9.8: Alternative text for Chapter 9

Figure 9.1.1: Sublimation of Dry Ice (Solid Carbon Dioxide) into Carbon Dioxide Gas

Figure showing stacked blocks of white dry ice (solid carbon dioxide) subliming. The blocks are arranged in a small pile with frosty, opaque surfaces and sharp edges. A visible cloud of white vapor forms around the blocks as the dry ice changes directly from solid to gas. The gas appears to flow downward and pool near the base, illustrating that cold carbon dioxide is denser than air. The background is plain and uncluttered to emphasize the sublimation process.

Figure 9.2.1: Transitions Between Solid, Liquid, and Gaseous States 

The image visually demonstrates how the physical state of a substance—solid, liquid, or gas—is determined by the balance between intermolecular forces (IMFs) and kinetic energy (KE). At the center are three stoppered Erlenmeyer flasks arranged horizontally. Each flask contains the same substance in a different phase. The flask on the left holds a crystalline solid, where particles are tightly packed in an orderly, rigid structure, reflecting strong IMFs and low kinetic energy. The middle flask contains a liquid, with particles still close together but more disordered and mobile. This phase represents a balance between moderate IMFs and increased kinetic energy. The flask on the right contains a gas, in which the particles are far apart and moving rapidly in random directions. In this phase, kinetic energy dominates, overcoming the weaker IMFs entirely.

Above the three flasks, a horizontal arrow points to the right and is labeled “increasing kinetic energy with temperature.” This arrow indicates that as temperature rises, so does the kinetic energy of the molecules, making phase transitions from solid to liquid to gas more likely. At the bottom of the image, another horizontal arrow points to the left and is labeled “increasing intermolecular forces.” This arrow shows that stronger IMFs favor the formation of more ordered, condensed phases such as liquids and solids. Together, the arrangement of flasks and directional arrows illustrates the dynamic balance between kinetic energy and intermolecular forces. A substance’s phase is determined by which of these two opposing factors is more influential under given conditions.

Figure 9.2.2: Condensation Forms When Water Vapor in The Air is Cooled

This figure visually demonstrates condensation, the phase change in which water vapor transitions into liquid water due to cooling. In image (a), a cold beverage glass sits in a warm, humid environment. The outer surface of the glass is covered with numerous visible droplets of liquid water. These droplets form when water vapor in the surrounding air comes into contact with the cold glass surface. The drop in temperature reduces the kinetic energy of the vapor molecules, allowing intermolecular forces—particularly hydrogen bonding—to draw the molecules together into the liquid phase. In image (b), condensation is shown in the form of fog. Here, tiny liquid water droplets appear suspended in the air, creating a misty appearance near the ground or over a surface. This occurs when warm, moist air cools rapidly, often due to contact with cooler air or surfaces, causing water vapor to condense without needing a solid surface. In both scenes, the decrease in temperature allows the attractive forces between H₂O molecules to overcome their motion, pulling them into the more condensed, liquid state and illustrating how condensation can occur both on surfaces and in open air.

Figure 9.2.3: Butane Gas Compression and Liquefaction in a Lighter

If the temperature of a liquid becomes sufficiently low or the pressure on it becomes high enough, the molecules lose enough kinetic energy (KE) that they can no longer overcome the intermolecular forces (IMFs) between them, causing the liquid to solidify into a fixed, orderly structure. The image shows a green disposable lighter containing liquid butane inside its fuel compartment. Normally a gas at standard temperature and pressure, butane is compressed inside the lighter, which forces the molecules closer together. This increased pressure strengthens the intermolecular attractions relative to their kinetic energy, causing the butane gas to condense into liquid form. This example indicates how gases like butane can be liquefied by compression if the temperature remains sufficiently low. A more detailed discussion of these and other phase transitions will be provided later in the chapter.

