Name: how bond energy predicts reaction heat
Uploaded: 2020-09-03T13:38:28.000Z
Duration: 2 min 49 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

Matter and its Interactions

PS.Chm.3.1

Bond Energies and Bond Lengths

Every chemical reaction is associated with a change in enthalpy, which helps to determine whether energy is released or required during the reaction. This enthalpy change can be estimated using average bond energies.

The energy required to break a specific bond in 1 mole of a gaseous chemical compound is called bond energy and is expressed in kJ/mol. This energy depends on the type of bonded atoms and the number of shared pairs of electrons. 

Bond energy is usually expressed as an average of bond energies of the same bond in different compounds. The average bond energy can be used to determine whether a chemical reaction is exothermic or endothermic. 

Consider ethene and bromine reacting to form 1,2-dibromoethane. Initially, the carbon double bond and the bromine single bond break &#8212; a process requiring energy input, which increases the potential energy of atoms. Therefore, bond-breaking is an endothermic process with a positive change in enthalpy.

Subsequently, new bonds are formed between carbon and bromine atoms yielding the product. Bond formation increases the stability of the molecule by reducing the potential energy. Therefore, bond formation is an exothermic process, causing a negative change in enthalpy. 

According to Hess&#8217;s law, the sum of enthalpy changes of reactants and products is equal to the overall enthalpy change of the reaction. The enthalpy change of reactants is the sum of enthalpy of bonds broken, whereas the enthalpy change of products is the sum of enthalpy of new bonds formed. 

Therefore, the formation of 1,2-dibromoethane, with an enthalpy of +255 kJ/mol, is an endothermic reaction. 

In addition to bond enthalpies, the type of bonds and bonded atoms also influence the bond length, which is the average distance between the nuclei of two bonded atoms.

Consider the different bonds between two nitrogen atoms and two carbon atoms. Atoms with multiple bonds, such as triple bonds, are held more tightly together, leading to shorter and stronger bonds. Consequently, the molecule is more stable and requires higher energy to dissociate. 

Generally, the bond strength is indirectly proportional to the bond length, with some exceptions.

bond energies and bond lengths

Stable molecules exist because covalent bonds hold the atoms together. The strength of a covalent bond is measured by the energy required to break it, that is, the energy necessary to separate the bonded atoms. Separating any pair of bonded atoms requires energy &#8212; the stronger a bond, the greater the energy required to break it.

The energy required to break a specific covalent bond in one mole of gaseous molecules is called the bond energy or the bond dissociation energy. The bond energy for a diatomic molecule, DX&#8211;Y, is defined as the standard enthalpy change for the endothermic reaction:

<img alt="Chemical equilibrium diagram, dissociation reaction XY(g) to X(g) + Y(g), enthalpy change &#916;H&#176;." src="/files/core/CDNSource/lessons/1067/1067_Equations-01.png" style="height: 40px;" data-altupdatedat="1761830439000" data-alt="Eq1" />

For example, the bond energy of the pure covalent H&#8211;H bond, DH&#8211;H, is 436 kJ/mol of H&#8211;H bonds broken:

<img alt="H2 gas dissociation equation; enthalpy change &#916;H=436kJ. Chemically descriptive formula." src="/files/core/CDNSource/lessons/1067/1067_Equations-02.png" style="height: 40px;" data-altupdatedat="1761830439000" data-alt="Eq2" />

Molecules with three or more atoms have two or more bonds. The sum of all bond energies in such a molecule is equal to the standard enthalpy change for the endothermic reaction that breaks all the bonds in the molecule. For example, the sum of the four C&#8211;H bond energies in CH4, 1660 kJ, is equal to the standard enthalpy change of the reaction:

<img alt="Chemical reaction equation: methane decomposition to carbon and hydrogen, &#916;H&#176;=1660 kJ, diagram." src="/files/core/CDNSource/lessons/1067/1067Image_1.svg" style="height: 60px;" data-altupdatedat="1761830439000" data-alt="Figure1" />

The average C&#8211;H bond energy, DC&#8211;H, is 1660/4 = 415 kJ/mol because there are four moles of C&#8211;H bonds broken per mole of the reaction. Although the four C&#8211;H bonds are equivalent in the original molecule, they do not each require the same energy to break; once the first bond is broken (which requires 439 kJ/mol), the remaining bonds are easier to break. The 415 kJ/mol value is the average, not the exact value required to break any one bond.

