Name: octet rule exceptions in lewis structures
Uploaded: 2020-09-03T13:38:28.000Z
Duration: 2 min 55 s
Channel: JoVE

K12 Standards

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

North Carolina K-12 Science Standard Course of Study

9-12

Chemistry

Matter and its Interactions

PS.Chm.3.1

Exceptions to the Octet Rule

Reguła oktetu wyjaśnia wiązania chemiczne w związkach grupy głównej, przewidując, że każdy atom osiąga konfigurację 8-elektronową. Istnieją jednak trzy główne wyjątki od tej reguły. 

Pierwszym wyjątkiem są gatunki o nieparzystych elektronach. Większość cząsteczek i jonów ma parzystą liczbę elektronów. Jednak niektóre cząsteczki, zwane rodnikami, mają jeden lub więcej niesparowanych elektronów. Rodniki z nieparzystą liczbą niesparowanych elektronów nie mogą osiągnąć oktetów. 

Anion ponadtlenkowy, rodnik z jednym niesparowanym elektronem, ma 13 elektronów walencyjnych. Może być reprezentowany przez dwie przyczyniające się struktury, w których jeden tlen ma tylko siedem elektronów, a zatem nie może dotrzeć do oktetu. 

Drugim wyjątkiem są atomy, które tworzą niekompletny oktet. Na przykład wodór, hel i lit mają tendencję do łączenia się w duet, podczas gdy pierwiastki z grupy 2 i 13, takie jak beryl i bor, często tworzą cząsteczki z odpowiednio czterema i sześcioma elektronami wokół siebie. 

Weźmy pod uwagę chlorek glinu, który ma 24 elektrony walencyjne. Podczas gdy wszystkie atomy chloru docierają do oktetu, aluminium otrzymuje tylko 6 elektronów walencyjnych - niekompletny oktet. 

Chociaż chlorek glinu jest stabilny, reaguje z cząsteczkami takimi jak amoniak, które mają niewspółdzieloną parę elektronów. Azot w amoniaku przekazuje swoją samotną parę aluminium, tworząc specjalne wiązanie zwane koordynacyjnym wiązaniem kowalencyjnym lub celownikowym . 

Trzecim wyjątkiem są pierwiastki, które mogą pomieścić więcej niż 8 elektronów walencyjnych lub rozszerzony oktet. Pierwiastki te znajdują się w trzecim rzędzie układu okresowego pierwiastków i poniżej. 

Pierwiastki, takie jak fosfor, siarka lub jod, mają dostęp do orbitali d, co pozwala im pomieścić więcej niż 8 elektronów walencyjnych - często do 12 lub 14. 

Weźmy pod uwagę anion tetrachlorojodku, który ma 36 elektronów walencyjnych. Nawet po przypisaniu wiążących par elektronów i spełnieniu oktetu dla wszystkich atomów, 4 elektrony walencyjne pozostają nieprzypisane. 

Elektrony te są umieszczane na centralnym atomie jodu, dając rozszerzony oktet z 12 elektronami. Cząsteczki z więcej niż 8 elektronami walencyjnymi wokół centralnego atomu nazywane są hiperwalencyjnymi.

Pamiętaj, że pierwiastki z drugiego rzędu układu okresowego pierwiastków, takie jak węgiel lub tlen, mają tylko orbitale s i p i nigdy nie tworzą związków hiperwalencyjnych, ponieważ łącznie mogą pomieścić tylko do 8 elektronów walencyjnych.

wyjątki od reguły oktetu

Wiele cząsteczek kowalencyjnych ma atomy centralne, które nie mają ośmiu elektronów w swoich strukturach Lewisa. Cząsteczki te dzielą się na trzy kategorie:

<ol style="display: flex; flex-direction: column; gap: 10px;">
	<li>Cząsteczki o nieparzystych elektronach mają nieparzystą liczbę elektronów walencyjnych i dlatego mają niesparowany elektron.</li>
	<li>Cząsteczki z niedoborem elektronów mają centralny atom z mniejszą liczbą elektronów, niż jest to potrzebne w przypadku konfiguracji gazu szlachetnego.</li>
	<li>Cząsteczki hiperwalentne mają atom centralny, który ma więcej elektronów niż potrzeba do konfiguracji gazu szlachetnego.</li>
</ol>

<h4>Cząsteczki nieparzystych elektronów</h4>

Cząsteczki zawierające nieparzystą liczbę elektronów nazywane są rodnikami. Tlenek azotu NO jest przykładem cząsteczki o elektronach nieparzystych; powstaje w silnikach spalinowych, gdy tlen i azot reagują w wysokich temperaturach.

