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How do I use Lewis structures to determine the oxidation numbers of #"N"# in compounds like #"N"_2"O", "NO", "N"_2"O"_3, "N"_2"O"_4#, and #"N"_2"O"_5#?

assigning oxidation numbers based on lewis dot structures

Explanation:

  • Lone pair electrons (LPs) belong entirely to the atom on which they reside.
  • Shared electrons (bonding pair electrons or BEs) between identical atoms are shared equally.
  • Shared electrons between different atoms belong entirely to the more electronegative atom.
  • Oxidation number (ON) is the difference between the valence electrons in the isolated atom (VE) and the valence electrons in the bound atom (LP + BE).
#color(blue)(bar(ul(|color(white)(a/a) ON = VE - LP - BE color(white)(a/a)|)))" "#

I will calculate the oxidation numbers for only one Lewis structure of each oxide.

An isolated #"N"# atom has five valence electrons ( #VE = 5# ).

Dinitrogen monoxide

The Lewis structure of #"N"_2"O"# is

(a) The left hand #"N"# atom

#LP = 2; BE = 3#

#ON = VE - LP - BE = 5 - 2 - 3 = 0#

(b) The central #"N"# atom

#LP = 0; BE = 3#

#ON = 5 - 0 - 3 = +2#

Each #"N"# atom has a different formal charge.

The average oxidation number on #"N" = (0 + 2)/2 = +1/2#

Nitrogen monoxide

The Lewis structure of #"NO"# is

NO

#LP = 3; BE = 0#

#ON = VE - LP - BE = 5 - 3 - 0 = +2#

Dinitrogen trioxide

The Lewis structure of #"N"_2"O"_3# is

N2O3

#LP = 0; BE = 1#

#ON = 5 - 0 - 1 = +4#

#LP = 2; BE = 1#

#ON = 5 - 2 - 1 = +2#

The average formal charge on #"N" = (+4 + 2)/2 = +6/2 = +3#

Dinitrogen Tetroxide

The Lewis structure of #"N"_2"O"_4# is

(a) The left-hand #"N"# atom

(b) The right-hand #"N"# atom

Each #"N"# atom has the same formal charge.

The average oxidation number on #"N" = (+4 + 4)/2 = +8/2 = +4#

Dinitrogen pentoxide

The Lewis structure of #"N"_2"O"_5# is

N2O5

#LP = 0; BE = 0#

#ON = 5 - 0 - 0 = +5#

The average oxidation number on #"N" = (+5 + 5)/2 = +10/2 = +5#

Related questions

  • What are lewis dot structures used for?
  • What is the lewis structure for #SO_2#?
  • How do you draw the lewis structure for ions?
  • How do you draw the Lewis structure for ionic compounds?
  • What are some examples of Lewis structures?
  • What is an example of a Lewis structures practice problem?
  • What are some common mistakes students make with Lewis structures?
  • What are some common mistakes students make when drawing Lewis structures?
  • How can I draw Lewis dot structures for ionic compounds?
  • How can I draw a Lewis structure of a compound?

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assigning oxidation numbers based on lewis dot structures

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29 7.2 Lewis Dot Structures

Learning objective.

By the end of this section, you will be able to:

  • Draw a Lewis electron dot diagram for any atom or a monatomic ion with an atomic number of less than 20.

In almost all cases, chemical bonds are formed by interactions of valence electrons in atoms. To facilitate our understanding of how valence electrons interact, a simple way of representing those valence electrons would be useful.

A Lewis dot structure is a representation of the valence electrons of an atom that uses dots around the symbol of the element. The number of dots equals the number of valence electrons in the atom. These dots are arranged to the right and left and above and below the symbol, with no more than two dots on a side. (It does not matter what order the positions are used.) For example, the Lewis electron dot diagram for hydrogen is simply

Hydrogen

Because the side is not important, the Lewis electron dot diagram could also be drawn as follows:

Hydrogen-Sides

The electron dot diagram for helium, with two valence electrons, is as follows:

Helium

By putting the two electrons together on the same side, we emphasize the fact that these two electrons are both in the first shell. The next atom, lithium, has three electrons total but only one electron in its valence shell. Its electron dot diagram resembles that of hydrogen, except the symbol for lithium is used:

Lithium

Carbon has four valence electrons. We draw the dots for the electrons on different sides. As such, the electron dot diagram for carbon is as follows:

assigning oxidation numbers based on lewis dot structures

With nitrogen, we distribute the five valence electrons around the atom, using four positions and not pairing up any electrons until necessary:

Nitrogen

Oxygen, with a total of six valence electrons, ends up with two unpaired electrons and two sets of paired electrons.