Figure 9.2.4: Intramolecular Forces Keep a Molecule Intact

The image in Figure 9.2.4 shows two hydrogen chloride (HCl) molecules positioned near each other to demonstrate the difference between intramolecular and intermolecular forces. Each molecule contains a small white cloud representing a hydrogen atom (H) and a larger green cloud representing a chlorine atom (Cl). In the left molecule, a solid straight line connects the hydrogen and chlorine atoms, labeled "intramolecular force (strong)," indicating the covalent bond that holds the atoms together within the molecule. Between the two molecules, a brown dashed line connects the hydrogen atom of one molecule to the chlorine atom of the other. This line is labeled "intermolecular force (weak)" and represents the dipole-dipole attraction between the molecules. The image emphasizes that intramolecular covalent bonds are significantly stronger—requiring about 430 kilojoules per mole to break—while the intermolecular forces between molecules are much weaker, requiring only about 17 kilojoules per mole to overcome.

Figure 9.2.5: Dispersion Forces Result From The Formation of Temporary Dipoles

The image illustrates the formation of dispersion forces between two nonpolar diatomic molecules. Both molecules are shown side by side, each consisting of two atoms with electron clouds surrounding them. In this image, only an uneven distribution of electrons is depicted—there is no initial stage showing even distribution. The electron cloud around each molecule is shown as slightly distorted, with more electron density concentrated on one side. This distortion is indicated by a directional line across each molecule, representing a temporary dipole, with one end labeled δ⁻ (partial negative) and the opposite end labeled δ⁺ (partial positive).

The temporary dipole in one molecule induces a similar dipole in the neighboring molecule, which is also labeled with δ⁻ and δ⁺ to reflect the shift in electron density. Between the two molecules, a brown dashed line connects the oppositely charged ends, indicating the intermolecular attraction known as the dispersion force. This visual representation shows how fluctuations in electron distribution can create temporary dipoles that lead to weak, short-lived electrostatic attractions, even in molecules that are overall nonpolar.

Figure 9.2.6: Boiling Points of Group 14 Tetrahydrides Increase Down The Group

A graph of the boiling points of group 14 tetrahydrides—CH₄, SiH₄, GeH₄, and SnH₄—plotted against the period number of the central group 14 element. The vertical axis represents temperature (boiling point), and the horizontal axis represents the period number. Each compound is represented by a blue sphere placed at its corresponding period on the x-axis, and the spheres are connected by a blue line to illustrate the trend.

The graph displays a clear upward trend, with CH₄ at period 2 showing the lowest boiling point and SnH₄ at period 5 showing the highest. This pattern reflects the increasing molar masses of the compounds: CH₄ (16 g/mol), SiH₄ (32 g/mol), GeH₄ (77 g/mol), and SnH₄ (123 g/mol). As the size and molar mass of the central atom increase down the group, so does the polarizability of the molecules.

This increase in polarizability leads to stronger dispersion forces, which in turn raise the boiling points. Since all four compounds are nonpolar and do not experience dipole-dipole interactions or hydrogen bonding, dispersion forces are the dominant intermolecular forces. The graph confirms the predicted order of boiling points—CH₄ < SiH₄ < GeH₄ < SnH₄—and demonstrates how dispersion forces become stronger in larger, more polarizable molecules.

Figure 9.2.7: The Strength of The Dispersion Forces Increases With Surface Area

The image presents a comparative visualization of the three isomers of C₅H₁₂—neopentane, isopentane, and n-pentane—arranged from left to right in order of increasing boiling point. Each molecule is shown with its mirror image placed directly above it to emphasize structural symmetry and shape. On the far left is neopentane, the most compact and spherical isomer, with its boiling point labeled below as 9.5 °C. Above its mirror image, a label reads "Small contact area → Weakest attraction," indicating that its compact shape limits the extent of intermolecular contact, resulting in the weakest London dispersion forces. In the center is isopentane, with a moderately branched structure and a boiling point of 27 °C. The label above it reads "Less surface area → Less attraction," showing that while it has more surface contact than neopentane, its dispersion forces remain weaker than those of n-pentane. On the far right is n-pentane, with a fully extended, linear structure. Its boiling point is listed as 36 °C, and the label above its mirror image states "Large contact area → Stronger attraction," highlighting that its elongated shape allows for maximum surface contact between molecules, resulting in the strongest dispersion forces among the three isomers. The image conveys that although all three compounds share the same molecular formula (C₅H₁₂), differences in molecular shape directly affect their intermolecular interactions and boiling points. This concept is analogous to VELCRO strips, where a larger contact area produces a stronger connection.

