Name: water and co2 phase behavior
Uploaded: 2020-09-24T15:06:49.000Z
Duration: 2 min 39 s
Channel: JoVE

K12 Standards

north-carolina-k-12-science-standard-course-of-study

NC K-12 Science Standard Course of Study

9-12

Chemistry

Energy

PS.Chm.7.1

Phase Diagrams

Carbon dioxide can exist as a solid, liquid, or gas under specific temperatures and pressures. This dependence is mapped into a phase diagram, which includes three general features: regions, lines, and points. 
Regions represent the temperature and pressure conditions for a single phase. At standard pressure, carbon dioxide can be solid or gaseous, depending on the temperature.
At pressures below 5.1 atm, if the temperature of solid carbon dioxide is raised, it will directly transition to the gaseous phase without passing through the liquid form.
A transition through all three phases &#8212; solid to liquid and liquid to gas &#8212; will occur at pressures above 5.1 atm. 
The lines or curves that separate the regions denote the temperatures and pressures at which the phases on either side of the curve are in equilibrium. 
For example, the point at 1 atm and &#8722;78.5 &#176;C, lies on the curve separating the solid and vapor phases, so a solid&#8211;vapor equilibrium exists under those conditions. Accordingly, this is called the sublimation curve. 
Similarly, the liquid-gas equilibrium exists on the vaporization curve, and the solid-liquid equilibrium exists on the fusion curve. These curves are known more genera

phase diagrams

A phase diagram combines plots of pressure versus temperature for the liquid-gas, solid-liquid, and solid-gas phase-transition equilibria of a substance. These diagrams indicate the physical states that exist under specific conditions of pressure and temperature and also provide the pressure dependence of the phase-transition temperatures (melting points, sublimation points, boiling points). Regions or areas labeled solid, liquid, and gas represent single phases, while lines or curves represent two phases coexisting in equilibrium (or phase change points). The triple point indicates conditions of pressure and temperature at which all three phases coexist. In contrast, a critical point indicates the temperature and pressure above which a single phase&#8212;whose physical properties are intermediate between the gaseous and liquid states&#8212;exists.
<img alt="Phase diagram showing states of matter; solid, liquid, gas, with transitions: melting, freezing." src="/files/ftp_upload/11352/11352_Image_1.png" style="text-align: center; height: 320px;" data-altupdatedat="1754904545000" data-alt="Phase diagram showing solid, liquid, gas states; transitions include melting, sublimation." />
Figure 1. A typical phase diagram.
A phase diagram identifies the physical state of a substance under specified conditions of pressure and temperature. To illustrate the utility of these plots, consider the phase diagram of water, shown below. 
<img alt="Phase diagram of water showing phases: solid, liquid, gas; includes triple and critical point." src="/files/ftp_upload/11352/11352_Image_2.png" style="text-align: center; height: 320px;" data-altupdatedat="1754904545000" data-alt="Phase diagram of water; solid, liquid, gas state regions; triple and critical points marked." />
Figure 2. Phase diagram of water.
A pressure of 50 kPa and a temperature of &#8722;10 &#176;C corresponds to the region of the diagram labeled &#8220;ice.&#8221; Under these conditions, water exists only as a solid. A pressure of 50 kPa and a temperature of 50 &#176;C corresponds to the region where water exists only as a liquid. At 25 kPa and 200 &#176;C, water exists only in the gaseous state. Curve BC is the liquid-vapor curve separating the liquid and gaseous regions of the phase diagram and provides the boiling point for water at any pressure. For example, at 1 atm, the boiling point is 100 &#176;C. Notice that the liquid-vapor curve terminates at a temperature of 374 &#176;C and a pressure of 218 atm, indicating that water cannot exist as a liquid above this temperature, regardless of the pressure. The physical properties of water under these conditions are intermediate between those of its liquid and gaseous phases. This unique state of matter is called a supercritical fluid. The solid-vapor curve labeled AB, indicates the temperatures and pressures at which ice and water vapor are in equilibrium. These temperature-pressure data pairs correspond to the sublimation, or deposition, points for water.
The solid-liquid curve labeled BD shows the temperatures and pressures at which ice and liquid water are in equilibrium, representing the melting/freezing points for water. Note that this curve exhibits a slight negative slope, indicating that the melting point for water decreases slightly as pressure increases. Water is an unusual substance in this regard, as most substances exhibit an increase in melting point with increasing pressure. The point of intersection of all three curves&#8212;labeled B&#8212;is the triple point of water, where all three phases coexist in equilibrium. At pressures lower than the triple point, water cannot exist as a liquid, regardless of the temperature.
Consider the phase diagram for carbon dioxide as another example.
<img alt="Phase diagram showing states: Solid, Liquid, Gas, and Supercritical Fluid (SCF) regions." src="/files/ftp_upload/11352/11352_Image_3.png" style="text-align: center; height: 350px;" data-altupdatedat="1754904545000" data-alt="Phase diagram displaying states: solid, liquid, gas, and supercritical fluid (SCF) regions." />
Figure 3. Phase diagram of carbon dioxide.
The solid-liquid curve exhibits a positive slope, indicating that the melting point for CO2 increases with pressure as it does for most substances. Notice that the triple point is well above 1 atm, indicating that carbon dioxide cannot exist as a liquid under ambient pressure conditions. Instead, cooling gaseous carbon dioxide at 1 atm results in its deposition into the solid-state. Likewise, solid carbon dioxide does not melt at 1 atm pressure but instead sublimes to yield gaseous CO2. Finally, the critical point for carbon dioxide is observed at a relatively modest temperature and pressure in comparison to water.
This text is adapted from <a href="https://openstax.org/books/chemistry-2e/pages/10-4-phase-diagrams">Openstax, Chemistry 2e, Section 10.4: Phase Diagrams</a>.