Aby narysować strukturę Lewisa dla cząsteczki o elektronach nieparzystych, takiej jak NO, rozważa się następujące kroki:

<ol style="display: flex; flex-direction: column; gap: 10px;">
<li>Określ całkowitą liczbę elektronów walencyjnych (powłoki zewnętrznej). Suma elektronów walencyjnych wynosi 5 (z N) + 6 (z O) = 11. Liczba nieparzysta oznacza, że jest to wolny rodnik, w którym nie każdy atom ma osiem elektronów na swojej powłoce walencyjnej.</li>

<li>Narysuj szkieletową strukturę cząsteczki. Można łatwo narysować strukturę szkieletową z pojedynczym wiązaniem N – O.</li>

<li>Rozłóż pozostałe elektrony jako wolne pary na atomach końcowych. W tym przypadku nie ma atomu centralnego, więc elektrony są rozmieszczone wokół obu atomów. W takich sytuacjach osiem elektronów jest przypisanych do atomu bardziej elektroujemnego; zatem tlen ma wypełnioną powłokę walencyjną:

	<img alt="Figure1" src="https://www.jove.com/files/core/CDNSource/lessons/1066/1066Image_1.svg" style="height: 40px;">


</li>


<li>Umieść wszystkie pozostałe elektrony na atomie centralnym. Ponieważ nie ma pozostałych elektronów, ten krok nie ma zastosowania.</li>

<li>Zmień rozmieszczenie elektronów, aby utworzyć wiązania wielokrotne z atomem centralnym, aby tam, gdzie to możliwe, uzyskać oktety. Chociaż cząsteczka o nieparzystych elektronach nie może mieć oktetu dla każdego atomu, każdy atom powinien otrzymywać elektrony jak najbliżej oktetu. W tym przypadku azot ma wokół siebie tylko pięć elektronów. Aby zbliżyć się do oktetu azotu, jedna z wolnych par tlenu jest wykorzystywana do utworzenia podwójnego wiązania NO. (Kolejnej wolnej pary elektronów nie można pobrać z tlenu i utworzyć wiązanie potrójne, ponieważ azot miałby wówczas dziewięć elektronów).

	<img alt="Figure2" src="https://www.jove.com/files/core/CDNSource/lessons/1066/1066Image_2.svg" style="height: 40px;">


</li>


</ol>

<h4>Cząsteczki z niedoborem elektronów</h4>

Niektóre cząsteczki zawierają jednak atomy centralne, które nie mają wypełnionej powłoki walencyjnej. Generalnie są to cząsteczki z atomami centralnymi z grup 2 i 13, i atomami zewnętrznymi (wodów lub inne atomy), które nie tworzą wiązań wielokrotnych. Na przykład w strukturach Lewisa diwodorku berylu, BeH2 i trifluorku boru, BF3, każdy z atomów berylu i boru ma odpowiednio tylko cztery i sześć elektronów. Możliwe jest narysowanie struktury z podwójnym wiązaniem pomiędzy atomem boru i atomem fluoru w BF3, spełniając regułę oktetu, ale dowody eksperymentalne wskazują, że długości wiązań są bliższe oczekiwanym dla pojedynczych wiązań B – F. Sugeruje to, że najlepsza struktura Lewisa zawiera trzy wiązania pojedyncze B – F i bor z niedoborem elektronów. Reaktywność związku jest również zgodna z borem z niedoborem elektronów. Jednakże wiązania B – F są nieco krótsze niż faktycznie oczekuje się dla pojedynczych wiązań B – F, co wskazuje, że w rzeczywistej cząsteczce występuje charakter wiązania podwójnego.

<img alt="Figure3" src="https://www.jove.com/files/core/CDNSource/lessons/1066/1066Image_3.svg" style="height: 80px;">


Atom taki jak atom boru w BF3, który nie ma ośmiu elektronów, jest bardzo reaktywny. Łatwo łączy się z cząsteczką zawierającą atom z wolną parą elektronów. Na przykład NH3 reaguje z BF3, ponieważ samotna para azotu może być współdzielona z atomem boru:

<img alt="Figure4" src="https://www.jove.com/files/core/CDNSource/lessons/1066/1066Image_4.svg" style="height: 80px;">


Cząsteczki hiperwalentne

Pierwiastki z drugiego okresu układu okresowego (n = 2) mogą pomieścić tylko osiem elektronów na swoich orbitali powłoki walencyjnej, ponieważ mają tylko cztery orbitale walencyjne (jeden orbital 2s i trzy orbitale 2p). Pierwiastki w trzecim i wyższych okresach (n ≥ 3) mają więcej niż cztery orbitale walencyjne i mogą dzielić więcej niż cztery pary elektronów z innymi atomami, ponieważ mają puste orbitale d na tej samej powłoce. Cząsteczki utworzone z tych pierwiastków są czasami nazywane cząsteczkami hiperwalentnymi, takimi jak PCl5 i SF6. W PCl5 centralny atom, fosfor, ma pięć par elektronów. W SF6 siarka ma sześć par elektronów.

<img alt="Figure5" src="https://www.jove.com/files/core/CDNSource/lessons/1066/1066Image_5.svg" style="height: 110px;">


W niektórych cząsteczkach hiperwalentnych, takich jak IF5 i XeF4, niektóre elektrony w zewnętrznej powłoce atomu centralnego stanowią samotne pary:

<img alt="Figure6" src="https://www.jove.com/files/core/CDNSource/lessons/1066/1066image_6.svg" style="height: 110px;">

W strukturach Lewisa tych cząsteczek pozostają elektrony pozostałe po wypełnieniu powłok walencyjnych zewnętrznych atomów ośmioma elektronami. Te dodatkowe elektrony muszą być przypisane do atomu centralnego.