Figure 9.2.8: Geckos’ Toes Contain Large Numbers of Tiny Hairs (setae)

The image visually explains the remarkable adhesive ability of geckos, progressing from a whole-animal view on the left to microscopic structures on the right. On the left, a full gecko is shown climbing a vertical or upside-down surface, illustrating its ability to stick without the use of suction or glue. In the center, a close-up of the gecko's toe reveals hundreds of thousands of tiny hair-like structures called setae. These setae are exaggerated in size for clarity, emphasizing their role in surface contact. On the right, an even closer view shows that each seta branches into hundreds of even tinier, flat, triangular structures called spatulae. These are also enlarged for visualization. Arrows or labels may indicate how the massive number of spatulae increases the gecko's total surface contact, enabling strong adhesion via dispersion forces —weak intermolecular forces that arise from temporary charge distributions.

Figure 9.2.9: Dipole–Dipole Interactions Between HCl Molecules in Different Orientations

Figure illustrating dipole–dipole interactions between two hydrogen chloride (HCl) molecules in two different orientations. On the left, the two HCl molecules are positioned side by side horizontally, each showing a partial positive charge on hydrogen (δ+) and a partial negative charge on chlorine (δ−); a horizontal dotted line connects the δ+ end of one molecule to the δ− end of the other, indicating electrostatic attraction. On the right, the two HCl molecules are positioned vertically next to each other (not stacked), with opposite ends aligned so that the δ+ region of one molecule is adjacent to the δ− region of the other; two vertical dotted lines indicate dipole–dipole attractions between the opposing partial charges.

Figure 9.2.10: Water Molecules Participate in Multiple Hydrogen-Bonding Interactions

Figure 9.2.10 illustrates dipole-dipole interactions using five water (H₂O) molecules as an example. Each molecule is composed of two white spheres labeled “H” for hydrogen and one red sphere labeled “O” for oxygen, reflecting the standard color coding used to represent water. The partial charges are not explicitly shown, but the polarity is understood: the oxygen atom is more electronegative and thus carries a partial negative character, while the hydrogen atoms carry a partial positive character.

In the image, the five water molecules are arranged to demonstrate how dipole-dipole interactions occur between them. The molecules are oriented so that the hydrogen atoms of one water molecule (white H spheres) are positioned near the oxygen atoms (red O spheres) of neighboring molecules. Between these interacting atoms, red dashed lines are drawn to indicate the intermolecular forces—specifically, dipole-dipole interactions, which in the case of water are also classified as hydrogen bonds.

Figure 9.2.11: Boiling Points of Group 15, 16, and 17 Hydrides Across Periods 3–5

The image is a line graph showing the boiling points of binary hydrides from groups 15, 16, and 17 plotted against the period number on the x-axis, spanning Period 2 to Period 5. The y-axis represents boiling points in degrees Celsius. Each group’s hydrides are represented by distinct colored lines for easy identification: green for group 15 (NH₃, PH₃, AsH₃, SbH₃), blue for group 16 (H₂O, H₂S, H₂Se, H₂Te), and red for group 17 (HF, HCl, HBr, HI).

Moving from Period 2 to Period 5, the boiling points for the heavier three hydrides in each group rise steadily. This trend reflects increasing molecular size and stronger dispersion forces as the dominant factor, despite a slight decrease in molecular polarity. For example, the green line representing group 15 shows PH₃, AsH₃, and SbH₃ boiling points increasing, while the blue and red lines for groups 16 and 17 show similar steady upward trends for their heavier hydrides.

However, the Period 2 hydrides—NH₃ (green), H₂O (blue), and HF (red)—are marked with question marks (?) at their data points to indicate that their boiling points deviate significantly from the expected trend. These points lie well above the smooth curves formed by the heavier hydrides, signaling an anomaly. This elevated boiling point is caused by strong hydrogen bonding, which is much stronger than the dispersion and dipole-dipole forces that dominate in the heavier hydrides.