A phase diagram combines plots of pressure versus temperature for the liquid-gas, solid-liquid, and solid-gas phase-transition equilibria of a substance. ...

Education

JoVE Core

Liquids, Solids, and Intermolecular Forces

KEY TERMS AND CONCEPTS

water and co2 phase behavior

Water and carbon dioxide phase behavior is shown with phase diagrams that map pressure against temperature. These plots show where a substance exists as a solid, liquid, or gas. They also show how melting points, sublimation points, and boiling points change with pressure.

In a phase diagram, the labeled regions for solid, liquid, and gas represent single phases. The lines or curves between them show two phases in equilibrium. The triple point is the one condition of pressure and temperature where all three phases coexist. The critical point is the temperature and pressure above which a single phase exists with properties between those of a liquid and a gas.

Water provides a useful example of how to read these diagrams. At 50 kPa and -10 °C, water lies in the ice region and exists only as a solid. At 50 kPa and 50 °C, it is in the liquid region. At 25 kPa and 200 °C, it is in the gaseous region.

The water liquid-vapor curve, BC, separates liquid and gas. It gives the boiling point at each pressure, including 100 °C at 1 atm. This curve ends at 374 °C and 218 atm. Above that point, water cannot remain a liquid, no matter the pressure. Instead, it becomes a supercritical fluid, which has properties between a liquid and a gas.

Water also shows a solid-vapor curve, AB, where ice and water vapor are in equilibrium. These pressure-temperature pairs are the sublimation and deposition points. The solid-liquid curve, BD, marks the melting and freezing points. For water, this curve has a slight negative slope, so the melting point decreases a little as pressure increases. Most substances show the opposite trend.

The triple point of water is the intersection of all three curves at B. At pressures below this point, water cannot exist as a liquid, even if the temperature changes. Carbon dioxide gives a different pattern. Its solid-liquid curve has a positive slope, so its melting point increases with pressure, like most substances.

The triple point of carbon dioxide lies well above 1 atm. That means carbon dioxide cannot exist as a liquid at normal atmospheric pressure. When gaseous carbon dioxide is cooled at 1 atm, it deposits directly into a solid. Solid carbon dioxide also does not melt at 1 atm. It sublimes to gas instead. Carbon dioxide also has a critical point, and it occurs at a relatively modest temperature and pressure compared with water.

