Name: how molecule shape affects intermolecular forces
Uploaded: 2020-09-24T15:06:49.000Z
Duration: 2 min 34 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

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

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

Why do molecules with increasing molar mass show a rising trend in boiling points? The difference can be attributed to their intermolecular forces.&#160;
All substances have dispersion forces &#8212; the weakest intermolecular force. The strength of dispersion forces depends on how easily the atoms can be polarized.
Atoms with higher masses have more electrons and larger electron clouds with electrons dispersing more freely, leading to increased dispersion forces. The stronger the forces, the more energy is required to break the interactions between neighboring atoms, resulting in higher boiling points.
This also explains why an increase in the molar mass in the alkane series translates to a corresponding increase in boiling points and melting points.
However, molar mass alone does not determine the strength of the dispersion forces.&#160;
Despite having the same masses, n-pentane and neopentane have different boiling points as a result of their differing shapes.&#160;
The straight n-pentane provides a larger surface area to interact with the neighboring molecules, while the spherical neopentane has a smaller surface area and consequently reduced dispersion forces. This leads to a lower boiling point.&#160;

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, and solubilities (miscibilities) of substances. Molar mass, molecular shape, and polarity affect the strength of different intermolecular forces, which influence the magnitude of physical properties across a family of molecules.
Temporary attractive forces like dispersion are present in all molecules, whether they are polar or nonpolar. They cause gases to condense (liquefy) and liquids to freeze (solidify) under very low temperature (or high pressure) conditions. Dispersion forces arise from temporary dipoles caused by the asymmetrical distribution of electrons around the atom&#39;s nucleus. Atoms (or molecules) with a greater number of electrons (higher molar mass) display stronger dispersion forces than lighter atoms (or molecules). The melting point and boiling point trend of halogens demonstrate this effect. Moving down the group, from fluorine to iodine, melting points and boiling points increase with increasing atomic size (or mass). This increase may be rationalized by considering how the strength of dispersion forces is affected by the electronic structure of the atoms or molecules in the substance. In a larger atom, the valence electrons are, on average, farther from the nuclei than in a smaller atom. Thus, they are less tightly held and can more easily form the temporary dipoles that produce the attraction. The measure of how easy or difficult it is for another electrostatic charge (for example, a nearby ion or polar molecule) to distort a molecule&#8217;s charge distribution (its electron cloud) is known as polarizability.
A molecule that has a charge cloud that is easily distorted is said to be very polarizable and will have large dispersion forces; one with a charge cloud that is difficult to distort is not very polarizable and will have small dispersion forces. The shapes of molecules also affect the magnitudes of the dispersion forces between them. For example, boiling points for the isomers n-pentane, isopentane, and neopentane are 36 &#176;C, 27 &#176;C, and 9.5 &#176;C, respectively. Even though these compounds are composed of molecules with the same chemical formula, C5H12, the difference in boiling points suggests that dispersion forces in the liquid phase are different, being greatest for n-pentane and least for neopentane. The elongated shape of n-pentane provides a greater surface area available for contact between molecules, resulting in correspondingly stronger dispersion forces. The more compact shape of isopentane offers a smaller surface area available for intermolecular contact and, therefore, weaker dispersion forces. Neopentane molecules are the most compact of the three, offering the least available surface area for intermolecular contact and, hence, the weakest dispersion forces.
Polar substances exhibit dipole&#8211;dipole attractions. The effect of this attraction is apparent when comparing the properties of polar HCl molecules to nonpolar F2 molecules. Both HCl and F2 consist of the same number of atoms and have approximately the same molecular mass. At a temperature of 150 K, molecules of both substances would have the same average KE. However, the dipole&#8211;dipole attractions between HCl molecules are sufficient to cause them to &#8220;stick together&#8221; to form a liquid, whereas the relatively weaker dispersion forces between nonpolar F2 molecules are not, and so this substance is gaseous at this temperature. The higher normal boiling point of HCl (188 K) compared to F2 (85 K) is a reflection of the greater strength of dipole&#8211;dipole attractions between HCl molecules, compared to the attractions between nonpolar F2 molecules.
A special type of dipole&#8211;dipole force&#8212;hydrogen bonds&#8212;have a pronounced effect on the properties of condensed phases (liquids and solids). For example, consider the trends in boiling points for the binary hydrides of group 15 (NH3, PH3, AsH3, and SbH3), group 16 hydrides (H2O, H2S, H2Se, and H2Te), and group 17 hydrides (HF, HCl, HBr, and HI). On progressing down the groups, the polarities of the molecules decrease slightly, whereas the sizes of the molecules increase substantially. The effect of increasingly stronger dispersion forces dominates that of increasingly weaker dipole&#8211;dipole attractions, and the boiling points are observed to increase steadily. Using this trend, the predicted boiling points for the lightest hydride for each group would be about &#8722;120 &#176;C (for NH3), &#8722;80 &#176;C (for H2O), and &#8722;110 &#176;C (for HF). However, the measured boiling points for these compounds are about &#8722;33.34 &#176;C (for NH3), 100 &#176;C (for H2O), and 19.5 &#176;C (for HF), all of which are dramatically higher than the predicted trends. The stark contrast between our na&#239;ve predictions and reality provides compelling evidence for the strength of hydrogen bonding.
Effect of Polarity on Miscibility
Liquids that can be homogeneously mixed in any proportion are said to be miscible. Miscible liquids have similar polarities. Consider, for example, methanol (CH3OH) and water (H2O), two liquids that are polar and capable of hydrogen bonding. On mixing, methanol and water will interact through intermolecular hydrogen bonds and mix; thus, they are miscible. Likewise, nonpolar liquids like hexane (C6H14) and bromine (Br2) are miscible with each other through dispersion forces. The chemical axiom &#8220;like dissolves like&#8221; is useful to predict the miscibility of compounds. Two liquids that do not mix to an appreciable extent are called immiscible. For example, nonpolar hexane is immiscible in polar water. Relatively weak attractive forces between the hexane and water do not adequately overcome the stronger hydrogen bonding forces between water molecules.
This text is adapted from<a href="https://openstax.org/books/chemistry-2e/pages/10-1-intermolecular-forces"> Openstax, Chemistry 2e, Section 10.1: Intermolecular forces</a>. and <a href="https://openstax.org/books/chemistry-2e/pages/11-3-solubility">Section 11.3: Solubility</a>.