The question marks highlight the unexpected nature of these Period 2 hydrides’ boiling points, emphasizing the key role of hydrogen bonding due to the highly electronegative atoms (N, O, and F) bonded to hydrogen. This graph effectively contrasts the smooth increase in boiling points driven by dispersion forces in heavier hydrides with the pronounced hydrogen-bonding effect seen in the lighter, Period 2 compounds.

Figure 9.2.12: In Comparison to Periods 3−5, The Binary Hydrides of Period 2  

The image is a line graph that plots the boiling points of binary hydrides from groups 15, 16, and 17 against their corresponding periods on the x-axis, from Period 2 to Period 5. The y-axis represents boiling point in degrees Celsius. Each group of hydrides is represented by a distinct colored curve: green for group 15 (NH₃, PH₃, AsH₃, SbH₃), blue for group 16 (H₂O, H₂S, H₂Se, H₂Te), and red for group 17 (HF, HCl, HBr, HI).

For all three groups, the curves display a general upward trend from Period 3 to Period 5, reflecting the increasing boiling points of heavier hydrides. This smooth rise is attributed to stronger dispersion forces that come with increasing molecular size and polarizability, even as molecular polarity slightly decreases. For example, in the green curve (group 15), the boiling points rise steadily from PH₃ to AsH₃ to SbH₃. Similarly, the blue and red curves for groups 16 and 17 show smooth increases from H₂S to H₂Te and from HCl to HI, respectively.

However, at Period 2, each curve makes a sharp upward bend, forming a noticeable spike for NH₃ (green), H₂O (blue), and HF (red). These Period 2 hydrides have significantly higher boiling points than expected based on the trend from heavier analogs. Rather than continuing the smooth progression seen in the rest of each curve, the lines bend sharply to accommodate these anomalously high values, creating a distinct “kink” or spike in each group’s trend line.

These sharp upward bends visually emphasize the exceptional impact of hydrogen bonding in NH₃, H₂O, and HF. Because nitrogen, oxygen, and fluorine are highly electronegative and small in size, they form strong hydrogen bonds with neighboring molecules—intermolecular forces much stronger than the dipole-dipole or dispersion forces that dominate in the heavier hydrides. The graph powerfully illustrates how hydrogen bonding causes a dramatic increase in boiling point, breaking the otherwise smooth trend set by molecular weight and dispersion forces alone.

Figure 9.2.13: Two Separate DNA Molecules Form a Double-Stranded Helix 

The image shows a detailed representation of the DNA double helix, emphasizing its structural components and molecular organization. The two strands of DNA spiral around each other, forming the well-known right-handed helical shape. Each strand features a sugar-phosphate backbone, visually represented by a light blue ribbon that curves smoothly along the outer edges of the helix. This ribbon traces the backbone of each strand, composed of alternating sugar (deoxyribose) and phosphate units.

Extending inward from the backbone are the nitrogenous bases, which are color-coded for clarity: adenine is red, thymine is orange, guanine is blue, and cytosine is green. These bases pair specifically—adenine with thymine, and guanine with cytosine—through hydrogen bonds, which are shown as horizontal lines connecting the bases across the two strands. This base-pairing forms the "rungs" of the DNA "ladder."

The image also includes directional labels: one end of each strand is marked 5′ (five-prime) and the other 3′ (three-prime), indicating the antiparallel orientation of the strands. One strand runs from 5′ to 3′, while the opposite strand runs from 3′ to 5′. This directionality is critical for DNA replication and transcription. Together, the light blue ribbon backbone, colored bases, and labeled ends give a clear and informative visualization of DNA’s double-stranded, complementary, and antiparallel structure.

Figure 9.2.14: The Geometries of The Base Molecules Result in Maximum Hydrogen Bonding 

The image shows a vertical molecular diagram of complementary base pairing in DNA, with clear representation of the structural orientation of the two antiparallel strands. At the top, the base pair between thymine (T) and adenine (A) is illustrated. Thymine, a single-ring pyrimidine, is on the left, and adenine, a double-ring purine, is on the right. These bases are connected by two dashed lines, representing the two hydrogen bonds that stabilize the A–T pairing.