The strength of a bond between two atoms increases as the number of electron pairs in the bond increases. Generally, as the bond strength increases, the bond length decreases. Thus, triple bonds are stronger and shorter than double bonds between the same two atoms; likewise, double bonds are stronger and shorter than single bonds between the same two atoms. When one atom bonds to various atoms in a group, the bond strength typically decreases as we move down the group. For example, C&#8211;F is 439 kJ/mol, C&#8211;Cl is 330 kJ/mol, and C&#8211;Br is 275 kJ/mol. 

Bond energies can be used to calculate approximate enthalpy changes for reactions, also called Bond enthalpies, where enthalpies of formation are not available. Calculations of this type will also tell whether a reaction is exothermic or endothermic. An exothermic reaction (&#916;H negative, heat produced) results when the bonds in the products are stronger than the bonds in the reactants. An endothermic reaction (&#916;H positive, heat absorbed) results when the bonds in the products are weaker than those in the reactants.

The enthalpy change, &#916;H, for a chemical reaction is approximately equal to the sum of the energy required to break all bonds in the reactants (energy &#8220;in&#8221;, positive sign) plus the energy released when all bonds are formed in the products (energy &#8220;out,&#8221; negative sign). This can be expressed mathematically in the following way:

<img alt="Enthalpy calculation formula, &#916;H=&#931;D bonds broken - &#931;D bonds formed, chemical reaction diagram." src="/files/core/CDNSource/lessons/1067/1067_Equations-03.png" style="height: 40px;" data-altupdatedat="1761830439000" data-alt="Eq3" />

In this expression, the symbol &#425; means &#8220;the sum of&#8221; and D represents the bond energy in kJ/mol, which is always a positive number. The bond energy will depend on whether the particular bond is a single, double, or triple bond. Thus, in calculating enthalpies in this manner, it is important that the bonding in all reactants and products are considered. Because D values are typically averages for one type of bond in many different molecules, this calculation provides a rough estimate, not an exact value, for the enthalpy of reaction.

Consider the following reaction:

<img alt="Chemical equation: H&#8322;(g) + Cl&#8322;(g) &#8594; 2HCl(g); synthesis reaction diagram." src="/files/core/CDNSource/lessons/1067/1067_Equations-04.png" style="height: 40px;" data-altupdatedat="1761830439000" data-alt="Eq4" />

or

<img alt="Chemical reaction equation for hydrogen chloride: H&#8322;(g) + Cl&#8322;(g) &#8594; 2HCl(g)." src="/files/core/CDNSource/lessons/1067/1067_Equations-05.png" style="height: 40px;" data-altupdatedat="1761830439000" data-alt="Eq5" />

To form two moles of HCl, one mole of H&#8211;H bonds and one mole of Cl&#8211;Cl bonds must be broken. The energy required to break these bonds is the sum of the bond energy of the H&#8211;H bond (436 kJ/mol) and the Cl&#8211;Cl bond (243 kJ/mol). During the reaction, two moles of H&#8211;Cl bonds are formed (bond energy = 432 kJ/mol), releasing 2 &#215; 432 kJ; or 864 kJ. Because the bonds in the products are stronger than those in the reactants, the reaction releases more energy than it consumes:

<img alt="&#916;H calculation formula; &#931;D bonds broken-&#931;D bonds formed; thermochemistry equation." src="/files/core/CDNSource/lessons/1067/1067_Equations-06.png" style="height: 40px;" data-altupdatedat="1761830439000" data-alt="Eq6" />

<img alt="Bond energy calculation formula, enthalpy change, chemical reaction equation, study energetics." src="/files/core/CDNSource/lessons/1067/1067_Equations-07.png" style="height: 40px;" data-altupdatedat="1761830439000" data-alt="Eq7" />

This excess energy is released as heat, so the reaction is exothermic.

This text is adapted from <a href="https://openstax.org/books/chemistry-2e/pages/7-5-strengths-of-ionic-and-covalent-bonds">Openstax, Chemistry 2e, Section 7.5: Bond Strength: Covalent Bonds.</a>

Stable molecules exist because covalent bonds hold the atoms together. The strength of a covalent bond is measured by the energy required to break it, ...