Oxygen

Fluorine and neon have seven and eight dots, respectively:

Fluoride-Neon

With the next element, sodium, the process starts over with a single electron because sodium has a single electron in its highest-numbered shell, the n = 3 shell.

What is the Lewis electron dot diagram for each element?      a) phosphorous         b) argon

Phosphorus-Argon

Elements in the same column of the periodic table have similar Lewis electron dot diagrams because they have the same valence shell electron configuration. Thus the electron dot diagrams for the first column of elements are as follows:

First-Column

Lewis dot structures are particularly useful for describing covalent bonding in compounds. They are somewhat but less helpful for describing ions. As a reminder, covalent bonding occurs when nonmetal elements form bonds, so drawing covalent compounds with Lewis structures only requires drawing atoms of these nonmetal elements. Unless you want to delve deeper, there is no need to learn about drawing Lewis structures for metals and ions.

Key Concepts and Summary

Lewis dot structures represent atoms with their atomic symbol surrounded by valence electrons, which are represented as dots.  This type of symbolic representation can help describe compound formation, especially for covalent compounds.

Review-Reflect, Extend

Review-reflect.

1. What column of the periodic table has Lewis electron dot diagrams with two electrons?

2. Draw the Lewis electron dot diagram for silicon.

We know that water can be broken down into elemental hydrogen and elemental oxygen, and that the ratio of these gases that form is 2:1. Since both of these gases are diatomic (H 2 and O 2 ), water is made of H and O atoms in a 2:1 ratio as well. The observable properties of water indicate that these atoms are bonded together by covalent, rather than ionic bonds.

Draw Lewis dot structures for two hydrogen atoms and one oxygen atom. Attempt to arrange these three atoms so that they are sharing electrons. A finished “correct” structure should have every atom in the structure, once the sharing arrangements are made, with an electron arrangement that could be seen as “complete” or a “full shell.”

Answers to Review-Reflect

1. the second column of the periodic table

Silicone

Introduction to Chemistry Copyright © 2020 by Carol Higginbotham is licensed under a Creative Commons Attribution-NonCommercial-ShareAlike 4.0 International License , except where otherwise noted.

