Name: melting, boiling, and state changes
Uploaded: 2020-09-24T15:06:49.000Z
Duration: 2 min 31 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 Transitions

Substances typically exist in one of three phases: solid, liquid, or gas. Transitioning from one phase to another significantly changes how ordered and tightly held the molecules are.
Molecules transition between different phases when their internal energy allows them to be in either state. This depends on factors like the strength of the intermolecular forces in the more-condensed state and the pressure exerted on the substance.
Temperature is a reflection of internal energy, so phase transition points are often described in terms of temperature at a certain pressure.
For instance, compare water and acetone. While water exhibits strong hydrogen bonds, acetone molecules display weak dipole-dipole forces.
Due to the stronger forces of attraction, more heat is required to turn water into steam. This explains why at any given pressure, the boiling point of acetone is lower than that of water.
Phase transitions happen molecule by molecule, so the phases coexist during the transition.
Until the transition of the bulk is complete, the temperature does not change even though heat is flowing to or from the substance.
A similar observation is made when heat is supplied to water. The tempera

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 solid pack closely together, generally in a pattern. The particles vibrate about their fixed positions but do not move or squeeze past their neighbors. In liquids, although the particles are closely spaced, they are randomly arranged. The position of the particles are not fixed&#8212;that is, they are free to move past their neighbors to occupy different locations. Because the particles are close together in the solid and liquid states, these are referred to as the condensed states or condensed phases. In these states, substances exhibit relatively strong intermolecular forces. In gases, the interparticle forces of attractions are weak. The particles of a gas are not constrained by their neighbors; the particles are free to move and, under normal conditions, are separated by large distances.
The internal energy of a substance&#8212;the total kinetic energy of all its molecules&#8212;depends on the strength of the intermolecular forces in the condensed phases and the pressure exerted on the substance. The internal energy of a substance is the highest in a gaseous state, the lowest in a solid state, and intermediate in a liquid.
Phase transitions are caused by changes in physical conditions, such as in temperature and/or pressure, which affect the strength of intermolecular forces. For example, the addition of heat to a substance causes its particle&#39;s thermal energy (or the energy of motion) to increase, overcoming the attractive intermolecular forces between them. A solid melts when its temperature rises to the point at which the particles vibrate fast enough to move out of their fixed positions. This phase transition is called melting, and the point at which it occurs is the solid&#39;s melting point. As the temperature increases further, the particles move faster until they finally escape into the gaseous state. This is vaporization, and the point at which it occurs is the liquid&#39;s boiling point.
The phase transition point and the energy change associated with the transition depend on the intermolecular forces that exist in the substance. At a given pressure, substances with stronger intermolecular forces require more energy to overcome them and, therefore, undergo phase changes at higher temperatures. The energy required to cause the complete phase transition of one mole of a substance without a change in temperature is called the molar heat or molar enthalpy of that transition. For example, the energy required to vaporize one mole of a liquid is called the molar enthalpy of vaporization.
Transitions that occur by absorbing energy are exothermic, and their enthalpy values are negative. On the other hand, transitions that occur by releasing energy are endothermic, and their enthalpy values are positive. For example, while the molar enthalpy of vaporization is positive, the molar enthalpy of condensation is negative.
Because a substance transforms from one phase to another molecule by molecule, during a phase transition, the two phases coexist; and the temperature of the substance stays constant, despite the continuous supply of heat. After the transition of the bulk completes, the temperature of the substance rises.
When phase transitions occur in a closed system, the opposing transitions occur at equal rates, leading to a state of dynamic equilibrium.

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 ...

Education

JoVE Core

Liquids, Solids, and Intermolecular Forces

KEY TERMS AND CONCEPTS

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 because the attractions between particles are weak. Gas particles are not held close by neighboring particles, so they move freely and are usually far apart under normal conditions. A substance’s internal energy, which is the total kinetic energy of its molecules, depends on intermolecular forces in condensed phases and on the pressure applied to the substance. Internal energy is highest in a gas, lowest in a solid, and intermediate in a liquid.

Changes in temperature or pressure can trigger a phase transition by changing the strength of intermolecular forces. When heat is added, particle thermal energy increases. The particles move faster, and the attractive forces between them can be overcome. A solid melts when its particles vibrate fast enough to leave their fixed positions. The temperature at which this happens is the melting point.

If heating continues, the particles gain even more energy until they escape into the gas phase. This process is vaporization, and the temperature where it occurs is the boiling point. The temperature of a transition and the energy needed for it depend on the intermolecular forces in the substance. At the same pressure, substances with stronger intermolecular forces need more energy and change phase at higher temperatures.

The energy required to change one mole of a substance completely from one phase to another without a temperature change is called the molar heat, or molar enthalpy, of the transition. The molar enthalpy of vaporization is the energy needed to vaporize one mole of a liquid. Transitions that absorb energy are endothermic and have negative enthalpy values. Transitions that release energy are exothermic and have positive enthalpy values. For example, the molar enthalpy of vaporization is positive, while the molar enthalpy of condensation is negative.

During a phase transition, a substance changes molecule by molecule, so two phases coexist while the change is happening. The temperature stays constant even if heat keeps being supplied. After the bulk of the substance has completed the transition, the temperature rises again. In a closed system, the forward and reverse transitions can happen at equal rates. This creates dynamic equilibrium.

Melting, Boiling, and State Changes

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...

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...

phase diagrams

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...

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 melting, boiling, and other state changes depend on particle motion, intermolecular forces, enthalpy, and dynamic equilibrium in solids, liquids, and gases.

Melting, Boiling, and State Changes | 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

ohios-learning-standards-and-model-curriculum

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

SC.912.P.12

SC.912.P.12.11

Chemistry and Physics

IPC.7

IPC.7.C

Georgia Standards of Excellence

SPS5.

Energy and Phase Transitions

SPS7.

Phase Changes and Energy

SC5.

Heating Curves and Phase Changes

Classifying Matter by State

Chemistry is the study of matter and the changes it undergoes. Matter is anything that has mass and occupies space. Matter is all around us; the air, ...

Classifying Matter by Composition

Matter: Pure Substances and Mixtures
According to its composition, the matter can be classified into two broad categories &mdash; pure substances and ...

physical-and-chemical-properties-of-matter

Physical and Chemical Properties of Matter

The characteristics that enable us to distinguish one substance from another are called properties.

Physical Properties of Matter

A physical property is ...

Molecular Comparison of Gases, Liquids, and Solids

Particles in a solid are tightly packed together (fixed shape) and often arranged in a regular pattern; in a liquid, they are close together with no ...

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: 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

When a substance&#8212;isolated from its environment&#8212;is subjected to heat changes, corresponding changes in temperature and phase of the substance ...

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. ...

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

For a chemical reaction (the system) carried out at constant pressure &ndash; with the only work done caused by expansion or contraction &ndash; the ...

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

Calorimeters are useful to determine the heat released or absorbed by a chemical reaction. Coffee cup calorimeters are designed to operate at constant ...

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
Differential Scanning Calorimetry (DSC) is a method of thermodynamic analysis based on ...

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

definition-measurement-pressure-atmospheric-pressure-barometer

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

Gas pressure is caused by force exerted by gas molecules colliding with the surfaces of objects. Although the force of each collision is very small, any ...

gas-laws-boyle-s-gay-lussac-charles-avogadro-s-and-ideal-gas-law

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

Through experiments, scientists established the mathematical relationships between pairs of variables, such as pressure and temperature, pressure and ...

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.

Melting, Boiling, and State Changes

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