Water and CO2 Phase Behavior

molecular comparison of gases, liquids, and solids

molecular-comparison-of-gases-liquids-and-solids

How Intermolecular Forces Shape States

Intermolecular forces and particle motion shape the states of matter: solids, liquids, and gases. In a solid, particles are tightly packed and often arranged in a regular pattern. A solid has a fixed shape, and its particles vibrate in place rather than flowing past one another. In a liquid, particles stay close together but do not have a regular arrangement. Liquids have no fixed shape, yet their particles move past each other while remaining in contact. That is why liquids can flow and take...

intermolecular vs intramolecular forces

How Forces Shape Liquids and Solids

Intermolecular forces are electrostatic attractions between molecules. Their strength depends on how far apart the molecules are. These forces matter most in solids and liquids, where particles are close together. In gases, they become important mainly under high pressure because the molecules are then closer together. Intermolecular forces help control key physical properties. They affect melting point, boiling point, density, and the enthalpies of fusion and vaporization. When a liquid is...

intermolecular forces

How Molecules Stick Together

Molecules and ions stick together through intermolecular forces. These are electrostatic attractions and repulsions caused by permanent, partial, or temporary charges. They act between ions, polar molecules, nonpolar molecules, and neutral atoms, and they vary in strength. Some intermolecular forces act between neutral particles and are called van der Waals forces. These include dipole-dipole forces, hydrogen bonding, and dispersion forces. Another important interaction is ion-dipole force,...

comparing intermolecular forces: melting point, boiling point, and miscibility

comparing-intermolecular-forces-melting-point-boiling-point

How Molecule Shape Affects Intermolecular Forces

Intermolecular forces are the attractions between molecules, and they help explain melting point, boiling point, and miscibility. Their strength depends on molar mass, molecular shape, and polarity. These forces shape the bulk behavior of substances that students see in chemistry labs. Dispersion forces are temporary attractions that occur in every substance, whether it is polar or nonpolar. They form when electrons are spread unevenly for a moment, creating a temporary dipole. Stronger...

surface tension, capillary action, and viscosity

surface-tension-capillary-action-and-viscosity

How Liquids Behave on Surfaces

Liquids behave in different ways on surfaces because of surface tension, adhesive forces, capillary action, and viscosity. These properties help explain why some liquids bead up, some spread out, and some flow more easily than others. Surface tension comes from cohesive forces, which are the attractions between identical molecules. Molecules inside a liquid are pulled equally in all directions, but molecules at the surface are pulled by fewer neighbors. That imbalance makes the liquid contract...

phase transitions

Melting, Boiling, and State Changes

Melting, boiling, and other state changes depend on how particles are arranged in a substance. In solids, particles pack closely in an ordered pattern. They vibrate in fixed positions but do not move past nearby particles. In liquids, particles are still close together, but they are randomly arranged and can move past one another. Solids and liquids are called condensed phases because their particles remain close and their intermolecular forces are relatively strong. Gases behave differently...

phase transitions: vaporization and condensation

phase-transitions-vaporization-and-condensation

Heat and Mass Transfer in Phase Changes

Heat and mass transfer drive phase changes in liquids. When a liquid is heated, its molecules move faster. Some surface molecules can then overcome intermolecular forces and escape into the gas phase. This process is called vaporization, and it happens at the surface of the liquid. Vaporization can also happen in a closed container. In that case, gas molecules cannot leave the container. The gas molecules move randomly and may collide with the surface of the condensed phase. Some of those...

vapor pressure

Vapor Pressure and Boiling Point

Vapor pressure and boiling point describe how a liquid behaves in a closed container and when it starts to boil. As a liquid vaporizes, some gas molecules stay above the liquid and can collide with the surface again. When those collisions turn gas back into liquid, the process is called condensation. If vaporization and condensation happen at the same rate, the liquid and vapor amounts stay constant. The vapor is then in equilibrium with the liquid. This is a dynamic equilibrium, which means...