Wyjątki od reguły oktetu

Wiele cząsteczek kowalencyjnych ma centralne atomy, które nie mają ośmiu elektronów w swoich strukturach Lewisa. Cząsteczki te dzielą się na trzy kategorie:

...

Education

JoVE Core

Chemia

Jeden

Chemical Bonding: Basic Concepts

Wiązania chemiczne: Podstawowe pojęcia

KEY TERMS AND CONCEPTS

Kluczowe terminy i pojęcia

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, and it forms in internal combustion engines when oxygen and nitrogen react at high temperature. To draw the Lewis structure for NO, first count the valence electrons. Nitrogen contributes 5 and oxygen contributes 6, for a total of 11.

The odd electron means not every atom can reach an octet. A single N–O bond can be drawn first, and the remaining electrons are placed as lone pairs on the terminal atom. In NO, oxygen is the more electronegative atom, so it gets a full valence shell first. Because no electrons remain for a central atom, the next step is to rearrange electrons to make multiple bonds where possible. In NO, one lone pair from oxygen forms an N=O double bond, which brings nitrogen closer to an octet. A triple bond is not used, because that would give nitrogen nine electrons.

Electron-deficient molecules are another exception. In these compounds, the central atom has fewer electrons than needed for a noble-gas configuration. This often happens with central atoms from groups 2 and 13, especially when the outer atoms are hydrogen or other atoms that do not form multiple bonds. Beryllium dihydride, BeH2, and boron trifluoride, BF3, are examples. In their Lewis structures, beryllium has only four electrons and boron has only six.

BF3 can be drawn with a boron-fluorine double bond, but experimental bond lengths are closer to B–F single bonds. That evidence supports the best Lewis structure as three B–F single bonds with electron-deficient boron. The compound is also very reactive for the same reason. Boron in BF3 readily combines with a molecule that has a lone pair, such as NH3, because the lone pair on nitrogen can be shared with boron.

Hypervalent molecules are the third group. These molecules have a central atom with more than eight electrons in its valence shell. Second-period elements, with only four valence orbitals, can hold only eight electrons. Third-period and higher elements can share more than four pairs of electrons because they have empty d orbitals in the same shell. Examples include PCl5 and SF6, where phosphorus shares five pairs of electrons and sulfur shares six pairs.

Some hypervalent molecules, such as IF5 and XeF4, also have lone pairs on the central atom. In these Lewis structures, electrons remain after the outer atoms have filled octets, and those extra electrons are placed on the central atom. Together, odd-electron, electron-deficient, and hypervalent molecules show the main ways Lewis structures can depart from the octet rule.

Octet Rule Exceptions in Lewis Structures

octet-rule-exceptions-in-lewis-structures

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

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 energies and bond lengths

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...

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 why some Lewis structures break the octet rule with odd-electron, electron-deficient, and hypervalent molecules like NO, BF3, and SF6.

Octet Rule Exceptions in Lewis Structures | 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.

ohios-learning-standards-and-model-curriculum

Ohio’s Learning Standards and Model Curriculum

Physical Science

Study Of Matter

Bonding and compounds

Structure And Properties Of Matter

Chemical bonding

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.3

CH.3.d

Physics

PH.9

PH.9.e

new-york-state-science-learning-standards

New York State Science Learning Standards

HS. Structure and Properties of Matter

HS-PS1-3.

HS. Chemical Reactions

HS-PS1-2.

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 ...

Electronegativity

Whether a bond is nonpolar or polar covalent is determined by a property of the bonding atoms called electronegativity.&#xA0;

Electronegativity values of ...

Bond Polarity, Dipole Moment, and Percent Ionic Character

Bond Polarity

The absolute value of the difference in electronegativity (&Delta;EN) of two bonded atoms provides a rough measure of the polarity to be ...

Lewis Structures of Molecular Compounds and Polyatomic Ions

To draw Lewis structures for complicated molecules and molecular ions, it is helpful to follow a step-by-step procedure as outlined:


	Determine the ...

Resonance

The Lewis structure of a nitrite anion (NO2&minus;) may actually be drawn in two different ways, distinguished by the locations of the N-O and N=O ...

Formal Charges

In some cases, there are seemingly more than one valid Lewis structures for molecules and polyatomic ions. The concept of formal charges can be used to ...

Many covalent molecules have central atoms that do not have eight electrons in their Lewis structures. These molecules fall into three categories:


	...

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, ...

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

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

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 ...

Construct an explanation to summarize the influences intermolecular forces have on the properties of chemical compounds.

PS.Chm.3.2

Analyze the structure of atoms and isotopes.

Octet Rule Exceptions in Lewis Structures

Description

Transcript

Tags

More videos by Standards