the energy us-ing:
Q = m Cfus or Q = m Cvap
Example 2: How much energy, in joules, is required to boil 75.8 g of water at 100oC?
When you solve a heat problem where the substance passes through a phase change or multiple phase changes, you will need to complete multiple steps to arrive at the final answer.
Use the following values for H2O: H2O – solid Cp = 2.060 J/goC Cfus = 334 J/g H2O – liquid Cp = 4.184 J/goC
H2O – gas Cp = 2.070 J/goC Cvap = 2260 J/g
Example 3: You have a sample of ice at –47oC and the sample has a mass of 357.9 g. You need to get the temperature of your sample to +118.6oC. What is the total amount of heat required for this change? (Show each part separately!!!)
In this example, you are starting at a temperature below the freezing point and you need to get to a temperature above the freezing point. Each segment on the change of state graph needs to be completed separately.
You will calculate the five steps: below freezing to freezing, from solid to liquid, from freezing to boiling, from liquid to gas, from boil-ing to above boiling. After each of the steps has been calculated, you will add together each portion to arrive at your final answer.
Below freezing point to freezing point:
At freezing point – from a solid to a liquid:
From freezing point to boiling point:
At boiling point – from a liquid to a gas:
From boiling point to above boiling point:
Total heat involved in the change:
Careful to keep T = Tfinal - Tinitial What does it mean if you get a negative value for T? Why is it important to keep the sign for T?