Education

JoVE Core

Chemical Bonding: Basic Concepts

KEY TERMS AND CONCEPTS

how bond energy predicts reaction heat

Bond energy helps explain why molecules stay together and how much heat a reaction may release or absorb. A covalent bond is stronger when more energy is needed to separate the atoms. The energy needed to break one specific covalent bond in one mole of gaseous molecules is called bond energy, or bond dissociation energy.

For a diatomic molecule, bond energy is the standard enthalpy change for an endothermic reaction that breaks that bond. The H-H bond is a common example. Its bond energy is 436 kJ/mol. Molecules with three or more atoms have several bonds, so the total bond energy is the sum of all the bonds that must be broken.

Methane shows why average bond energy is useful. Breaking all four C-H bonds in CH4 requires 1660 kJ, so the average C-H bond energy is 415 kJ/mol. That is an average value, not the exact energy for every single C-H bond. After the first bond breaks, which requires 439 kJ/mol, the remaining bonds are easier to break.

Bond strength usually increases as the number of electron pairs in the bond increases. As bond strength increases, bond length usually decreases. Triple bonds are therefore stronger and shorter than double bonds between the same atoms, and double bonds are stronger and shorter than single bonds. Bond strength can also decrease down a group. For example, C-F is 439 kJ/mol, C-Cl is 330 kJ/mol, and C-Br is 275 kJ/mol.

Bond energies can also estimate reaction enthalpy, especially when enthalpies of formation are not available. These calculations can show whether a reaction is exothermic or endothermic. An exothermic reaction has a negative ΔH and releases heat. An endothermic reaction has a positive ΔH and absorbs heat.

To estimate ΔH, add the energy needed to break all bonds in the reactants and subtract the energy released when bonds form in the products. In equation form, this is the sum of bond energies broken plus the sum of bond energies formed, with formed bonds counted as negative. Because bond energy values are usually averages for a bond type, the result is an estimate, not an exact value.

The formation of hydrogen chloride is a useful example. To make two moles of HCl, one mole of H-H bonds and one mole of Cl-Cl bonds must be broken. That requires 436 kJ/mol plus 243 kJ/mol. At the same time, two moles of H-Cl bonds form and release 864 kJ because each H-Cl bond has a bond energy of 432 kJ/mol. Since the product bonds are stronger than the reactant bonds, the reaction releases more energy than it uses and is exothermic.

How Bond Energy Predicts Reaction Heat

types of chemical bonds

Chemical Bonding and Bond Properties

Chemical bonding explains how atoms and ions join to form molecules and compounds. American chemist Gilbert N. Lewis pioneered chemical bonding theory with the Lewis model. This model helps explain why bonds form, why some are stronger than others, and why compounds have fixed formulas like H2O instead of H4O. Ionic bonding involves ions, which are atoms or molecules with an electrical charge. A cation is a positive ion formed when a neutral atom loses one or more electrons from its valence...

the born-haber cycle

Predicting Ionic Stability with Lattice Energy

Lattice energy helps explain why ionic compounds are stable. It measures the strength of attraction between positive and negative ions in a crystal. In one common definition, lattice energy is the energy required to separate one mole of an ionic solid into gaseous ions, so it has a positive value. Another definition uses the opposite sign and describes the energy released when gaseous ions form a lattice. When checking data, it is important to confirm which convention is being used. In either...

lewis symbols and the octet rule

Valence Electrons, Lewis Symbols, and Bonding

Chemical bonds form when atoms or ions interact in ways that lower the potential energy of a molecule. Gilbert N. Lewis created a simple model to show bond formation using valence electrons, the electrons in the outer shell. In this model, electrons are shared or transferred to help atoms reach a stable electron configuration. The Lewis model helps describe two common bond types. An ionic bond forms when electrons move from a metal to a nonmetal. A covalent bond forms when two nonmetals share...

electronegativity

Electronegativity and Polar Bonds

Electronegativity helps explain why some covalent bonds are nonpolar and others are polar covalent. It is a measure of how strongly an atom attracts shared electrons toward itself. That pull controls how electron density is distributed in a bond. Linus Pauling proposed the first electronegativity values by studying bond energies in heteronuclear molecules such as hydrogen fluoride. He expected the bond energy of HF to be about the average of the bond energies of H2 and F2, or 296 kJ/mol. The...