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  • 7.3 Lewis Symbols and Structures
  • Introduction
  • 1.1 Chemistry in Context
  • 1.2 Phases and Classification of Matter
  • 1.3 Physical and Chemical Properties
  • 1.4 Measurements
  • 1.5 Measurement Uncertainty, Accuracy, and Precision
  • 1.6 Mathematical Treatment of Measurement Results
  • Key Equations
  • 2.1 Early Ideas in Atomic Theory
  • 2.2 Evolution of Atomic Theory
  • 2.3 Atomic Structure and Symbolism
  • 2.4 Chemical Formulas
  • 2.5 The Periodic Table
  • 2.6 Ionic and Molecular Compounds
  • 2.7 Chemical Nomenclature
  • 3.1 Formula Mass and the Mole Concept
  • 3.2 Determining Empirical and Molecular Formulas
  • 3.3 Molarity
  • 3.4 Other Units for Solution Concentrations
  • 4.1 Writing and Balancing Chemical Equations
  • 4.2 Classifying Chemical Reactions
  • 4.3 Reaction Stoichiometry
  • 4.4 Reaction Yields
  • 4.5 Quantitative Chemical Analysis
  • 5.1 Energy Basics
  • 5.2 Calorimetry
  • 5.3 Enthalpy
  • 6.1 Electromagnetic Energy
  • 6.2 The Bohr Model
  • 6.3 Development of Quantum Theory
  • 6.4 Electronic Structure of Atoms (Electron Configurations)
  • 6.5 Periodic Variations in Element Properties
  • 7.1 Ionic Bonding
  • 7.2 Covalent Bonding
  • 7.4 Formal Charges and Resonance
  • 7.5 Strengths of Ionic and Covalent Bonds
  • 7.6 Molecular Structure and Polarity
  • 8.1 Valence Bond Theory
  • 8.2 Hybrid Atomic Orbitals
  • 8.3 Multiple Bonds
  • 8.4 Molecular Orbital Theory
  • 9.1 Gas Pressure
  • 9.2 Relating Pressure, Volume, Amount, and Temperature: The Ideal Gas Law
  • 9.3 Stoichiometry of Gaseous Substances, Mixtures, and Reactions
  • 9.4 Effusion and Diffusion of Gases
  • 9.5 The Kinetic-Molecular Theory
  • 9.6 Non-Ideal Gas Behavior
  • 10.1 Intermolecular Forces
  • 10.2 Properties of Liquids
  • 10.3 Phase Transitions
  • 10.4 Phase Diagrams
  • 10.5 The Solid State of Matter
  • 10.6 Lattice Structures in Crystalline Solids
  • 11.1 The Dissolution Process
  • 11.2 Electrolytes
  • 11.3 Solubility
  • 11.4 Colligative Properties
  • 11.5 Colloids
  • 12.1 Chemical Reaction Rates
  • 12.2 Factors Affecting Reaction Rates
  • 12.3 Rate Laws
  • 12.4 Integrated Rate Laws
  • 12.5 Collision Theory
  • 12.6 Reaction Mechanisms
  • 12.7 Catalysis
  • 13.1 Chemical Equilibria
  • 13.2 Equilibrium Constants
  • 13.3 Shifting Equilibria: Le Châtelier’s Principle
  • 13.4 Equilibrium Calculations
  • 14.1 Brønsted-Lowry Acids and Bases
  • 14.2 pH and pOH
  • 14.3 Relative Strengths of Acids and Bases
  • 14.4 Hydrolysis of Salts
  • 14.5 Polyprotic Acids
  • 14.6 Buffers
  • 14.7 Acid-Base Titrations
  • 15.1 Precipitation and Dissolution
  • 15.2 Lewis Acids and Bases
  • 15.3 Coupled Equilibria
  • 16.1 Spontaneity
  • 16.2 Entropy
  • 16.3 The Second and Third Laws of Thermodynamics
  • 16.4 Free Energy
  • 17.1 Review of Redox Chemistry
  • 17.2 Galvanic Cells
  • 17.3 Electrode and Cell Potentials
  • 17.4 Potential, Free Energy, and Equilibrium
  • 17.5 Batteries and Fuel Cells
  • 17.6 Corrosion
  • 17.7 Electrolysis
  • 18.1 Periodicity
  • 18.2 Occurrence and Preparation of the Representative Metals
  • 18.3 Structure and General Properties of the Metalloids
  • 18.4 Structure and General Properties of the Nonmetals
  • 18.5 Occurrence, Preparation, and Compounds of Hydrogen
  • 18.6 Occurrence, Preparation, and Properties of Carbonates
  • 18.7 Occurrence, Preparation, and Properties of Nitrogen
  • 18.8 Occurrence, Preparation, and Properties of Phosphorus
  • 18.9 Occurrence, Preparation, and Compounds of Oxygen
  • 18.10 Occurrence, Preparation, and Properties of Sulfur
  • 18.11 Occurrence, Preparation, and Properties of Halogens
  • 18.12 Occurrence, Preparation, and Properties of the Noble Gases
  • 19.1 Occurrence, Preparation, and Properties of Transition Metals and Their Compounds
  • 19.2 Coordination Chemistry of Transition Metals
  • 19.3 Spectroscopic and Magnetic Properties of Coordination Compounds
  • 20.1 Hydrocarbons
  • 20.2 Alcohols and Ethers
  • 20.3 Aldehydes, Ketones, Carboxylic Acids, and Esters
  • 20.4 Amines and Amides
  • 21.1 Nuclear Structure and Stability
  • 21.2 Nuclear Equations
  • 21.3 Radioactive Decay
  • 21.4 Transmutation and Nuclear Energy
  • 21.5 Uses of Radioisotopes
  • 21.6 Biological Effects of Radiation
  • A | The Periodic Table
  • B | Essential Mathematics
  • C | Units and Conversion Factors
  • D | Fundamental Physical Constants
  • E | Water Properties
  • F | Composition of Commercial Acids and Bases
  • G | Standard Thermodynamic Properties for Selected Substances
  • H | Ionization Constants of Weak Acids
  • I | Ionization Constants of Weak Bases
  • J | Solubility Products
  • K | Formation Constants for Complex Ions
  • L | Standard Electrode (Half-Cell) Potentials
  • M | Half-Lives for Several Radioactive Isotopes