clausius-clapeyron equation

Predicting Vapor Pressure with Temperature

Predicting vapor pressure with temperature is a key use of the Clausius-Clapeyron equation. It describes how the equilibrium between a liquid and its vapor changes as temperature changes. When temperature rises, the vapor pressure of the liquid also rises. The equation gives a quantitative link between vapor pressure, P, and temperature, T. It shows the rate at which vapor pressure increases for each unit increase in temperature. In this relationship, ΔHvap is the enthalpy of vaporization, R...

phase transitions: melting and freezing

Melting Point and Freezing Equilibrium

Melting point and freezing equilibrium describe how a crystalline solid changes between solid and liquid states. As a solid is heated, the average energy of its atoms, molecules, or ions increases, and the solid becomes hotter. At a certain point, the added energy becomes large enough to partly overcome the forces that hold the particles in fixed positions. The solid then starts melting, or changing into a liquid. During melting, the temperature stops rising even though heat is still being...

phase transitions: sublimation and deposition

phase-transitions-sublimation-and-deposition

Sublimation, Frost, and Enthalpy Changes

Sublimation is a phase change in which a solid turns directly into a gas. It skips the liquid state. Dry ice is a clear example at room temperature and standard pressure, where solid carbon dioxide seems to disappear as it sublimes without forming a liquid first. Some other solids also sublime under the right conditions. Snow and ice can sublimate below water’s melting point, and wind or the lower air pressure at high altitudes can speed up that slow process. Solid iodine does the same when...

heating and cooling curves

Phase Changes on Heating and Cooling Curves

Phase changes on heating and cooling curves show how a substance responds when it absorbs or releases heat while isolated from its surroundings. The graph tracks changes in temperature and state as the substance is heated or cooled. When a solid is heated, its temperature rises first. The amount of heat absorbed depends on the solid’s heat capacity, which is written as q = mcΔT. In this equation, m is the mass and c is the specific heat. This relationship applies when matter changes...

Water and carbon dioxide phase behavior is shown with phase diagrams that map pressure against temperature. These plots show where a substance exists as a solid, liquid, or gas. They also show how melting points, sublimation points, and boiling points change with pressure. In a phase diagram, the labeled regions for solid, liquid, and gas represent single phases. The lines or curves between them show two phases in equilibrium. The triple point is the one condition of pressure and temperature...

structures of solids

Crystalline and Amorphous Solid Patterns

Crystalline and amorphous solids differ in how their particles are arranged. In crystalline solids, atoms, ions, or molecules follow a definite repeating pattern. Metals and ionic compounds typically form these ordered solids. Because the same particles are held in place by the same forces or energy, crystalline solids have a precise melting temperature. Amorphous solids, also called non-crystalline solids or glasses, do not have an ordered internal structure. Their particles are arranged...

molecular and ionic 
solids

Molecular and Ionic Crystal Properties

Crystalline solids are grouped into four types: molecular, ionic, metallic, and covalent network solids. This page focuses on molecular and ionic crystals and how their particle type and attractions shape their properties. Molecular crystalline solids are made of neutral molecules. Common examples include ice, sucrose, and iodine. The molecules are held together by weak intermolecular forces, including London dispersion forces, dipole-dipole interactions, and hydrogen bonds. These forces...

ionic crystal structures

Ionic Crystal Packing and Ion Holes

Ionic crystal packing depends on ion size, charge, and the ratio of cations to anions. In these solids, the ions usually behave like charged spheres. Stable structures form when each ion is surrounded by as many oppositely charged ions as possible and the ions remain in contact. In simple ionic crystals, the larger anions usually form a close-packed array. The smaller cations fit into the spaces between them. Two common spaces are tetrahedral holes, which are surrounded by four anions, and...

metallic solids

Metallic Bonding and Crystal Packing

Metallic solids are made of metal atoms arranged in crystals. Copper, aluminum, and iron are common examples. These solids are often described as metal nuclei spread through a sea of delocalized electrons. The atoms are held together by metallic bonding, and that bonding gives metals many of their useful bulk properties. Metallic solids are shiny, conduct heat and electricity well, and can be shaped without breaking. This shaping ability is called malleability, which means a metal can deform...