ionic bonding and electron transfer

Ionic Compounds and Crystal Lattices

Ionic compounds form when positive and negative ions attract each other. A cation is a positive ion, and it forms when an atom loses one or more valence electrons. An anion is a negative ion, and it forms when an atom gains one or more valence electrons. Compounds made of these ions are called ionic compounds, or salts, and the force between them is an ionic bond. The properties of ionic solids help show how strong this attraction is. They usually form rigid, brittle crystals and have high...

trends in lattice energy: ion size and charge

trends-in-lattice-energy-ion-size-and-charge

Lattice Energy Patterns in Ionic Solids

Lattice energy shows how strongly ions hold an ionic solid together. It is the energy required to separate one mole of a solid into its gaseous ions. For sodium chloride, this process requires 769 kJ per mole. The opposite convention may also be used, where lattice energy is the energy released when gaseous ions form the solid. In either case, a larger magnitude means a more stable ionic compound. For ionic crystals, lattice energy follows a pattern based on charge and ion size. The...

covalent bonding and lewis structures

Shared Electrons in Covalent Molecules

Covalent molecules form when atoms share electrons instead of transferring them. This bonding usually happens between nonmetal atoms with similar tendencies to attract electrons. Because the atoms pull on the electrons in a shared way, the bond is held together by mutual attraction. Covalent compounds often have physical properties that differ from ionic compounds. The attraction between neutral molecules is weaker than the attraction between charged ions. As a result, many covalent compounds...

bond polarity, dipole moment, and percent ionic character

bond-polarity-dipole-moment-and-percent-ionic-character

How Bond Polarity Shapes Molecular Dipoles

Bond polarity depends on electronegativity, the ability of an atom to attract shared electrons. The absolute value of the electronegativity difference, or ΔEN, gives a rough guide to bond type. When the difference is very small or zero, the bond is covalent and nonpolar. When it is larger, the bond may be polar covalent or ionic. Examples show how this works in real compounds. H–H has a ΔEN of 0, so it is nonpolar. H–Cl has a ΔEN of 0.9, so it is polar covalent. Na–Cl has a ΔEN of 2.1, so it...

lewis structures of molecular compounds and polyatomic ions

lewis-structures-of-molecular-compounds-and-polyatomic-ions

Drawing Lewis Structures with Valence Electron Counts

Drawing Lewis structures starts with counting valence electrons and building the atom arrangement step by step. Valence electrons are the outer-shell electrons that help show how atoms bond. For ions, a positive charge means one electron is removed for each plus sign, and a negative charge means one electron is added for each minus sign. The next step is to sketch the skeleton structure. The least electronegative atom usually goes in the center, while hydrogen is almost never central. Fluorine...

resonance

Resonance Forms in Lewis Structures

Resonance forms in Lewis structures help explain molecules and ions that cannot be shown correctly with just one drawing. In nitrite ion, NO2−, the Lewis structure can be drawn in two ways. The N-O and N=O bonds switch positions in the two drawings, but the atoms stay in the same arrangement. If nitrite really had one single bond and one double bond, the two bond lengths should be different. A double bond is shorter and stronger than a single bond between the same atoms. Experiments show that...

formal charges

Lewis Structures with Formal Charge

Lewis structures can sometimes show more than one reasonable arrangement for a molecule or polyatomic ion. Formal charge helps chemists choose the best Lewis structure when that happens. It is a useful way to compare possible structures in chemistry. Formal charge is the hypothetical charge an atom would have if bonding electrons were shared evenly between the atoms. It can also be found by starting with the valence electrons of the neutral atom, subtracting the nonbonding electrons, and then...

exceptions to the octet rule

Octet Rule Exceptions in Lewis Structures

Some covalent molecules do not follow the octet rule in their Lewis structures. Their central atoms can fall into three groups: odd-electron molecules, electron-deficient molecules, and hypervalent molecules. These cases show why a Lewis structure may not always give every atom eight valence electrons. Odd-electron molecules have an odd number of valence electrons, so one electron remains unpaired. Molecules with an odd number of electrons are called radicals. Nitric oxide, NO, is one example,...