Learning Objectives

By the end of this section, you will be able to:

  • Write Lewis symbols for neutral atoms and ions
  • Draw Lewis structures depicting the bonding in simple molecules

Thus far in this chapter, we have discussed the various types of bonds that form between atoms and/or ions. In all cases, these bonds involve the sharing or transfer of valence shell electrons between atoms. In this section, we will explore the typical method for depicting valence shell electrons and chemical bonds, namely Lewis symbols and Lewis structures.

Lewis Symbols

We use Lewis symbols to describe valence electron configurations of atoms and monatomic ions. A Lewis symbol consists of an elemental symbol surrounded by one dot for each of its valence electrons:

Figure 7.9 shows the Lewis symbols for the elements of the third period of the periodic table.

Lewis symbols can also be used to illustrate the formation of cations from atoms, as shown here for sodium and calcium:

Likewise, they can be used to show the formation of anions from atoms, as shown here for chlorine and sulfur:

Figure 7.10 demonstrates the use of Lewis symbols to show the transfer of electrons during the formation of ionic compounds.

Lewis Structures

We also use Lewis symbols to indicate the formation of covalent bonds, which are shown in Lewis structures , drawings that describe the bonding in molecules and polyatomic ions. For example, when two chlorine atoms form a chlorine molecule, they share one pair of electrons:

The Lewis structure indicates that each Cl atom has three pairs of electrons that are not used in bonding (called lone pairs ) and one shared pair of electrons (written between the atoms). A dash (or line) is sometimes used to indicate a shared pair of electrons:

A single shared pair of electrons is called a single bond . Each Cl atom interacts with eight valence electrons: the six in the lone pairs and the two in the single bond.

The Octet Rule

The other halogen molecules (F 2 , Br 2 , I 2 , and At 2 ) form bonds like those in the chlorine molecule: one single bond between atoms and three lone pairs of electrons per atom. This allows each halogen atom to have a noble gas electron configuration. The tendency of main group atoms to form enough bonds to obtain eight valence electrons is known as the octet rule .

The number of bonds that an atom can form can often be predicted from the number of electrons needed to reach an octet (eight valence electrons); this is especially true of the nonmetals of the second period of the periodic table (C, N, O, and F). For example, each atom of a group 14 element has four electrons in its outermost shell and therefore requires four more electrons to reach an octet. These four electrons can be gained by forming four covalent bonds, as illustrated here for carbon in CCl 4 (carbon tetrachloride) and silicon in SiH 4 (silane). Because hydrogen only needs two electrons to fill its valence shell, it is an exception to the octet rule. The transition elements and inner transition elements also do not follow the octet rule:

Group 15 elements such as nitrogen have five valence electrons in the atomic Lewis symbol: one lone pair and three unpaired electrons. To obtain an octet, these atoms form three covalent bonds, as in NH 3 (ammonia). Oxygen and other atoms in group 16 obtain an octet by forming two covalent bonds:

Double and Triple Bonds

As previously mentioned, when a pair of atoms shares one pair of electrons, we call this a single bond. However, a pair of atoms may need to share more than one pair of electrons in order to achieve the requisite octet. A double bond forms when two pairs of electrons are shared between a pair of atoms, as between the carbon and oxygen atoms in CH 2 O (formaldehyde) and between the two carbon atoms in C 2 H 4 (ethylene):

A triple bond forms when three electron pairs are shared by a pair of atoms, as in carbon monoxide (CO) and the cyanide ion (CN – ):

Writing Lewis Structures with the Octet Rule

For very simple molecules and molecular ions, we can write the Lewis structures by merely pairing up the unpaired electrons on the constituent atoms. See these examples:

For more complicated molecules and molecular ions, it is helpful to follow the step-by-step procedure outlined here:

  • Determine the total number of valence (outer shell) electrons. For cations, subtract one electron for each positive charge. For anions, add one electron for each negative charge.
  • Draw a skeleton structure of the molecule or ion, arranging the atoms around a central atom. (Generally, the least electronegative element should be placed in the center.) Connect each atom to the central atom with a single bond (one electron pair).
  • Distribute the remaining electrons as lone pairs on the terminal atoms (except hydrogen), completing an octet around each atom.
  • Place all remaining electrons on the central atom.
  • Rearrange the electrons of the outer atoms to make multiple bonds with the central atom in order to obtain octets wherever possible.