At the bottom, the base pair between guanine (G) and cytosine (C) is shown, with guanine on the left and cytosine on the right. These are held together by three dashed hydrogen bonds, indicating a stronger interaction than the A–T pair. All hydrogen bonds are shown in the center of the image, highlighting the specific and directional pairing between purines and pyrimidines.

Flanking the base pairs on both sides are vertical straight-line structures representing the sugar-phosphate backbone of the DNA strands. These backbones are not depicted as ribbons but as simple linear bonds, emphasizing the chemical framework rather than the spatial twist of the helix. The left strand is labeled 3′ to 5′, with the direction running bottom to top, while the right strand is labeled 5′ to 3′, also oriented bottom to top. This reflects the antiparallel arrangement of DNA strands, crucial for DNA replication and base pairing.

Figure 9.3.1: The Kinetic Theory of Gases

The image shows a transparent cubic box representing a closed container filled with gas particles. Inside the box are many small spheres, each symbolizing an individual gas particle. These spheres are spaced widely apart, highlighting that gas particles occupy a very small fraction of the container’s volume and are separated by large distances.

Each particle is shown with arrows pointing in different directions, indicating their constant, random motion within the container. The arrows vary in length, symbolizing the different speeds of the particles as they move independently of one another. Several particles are depicted colliding with the container walls or with each other, demonstrating the concept of elastic collisions where kinetic energy is conserved and there is no net energy loss.

The image intentionally shows no clustering or attraction between particles, reflecting the principle that there are no intermolecular forces (neither attraction nor repulsion) acting between gas particles. The container itself is static, with no changes in size or shape, indicating a rigid boundary confining the gas.

Overall, the image visualizes the fundamental postulates of the Kinetic Molecular Theory: gases consist of tiny particles in continuous, random motion; these particles are separated by large distances relative to their size; they collide elastically with each other and the container walls; and they do not exert attractive or repulsive forces on each other. This model explains the physical behavior of gases under many conditions.

Figure 9.4.1: Atmospheric Pressure and How Atmospheric Mass is Distributed With Altitude

This image visually represents the concept of atmospheric pressure and how atmospheric mass is distributed with altitude. On the left side of the image is a tall vertical column of air extending from sea level up to 40 kilometers, representing a 1-square-meter column of the atmosphere. This column symbolizes the entire weight of the air pressing down on a unit area at Earth’s surface. At the base of the column is a label indicating a total mass of 10,000 kilograms, which corresponds to the weight of the air above one square meter and is responsible for exerting a pressure of 1 atmosphere (1 atm), equivalent to 101 kilopascals, 760 mmHg, or 1.01 × 10⁵ newtons per square meter. A bold downward arrow labeled “1 atm” points from the top of the column to the surface, emphasizing that atmospheric pressure results from the weight of all the air above.

To the right of the column, a person is shown standing at sea level, illustrating that we live at the bottom of this “ocean of air.” Near the person are silhouettes of two elephants, used as a visual comparison to help convey the magnitude of atmospheric pressure—the weight of the air column is equivalent to having two elephants standing on each square meter of Earth’s surface. Above the person is a graph displaying the relationship between cumulative atmospheric mass and altitude. The vertical axis of the graph shows altitude in kilometers, ranging from 0 to 40 km, while the horizontal axis shows atmospheric mass as a percentage, increasing from 0% to 100%.

A yellow curve runs diagonally across the graph, showing an exponential relationship where atmospheric mass increases rapidly as altitude decreases. The curve begins at 0% atmospheric mass at 40 km and rises smoothly to 100% at sea level. Key data points are labeled along the curve: 50% of the atmosphere’s mass lies below 5.5 km, 90% lies below 16.5 km, and 99% lies below 30 km. These values highlight how the vast majority of atmospheric mass is concentrated in the lowest part of the atmosphere. The graph also labels the troposphere, which extends from 0 to 13 km, and the stratosphere, which ranges from 13 to 44 km, indicating how the atmosphere is layered. While the stratosphere continues beyond the top of the graph, the figure focuses on the first 40 km, where nearly all atmospheric mass is located.