Bond energy helps explain why molecules stay together and how much heat a reaction may release or absorb. A covalent bond is stronger when more energy is needed to separate the atoms. The energy needed to break one specific covalent bond in one mole of gaseous molecules is called bond energy, or bond dissociation energy. For a diatomic molecule, bond energy is the standard enthalpy change for an endothermic reaction that breaks that bond. The H-H bond is a common example. Its bond energy is...

bonding in metals

Electron Sea Model in Metals

Electron sea model explains how metallic bonding holds metals together. In this model, most metal atoms do not have enough valence electrons to form ionic or covalent bonds in the usual way. Their valence electrons are held loosely because metals have low electronegativity, and their ionization energy is low, so electrons are removed easily. When a metal atom loses these outer electrons, it becomes a positively charged metal ion. The freed electrons form a delocalized electron cloud around...

key terms and concepts

Learn how bond energy, bond length, and bond dissociation energy help predict whether reactions absorb or release heat, with examples like H-H, CH4, and HCl.

How Bond Energy Predicts Reaction Heat | 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.

Pennsylvania State Academic Standards

Physical Science

Chemical Reactions

3.2.9-12.D

3.2.9-12.E

Louisiana Student Standards for Science

High School

Matter And Its Interactions

HS-PS1-4

ohios-learning-standards-and-model-curriculum

Ohio’s Learning Standards and Model Curriculum

Study Of Matter

Bonding and compounds

Structure And Properties Of Matter

Chemical bonding

Tennessee Academic Standards

Energy

CHEM1.PS3.3

Virginia  Standards of Learning

HIgh School

CH.3

CH.3.d

Physics

PH.9

PH.9.e

Minnesota Academic Standards in Science

3.1.1 9C.2.2.1.1

new-jersey-student-learning-standards-for-science

New Jersey Student Learning Standards for Science

Matter and Its Interactions

Alabama Course of Study Science

motion-and-stability-forces-and-interactions

Motion and Stability: Forces and Interactions

9-12.C.2.2.b

Next Generation Science Standards

Texas Essential Knowledge and Skills

Chemistry and Physics

IPC.7

IPC.7.A

south-carolina-college-and-career-ready-science-standards

South Carolina College and Career Ready Science Standards

C-PS1-4

new-york-state-science-learning-standards

New York State Science Learning Standards

HS. Chemical Reactions

HS-PS1-2.

HS-PS1-4.

HS-PS1-12.

Georgia Standards of Excellence

SC2.

Energy Changes in Chemical Reactions

Scientific Laws and Theories

Scientific Laws

In science, a law is defined as a concise, verbal or mathematical, statement that summarizes a vast number of experimental observations. ...

The Scientific Method

Chemistry is an empirical science. Scientists often pose questions to understand the chemistry in everyday life and seek answers to these questions. To ...

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 ...

The Atomic Theory of Matter

The earliest recorded discussion of the basic structure of matter comes from ancient Greek philosophers. Leucippus and Democritus argued that all matter ...

Subatomic Particles

Dalton was only partially correct about the particles that make up matter. All matter is composed of atoms, and atoms are composed of three smaller ...

Elements: Chemical Symbols and Isotopes

A chemical symbol is an abbreviation used to indicate an element or an atom of an element. For example, the symbol for mercury is Hg. The same symbol is ...

Use models to explain how the scientific understanding of atomic structure has evolved.

PS.Chm.1.1

Radioactivity and Nuclear Equations

Nuclear chemistry is the study of reactions that involve changes in nuclear structure. The nucleus of an atom is composed of protons and, except for ...

Types of Radioactivity

The most common types of radioactivity are &#945; decay, &#946; decay, &#947; decay, neutron emission, and electron capture.
Alpha (&#945;) decay is the ...

Nuclear Stability

Protons and neutrons, collectively called nucleons, are packed together tightly in a nucleus. With a radius of about 10&#8722;15 meters, a nucleus is ...

Nuclear Binding Energy

The difference between the calculated and experimentally measured masses is known as the mass defect of the atom. In the case of helium-4, the mass defect ...

Radioactive Decay and Radiometric Dating

Radioactivity is a spontaneous disintegration of an unstable nuclide and is a random process, as all the nuclei in the sample do not decay simultaneously. ...

Nuclear Fission

Many heavier elements with smaller binding energies per nucleon can decompose into more stable elements that have intermediate mass numbers and larger ...