Let us determine the Lewis structures of SiH 4 , CHO 2 − , CHO 2 − , NO + , and OF 2 as examples in following this procedure:

  • For a molecule, we add the number of valence electrons on each atom in the molecule: SiH 4 Si: 4 valence electrons/atom × 1 atom = 4 + H: 1 valence electron/atom × 4 atoms = 4 ¯ = 8 valence electrons SiH 4 Si: 4 valence electrons/atom × 1 atom = 4 + H: 1 valence electron/atom × 4 atoms = 4 ¯ = 8 valence electrons
  • For a negative ion , such as CHO 2 − , CHO 2 − , we add the number of valence electrons on the atoms to the number of negative charges on the ion (one electron is gained for each single negative charge): CHO 2 − C: 4 valence electrons/atom × 1 atom = 4 H: 1 valence electron/atom × 1 atom = 1 O: 6 valence electrons/atom × 2 atoms = 12 + 1 additional electron = 1 ¯ = 18 valence electrons CHO 2 − C: 4 valence electrons/atom × 1 atom = 4 H: 1 valence electron/atom × 1 atom = 1 O: 6 valence electrons/atom × 2 atoms = 12 + 1 additional electron = 1 ¯ = 18 valence electrons
  • For a positive ion , such as NO + , we add the number of valence electrons on the atoms in the ion and then subtract the number of positive charges on the ion (one electron is lost for each single positive charge) from the total number of valence electrons: NO + N: 5 valence electrons/atom × 1 atom = 5 O: 6 valence electron/atom × 1 atom = 6 + −1 electron (positive charge) = −1 ¯ = 10 valence electrons NO + N: 5 valence electrons/atom × 1 atom = 5 O: 6 valence electron/atom × 1 atom = 6 + −1 electron (positive charge) = −1 ¯ = 10 valence electrons
  • Since OF 2 is a neutral molecule, we simply add the number of valence electrons: OF 2 O: 6 valence electrons/atom × 1 atom = 6 + F: 7 valence electrons/atom × 2 atoms = 14 ¯ = 20 valence electrons OF 2 O: 6 valence electrons/atom × 1 atom = 6 + F: 7 valence electrons/atom × 2 atoms = 14 ¯ = 20 valence electrons
  • For SiH 4 , CHO 2 − , CHO 2 − , and NO + , there are no remaining electrons; we already placed all of the electrons determined in Step 1.
  • SiH 4 : Si already has an octet, so nothing needs to be done.
  • In OF 2 , each atom has an octet as drawn, so nothing changes.

Example 7.4

Writing lewis structures.

  • Step 1. Calculate the number of valence electrons. HCN: (1 × × 1) + (4 × × 1) + (5 × × 1) = 10 H 3 CCH 3 : (1 × × 3) + (2 × × 4) + (1 × × 3) = 14 HCCH: (1 × × 1) + (2 × × 4) + (1 × × 1) = 10 NH 3 : (5 × × 1) + (3 × × 1) = 8

Check Your Learning

How sciences interconnect, fullerene chemistry.

Carbon, in various forms and compounds, has been known since prehistoric times, . Soot has been used as a pigment (often called carbon black) for thousands of years. Charcoal, high in carbon content, has likewise been critical to human development. Carbon is the key additive to iron in the steelmaking process, and diamonds have a unique place in both culture and industry. With all this usage came significant study, particularly with the emergence of organic chemistry. And even with all the known forms and functions of the element, scientists began to uncover the potential for even more varied and extensive carbon structures.

As early as the 1960s, chemists began to observe complex carbon structures, but they had little evidence to support their concepts, or their work did not make it into the mainstream. Eiji Osawa predicted a spherical form based on observations of a similar structure, but his work was not widely known outside Japan. In a similar manner, the most comprehensive advance was likely computational chemist Elena Galpern's, who in 1973 predicted a highly stable, 60-carbon molecule; her work was also isolated to her native Russia. Still later, Harold Kroto, working with Canadian radio astronomers, sought to uncover the nature of long carbon chains that had been discovered in interstellar space.