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

Education

JoVE Core

Liquids, Solids, and Intermolecular Forces

KEY TERMS AND CONCEPTS

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 dispersion forces are found in atoms or molecules with more electrons, which usually means a higher molar mass.

The halogens show this pattern clearly. As you move from fluorine to iodine, melting points and boiling points rise as atomic size and mass increase. In larger atoms, valence electrons are farther from the nucleus and are easier to distort, so temporary dipoles form more easily. This ease of distortion is called polarizability.

Molecular shape also changes dispersion forces. The isomers n-pentane, isopentane, and neopentane have the same formula, C5H12, but different boiling points: 36 °C, 27 °C, and 9.5 °C. N-pentane is elongated and has more surface area for contact, so its dispersion forces are strongest. Neopentane is the most compact and has the weakest intermolecular contact.

Polar molecules also attract each other through dipole-dipole forces. A good example is HCl compared with nonpolar F2. At 150 K, both substances have the same average kinetic energy, but HCl molecules stick together well enough to form a liquid, while F2 stays a gas. HCl also has a higher normal boiling point, 188 K, than F2, which boils at 85 K.

Hydrogen bonding is a special type of dipole-dipole attraction, and it strongly affects condensed phases such as liquids and solids. The binary hydrides of group 15, group 16, and group 17 show boiling point trends that rise down each group as dispersion forces become stronger and dipole-dipole attractions become weaker. The lightest members would be predicted to boil near -120 °C for NH3, -80 °C for H2O, and -110 °C for HF, but the measured values are much higher: -33.34 °C, 100 °C, and 19.5 °C. That difference is strong evidence for hydrogen bonding.

Polarity also helps predict miscibility, which means liquids mix uniformly in any proportion. Methanol and water are both polar and can hydrogen bond, so they are miscible. Nonpolar liquids such as hexane and bromine are also miscible with each other through dispersion forces. The rule &quot;like dissolves like&quot; is useful here, and it explains why nonpolar hexane is immiscible in polar water.

How Molecule Shape Affects Intermolecular Forces

how-molecule-shape-affects-intermolecular-forces

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,...

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...

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 molecular shape, polarity, and dispersion forces affect melting point, boiling point, hydrogen bonding, and miscibility in high school chemistry.

How Molecule Shape Affects Intermolecular Forces | 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.

Matter and its Interactions

PS.Chm.3.2

Physical Science

PS.PSc.1.3

ohios-learning-standards-and-model-curriculum

Ohio’s Learning Standards and Model Curriculum

Study Of Matter

Classification of matter

Structure And Properties Of Matter

Intermolecular forces of attraction

Tennessee Academic Standards

High School

motion-and-stability-forces-and-interactions

Motion and Stability: Forces and Interactions

CHEM1.PS2.2

Virginia  Standards of Learning

HIgh School

CH.6

CH.6.c

PS.3

PS.3.a

Physics

PH.9

PH.9.e

Arizona Science Standards

HS+C.P1

HS+C.P1U1.4

Texas Essential Knowledge and Skills

CHEM.7

CHEM.7.D

CHEM.11

CHEM.11.C

Alabama Course of Study Science

Matter and Its Interactions

9-12.C.1.1.d

9-12.C.2.2

9-12.C.1.5.e

9-12.PS.3.5.a

new-york-state-science-learning-standards

New York State Science Learning Standards

HS. Structure and Properties of Matter

HS-PS1-3.

Georgia Standards of Excellence

SPS5.

States of Matter

SC2.

chemical-and-physical-properties-of-substances

Chemical and Physical Properties of Substances

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 ...

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

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

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

How Molecule Shape Affects Intermolecular Forces

Description

Transcript

Tags

More videos by Standards