band theory

Band Gaps in Conductors and Insulators

Band gaps in solids help explain why some materials conduct electricity and others do not. When many atoms come together, their atomic orbitals combine into molecular orbitals with distinct energies. In a solid, so many orbitals are involved that these energies merge into continuous regions called bands. The energy difference between bands is called the band gap. Electrons in a solid must move across orbitals of different energies to travel through the material. In conductors, the valence...

network covalent solids

Why Network Covalent Solids Are So Hard

Network covalent solids are made of atoms joined in a three-dimensional network of covalent bonds. They appear in the crystal structures of nonmetals such as diamond, graphite, and silicon. Some covalent compounds also form this kind of structure, including silicon dioxide, which is found in sand, and silicon carbide, the abrasive used on sandpaper. Many minerals also contain networks of covalent bonds. These solids are hard to melt or break because their covalent bonds must be broken first.

x-ray crystallography

X-ray Diffraction in Crystal Structure

X-ray diffraction can reveal the unit cell size and atom arrangement in a crystal. This method is called X-ray crystallography. It uses diffraction, which is the bending of an electromagnetic wave when it meets a barrier about the same size as its wavelength. X-rays work well for this method because their wavelengths are about the same as the spacing between neighboring atoms in crystals. They are measured on the scale of a few angstroms. When monochromatic X-rays, or X-rays with one...

lattice centering and coordination number

lattice-centering-and-coordination-number

Crystal Packing in Cubic Unit Cells

Crystal packing in cubic unit cells shows how atoms are arranged in a repeating solid structure. A crystal is best described by its unit cell, which is the smallest repeating part of the lattice. Lattice points mark where atoms or ions are located. The full crystal is built when this unit cell repeats in three dimensions. The cubic lattice has three common unit cells. A simple cubic structure, also called a primitive cubic structure, can be pictured as identical spheres stacked in layers with...

key terms and concepts

Explore how phase diagrams show solid, liquid, and gas regions, plus triple and critical points, using water and carbon dioxide as examples.

Water and CO2 Phase Behavior | JoVE K12

chemistry

JoVE Core Chemistry brings general chemistry concepts to life using concise and easy-to-understand animated video lessons. Plus, in the scientist-in-action videos, we demonstrate classical and original research experiments performed in today's laboratories worldwide.

Physical Science

Matter and its Interactions

PS.PSc.1.2

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Ohio’s Learning Standards and Model Curriculum

Study Of Matter

Classification of matter

Energy And Waves

Thermal energy

Structure And Properties Of Matter

Intermolecular forces of attraction

Virginia  Standards of Learning

HIgh School

CH.6

CH.6.a

CH.7

CH.7.b

PS.3

PS.3.b

Texas Essential Knowledge and Skills

High School

CHEM.11

CHEM.11.C

Alabama Course of Study Science

Matter and Its Interactions

9-12.C.1.5.d

9-12.PS.3.5.b

Florida Department of Education

SC.912.P.8

SC.912.P.8.2

Georgia Standards of Excellence

SPS5.

Energy and Phase Transitions

SPS7.

Phase Changes and Energy

Classifying Matter by State

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Classifying Matter by Composition

Matter: Pure Substances and Mixtures
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physical-and-chemical-properties-of-matter

Physical and Chemical Properties of Matter

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Physical Properties of Matter

A physical property is ...

Molecular Comparison of Gases, Liquids, and Solids

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Intermolecular vs Intramolecular Forces

Intermolecular forces (IMF) are electrostatic attractions arising from charge-charge interactions between molecules. The strength of the intermolecular ...

Intermolecular Forces

Atoms and molecules interact through bonds (or forces): intramolecular and intermolecular. The forces are electrostatic as they arise from interactions ...

Comparing Intermolecular Forces: Melting Point, Boiling Point, and Miscibility

Intermolecular forces are attractive forces that exist between molecules. They dictate several bulk properties, such as melting points, boiling points, ...

Surface Tension, Capillary Action, and Viscosity

Surface Tension
The various IMFs between identical molecules of a substance are examples of cohesive forces. The molecules within a liquid are surrounded ...