Kroto sought to use a machine developed by Richard Smalley's team at Rice University to learn more about these structures. Together with Robert Curl, who had introduced them, and three graduate students—James Heath, Sean O’Brien, and Yuan Liu—they performed an intensive series of experiments that led to a major discovery.

In 1996, the Nobel Prize in Chemistry was awarded to Richard Smalley ( Figure 7.11 ), Robert Curl, and Harold Kroto for their work in discovering a new form of carbon, the C 60 buckminsterfullerene molecule ( Figure 7.1 ). An entire class of compounds, including spheres and tubes of various shapes, were discovered based on C 60. This type of molecule, called a fullerene, shows promise in a variety of applications. Because of their size and shape, fullerenes can encapsulate other molecules, so they have shown potential in various applications from hydrogen storage to targeted drug delivery systems. They also possess unique electronic and optical properties that have been put to good use in solar powered devices and chemical sensors.

Exceptions to the Octet Rule

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

  • Odd-electron molecules have an odd number of valence electrons, and therefore have an unpaired electron.
  • Electron-deficient molecules have a central atom that has fewer electrons than needed for a noble gas configuration.
  • Hypervalent molecules have a central atom that has more electrons than needed for a noble gas configuration.

Odd-electron Molecules

We call molecules that contain an odd number of electrons free radicals . Nitric oxide, NO, is an example of an odd-electron molecule; it is produced in internal combustion engines when oxygen and nitrogen react at high temperatures.

To draw the Lewis structure for an odd-electron molecule like NO, we follow the same five steps we would for other molecules, but with a few minor changes:

  • Determine the total number of valence (outer shell) electrons . The sum of the valence electrons is 5 (from N) + 6 (from O) = 11. The odd number immediately tells us that we have a free radical, so we know that not every atom can have eight electrons in its valence shell.
  • Draw a skeleton structure of the molecule . We can easily draw a skeleton with an N–O single bond: N–O
  • Place all remaining electrons on the central atom . Since there are no remaining electrons, this step does not apply.

Electron-deficient Molecules

We will also encounter a few molecules that contain central atoms that do not have a filled valence shell. Generally, these are molecules with central atoms from groups 2 and 13, outer atoms that are hydrogen, or other atoms that do not form multiple bonds. For example, molecules of beryllium dihydride, BeH 2 , and boron trifluoride, BF 3 , contain central atoms with only two and three valence electrons, respectively. With only H as outer atoms, there is just one possibility for the Lewis structure of BeH 2 , and it does not satisfy the octet rule for the central Be atom (see as follows). It is possible to draw a structure with a double bond between a boron atom and a fluorine atom in BF 3 , satisfying the octet rule, but experimental evidence indicates the bond lengths are closer to that expected for B–F single bonds. This suggests the best Lewis structure has three B–F single bonds and an electron deficient boron. The reactivity of the compound is also consistent with an electron deficient boron. However, the B–F bonds are slightly shorter than what is actually expected for B–F single bonds, indicating that some double bond character is found in the actual molecule.

An atom like the boron atom in BF 3 , which does not have eight electrons, is very reactive. It readily combines with a molecule containing an atom with a lone pair of electrons. For example, NH 3 reacts with BF 3 because the lone pair on nitrogen can be shared with the boron atom:

Hypervalent Molecules

Elements in the second period of the periodic table ( n = 2) can accommodate only eight electrons in their valence shell orbitals because they have only four valence orbitals (one 2 s and three 2 p orbitals). Elements in the third and higher periods ( n ≥ 3) have more than four valence orbitals and can share more than four pairs of electrons with other atoms because they have empty d orbitals in the same shell. Molecules formed from these elements are sometimes called hypervalent molecules . Figure 7.12 shows the Lewis structures for two hypervalent molecules, PCl 5 and SF 6.

In some hypervalent molecules, such as IF 5 and XeF 4 , some of the electrons in the outer shell of the central atom are lone pairs:

When we write the Lewis structures for these molecules, we find that we have electrons left over after filling the valence shells of the outer atoms with eight electrons. These additional electrons must be assigned to the central atom.