Together, the visual components of the figure—the air column, the person and elephants, the pressure arrow, and the graph—illustrate that atmospheric pressure is the direct result of the mass of air above, and that most of this mass is located within the first few kilometers of the atmosphere. The steadily rising yellow curve visually reinforces that the accumulation of atmospheric mass follows an exponential trend with decreasing altitude, reaching nearly 100% at around 40 km above Earth’s surface.

Figure 9.4.2: A Mercury Barometer

A diagram of a mercury barometer shows a long glass tube closed at the top, filled with mercury, and inverted into a dish of mercury. Some mercury stays in the tube, forming a tall column, while a vacuum exists above it. Arrows indicate atmospheric pressure pushing down on the mercury in the dish and gravity pulling the mercury column downward, illustrating how the column height balances atmospheric pressure.

Figure 9.4.3: Closed-End and Open-End Manometers: Measuring Gas Pressure

Figure showing two types of mercury manometers used to measure gas pressure. In part (a), a closed-end manometer is shown with a U-shaped tube partially filled with mercury; one arm is connected to a gas bulb and the other arm is sealed, creating a vacuum above the mercury in the reference arm (P ≈ 0). The difference in mercury heights between the two arms directly represents the pressure of the gas in the bulb. In part (b), an open-end manometer is shown with one arm connected to a gas bulb and the other arm open to the atmosphere (P ≈ 1 atm). The mercury levels differ depending on the relative pressures: if the gas pressure is greater than atmospheric pressure, mercury is pushed down in the gas-connected arm and rises in the open arm; if the gas pressure is lower than atmospheric pressure, mercury is higher in the gas-connected arm. The height difference between the columns indicates the pressure difference between the gas and atmospheric pressure.

Figure 9.5.1: Demonstration of Boyle's Law

Figure illustrates Boyle’s law with three parts. On the left, a piston-and-cylinder contains very small spots of gas at large volume and low pressure. In the middle, the piston is compressed, reducing the volume by half and increasing the pressure of the gas. On the far right, a graph plots volume versus pressure, with volume on the y-axis and pressure on the x-axis, showing the inverse relationship between pressure and volume for the gas at constant temperature.

Figure 9.5.2: Demonstration of Charles's Law

Figure illustrates Charles’s law. Two piston-and-cylinder setups are shown, each with the same mass weight on top of the piston to maintain constant pressure, a temperature gauge attached to the cylinder, and a burner beneath to heat the gas. On the left, the gas has a lower temperature and smaller volume, and on the right, after heating, the gas expands to a larger volume, demonstrating that the volume of a gas increases with temperature at constant pressure.

Figure 9.5.3: Gas Volume versus Absolute Temperature Relationship

Image shows a graph of gas volume versus absolute temperature using data from Table 13.5.1. Volume (in mL) is on the y-axis, extending from 0 to 400 mL, and temperature (in K) is on the x-axis, extending from 0 to 1000 K. The straight-line graph demonstrates a direct relationship between volume and temperature, illustrating Charles’s law, with the constant K remaining at 0.4.

Figure 9.5.4: Demonstrates Avogadro’s Hypothesis

Figure demonstrates Avogadro’s hypothesis. Four cubic containers are shown from left to right holding helium (He), nitrogen (N₂), ammonia (NH₃), and methane (CH₄). Each container contains one spot representing a gas particle. The gases are at the same volume, pressure, and temperature (0 K). Each gas contains the same number of particles (1 mole). At the bottom of each container, the volume (V), pressure (P), temperature (T), number of moles, and mass of the gas are indicated, showing that while all other properties are equal, the mass differs due to differing molar masses.

Figure 9.5.5: Carbonated Beverage: Dissolved CO₂ Bubbles

The figure shows a glass filled with a carbonated beverage. Small CO₂ bubbles can be seen rising through the liquid, illustrating the gas dissolved in the beverage.

Figure 9.5.6: Scuba Diver Underwater at Depth

A scuba diver is shown underwater positioned below the sea surface. The diver represents being at depth in the ocean, emphasizing increased surrounding water pressure compared to the surface.

Figure 9.5.7: Combined Atmospheric and Hydrostatic Pressure with Depth in Water