Nuclear Power

Controlled nuclear fission reactions are used to generate electricity. Any nuclear reactor that produces power via the fission of uranium or plutonium by ...

Nuclear Fusion

The process of converting very light nuclei into heavier nuclei is also accompanied by the conversion of mass into large amounts of energy, a process ...

Nuclear Transmutation

Nuclear transmutation is the conversion of one nuclide into another. It can occur by the radioactive decay of a nucleus, or the reaction of a nucleus with ...

Use models to compare nuclear reactions including alpha decay, beta decay and gamma decay; nuclear fusion and nuclear fission.

PS.Chm.1.2

Atomic Orbitals

An atomic orbital represents the three-dimensional regions in an atom where an electron has the highest probability to reside. The radial distribution ...

The Pauli Exclusion Principle

The arrangement of electrons in the orbitals of an atom is called its electron configuration. We describe an electron configuration with a symbol that ...

The Energies of Atomic Orbitals

In an atom, the negatively charged electrons are attracted to the positively charged nucleus. In a multielectron atom, electron-electron repulsions are ...

The Aufbau Principle and Hund's Rule

To determine the electron configuration for any particular atom, we can build the structures in the order of atomic numbers. Beginning with hydrogen, and ...

electron-configuration-of-multielectron-atoms

Electron Configuration of Multielectron Atoms

The alkali metal sodium (atomic number 11) has one more electron than the neon atom. This electron must go into the lowest-energy subshell available, the ...

Use models to explain how electrons are distributed in atoms.

PS.Chm.1.3

Periodic Classification of the Elements

The periodic table arranges atoms based on increasing atomic number so that elements with the same chemical properties recur periodically. When their ...

The Periodic Table

As early chemists discovered more elements, they realized that various elements could be grouped by their similar chemical behaviors. One such grouping ...

Atomic Mass

Atoms &mdash; and the protons, neutrons, and electrons that compose them &mdash; are extremely small. For example, a carbon atom weighs less than 2 ...

Molar Mass

The identity of a substance is defined not only by the types of atoms or ions it contains but by the quantity of each type of atom or ion. For example, ...

Use the Periodic Table as a model to predict the relative properties of elements based on the pattern of valence electrons in the outermost energy levels of atoms.

PS.Chm.2.1

atomic-radii-and-effective-nuclear-charge

Atomic Radii and Effective Nuclear Charge

The elements in groups of the periodic table exhibit similar chemical behavior. This similarity occurs because the members of a group have the same number ...

Ionic Radii

Ionic radius is the measure used to describe the size of an ion. A cation always has fewer electrons and the same number of protons as the parent atom; it ...

Ionization Energy

The amount of energy required to remove the most loosely bound electron from a gaseous atom in its ground state is called its first ionization energy ...

Electron Affinity

The electron affinity (EA) is the energy change for adding an electron to a gaseous atom to form an anion (negative ion).



This process can be either ...

Construct an explanation to infer the atomic size, reactivity, electronegativity, and ionization energy of an element based on its position in the Periodic Table.

PS.Chm.2.2

Types of Chemical Bonds

Chemical bonding theories were pioneered by American chemist Gilbert&#xA0;N. Lewis. He developed a model called the Lewis model to explain the type and ...

Lewis Symbols and the Octet Rule

Chemical bonds are complex interactions between two or more atoms or ions, which reduce the potential energy of the molecule. Gilbert N. Lewis developed a ...

Ionic Bonding and Electron Transfer

Ions are atoms or molecules bearing an electrical charge. A cation (a positive ion) forms when a neutral atom loses one or more electrons from its valence ...

The Born-Haber Cycle

Lattice Energy&#xA0;

An ionic compound is stable because of the electrostatic attraction between its positive and negative ions. The lattice energy of a ...

Trends in Lattice Energy: Ion Size and Charge

An ionic compound is stable because of the electrostatic attraction between its positive and negative ions. The lattice energy of a compound is a measure ...

Covalent Bonding and Lewis Structures

Compared to ionic bonds, which results from the transfer of electrons between metallic and nonmetallic atoms, covalent bonds result from the mutual ...

Analyze and interpret data to explain the mechanisms and properties of the two main types of intramolecular (ionic and covalent) bonds.

Analyze the structure of atoms and isotopes.

How Bond Energy Predicts Reaction Heat

Description

Transcript

Tags

More videos by Standards