States of Matter and Phase Changes

The internal energy of a substance&#8212;the total kinetic energy of all its molecules and the potential energy of their associated forces&#8212;depends ...

Phase Transitions

Whether solid, liquid, or gas, a substance&#39;s state depends on the order and arrangement of its particles (atoms, molecules, or ions). Particles in the ...

Phase Transitions: Vaporization and Condensation

The physical form of a substance changes on changing its temperature. For example, raising the temperature of a liquid causes the liquid to vaporize ...

Vapor Pressure

When a liquid vaporizes in a closed container, gas molecules cannot escape. As these gas phase molecules move randomly about, they will occasionally ...

Clausius-Clapeyron Equation

The equilibrium between a liquid and its vapor depends on the temperature of the system; a rise in temperature causes a corresponding rise in the vapor ...

Phase Transitions: Melting and Freezing

Heating a crystalline solid increases the average energy of its atoms, molecules, or ions, and the solid gets hotter. At some point, the added energy ...

Phase Transitions: Sublimation and Deposition

Some solids can transition directly into the gaseous state, bypassing the liquid state, via a process known as sublimation. At room temperature and ...

Heating and Cooling Curves

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Use models to explain how changes in energy affect the arrangement and movement of the particles in solids, liquids, and gases, as well as the relative strengths of their intermolecular forces.

Energy Basics

Chemical reactions, such as those that occur when you light a match, involve changes in energy as well as matter.

Chemical changes and their accompanying ...

First Law of Thermodynamics

Energy Conservation

Energy can be converted from one form into another, but all of the energy present before a change occurs always exists in some form ...

Internal Energy

The total of all possible kinds of energy present in a substance is called the internal energy (U), sometimes symbolized as E. Suppose a system with ...

Quantifying Heat

Thermal Energy&#160;

Microscopically, thermal energy is the kinetic energy associated with the random motion of atoms and molecules. Temperature is a ...

Quantifying Work

As a system undergoes a change, its internal energy can change, and energy can be transferred from the system to the surroundings, or from the ...

Enthalpy

Chemists ordinarily use a property known as enthalpy (H) to describe the thermodynamics of chemical and physical processes. Enthalpy is defined as the sum ...

Thermochemical Equations

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Constant Pressure Calorimetry

Calorimetry is a technique used to measure the amount of heat involved in a chemical or physical process or to measure the heat transferred to or from a ...

Constant Volume Calorimetry

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using-differential-scanning-calorimetry-to-measure-changes-in-enthalpy

Using Differential Scanning Calorimetry to Measure Changes in Enthalpy

Source: Laboratory of Dr. Terry Tritt &#8212; Clemson University
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Hess's Law

There are two ways to determine the amount of heat involved in a chemical change: measure it experimentally, or calculate it from other experimentally ...

Standard Enthalpy of Formation

Enthalpy changes are typically tabulated for reactions in which both the reactants and products are at the same conditions. A standard state is a commonly ...

Enthalpies of Reaction

Hess&rsquo;s law can be used to determine the enthalpy change of any reaction if the corresponding enthalpies of formation of the reactants and products ...

Use mathematics and computational thinking to execute simple calorimetric calculations based on the Law of Conservation of Energy.

PS.Chm.7.2

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Definition and Measurement of Pressure: Atmospheric Pressure, Barometer, and Manometer

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Gas Laws: Boyle's, Gay-Lussac, Charles', Avogadro's, and Ideal Gas Law

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Applications of the Ideal Gas Law: Molar Mass, Density, and Volume

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Mixtures of Gases: Dalton's Law of Partial Pressures and Mole Fractions

Unless individual gases chemically react with each other, the individual gases in a mixture of gases do not affect each other&rsquo;s pressure. Each gas ...

Use mathematics and computational thinking to explain the relationships among pressure, temperature, volume, and quantity of gas, both qualitatively and quantitatively.

PS.Chm.7.3

Understand the relationship among pressure, temperature, volume, and phase.

Water and CO2 Phase Behavior

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