Example 7.5

Writing lewis structures: octet rule violations.

  • Step 1. Calculate the number of valence electrons: XeF 2 : 8 + (2 × × 7) = 22 XeF 6 : 8 + (6 × × 7) = 50

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CHEM101: General Chemistry I

Formal charge and dot structures.

Sometimes we can draw more than one Lewis structure for a given molecule or polyatomic ion that are not equivalent. In this case, we use a concept called formal charge to determine, which Lewis structure is best. Formal charge is not actually a charge; rather, it is just a system to keep track of electrons in a given Lewis structure.

Watch these two videos, which teach you the rules for assigning formal charge to atoms in a Lewis structure, and show examples of using formal charge to determine the best possible Lewis structure for a given molecule.

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4.2: Lewis Structures (Problems)

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PROBLEM \(\PageIndex{1}\)

Write Lewis structures for the following: (please note, none of the solutions are using the expanded octet rule or formal charges)

Screen Shot 2018-12-31 at 2.21.44 PM.png

PROBLEM \(\PageIndex{2}\)

Write Lewis structures for the following: (please note, none of the solutions are using the expanded octet rule or formal charges)

  • \(\ce{NH4+}\)
  • \(\ce{BF4-}\)
  • \(\ce{C2^2+}\)

PROBLEM \(\PageIndex{3}\)

Write Lewis structures for: (please note, none of the solutions are using the expanded octet rule or formal charges)

a. \(\ce{PO4^3-}\) c. \(\ce{SO3^2-}\) d. HONO

Screen Shot 2018-12-31 at 2.22.34 PM.png

PROBLEM \(\PageIndex{4}\)

Methanol, H 3 COH, is used as the fuel in some race cars. Ethanol, C 2 H 5 OH, is used extensively as motor fuel in Brazil. Both methanol and ethanol produce CO 2 and H 2 O when they burn. Write the chemical equations for these combustion reactions using Lewis structures instead of chemical formulas.

PROBLEM \(\PageIndex{5}\)

Many planets in our solar system contain organic chemicals including methane (CH 4 ) and traces of ethylene (C 2 H 4 ), ethane (C 2 H 6 ), propyne (H 3 CCCH), and diacetylene (HCCCCH). Write the Lewis structures for each of these molecules. (diacetylene may be a little tricky!)

Screen Shot 2018-12-31 at 2.22.47 PM.png

PROBLEM \(\PageIndex{6}\)

Carbon tetrachloride was formerly used in fire extinguishers for electrical fires. It is no longer used for this purpose because of the formation of the toxic gas phosgene, Cl 2 CO. Write the Lewis structures for carbon tetrachloride and phosgene.

PROBLEM \(\PageIndex{7}\)

The arrangement of atoms in several biologically important molecules is given here. Complete the Lewis structures of these molecules by adding multiple bonds and lone pairs. Do not add any more atoms.

a. the amino acid serine:

c. pyruvic acid:

e. carbonic acid:

PROBLEM \(\PageIndex{8}\)

How are single, double, and triple bonds similar? How do they differ?

Each bond includes a sharing of electrons between atoms. Two electrons are shared in a single bond; four electrons are shared in a double bond; and six electrons are shared in a triple bond.

Contributors

Paul Flowers (University of North Carolina - Pembroke), Klaus Theopold (University of Delaware) and Richard Langley (Stephen F. Austin State University) with contributing authors.  Textbook content produced by OpenStax College is licensed under a Creative Commons Attribution License 4.0 license. Download for free at http://cnx.org/contents/[email protected] ).

  • Adelaide Clark, Oregon Institute of Technology

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  1. How to Determine the Oxidation Number from Lewis Structure Examples

    assigning oxidation numbers based on lewis dot structures

  2. How to Assign Oxidation Numbers

    assigning oxidation numbers based on lewis dot structures

  3. Assigning Oxidation Numbers

    assigning oxidation numbers based on lewis dot structures

  4. Assigning oxidation numbers

    assigning oxidation numbers based on lewis dot structures

  5. Rules to assign oxidation numbers

    assigning oxidation numbers based on lewis dot structures

  6. Oxidation Numbers

    assigning oxidation numbers based on lewis dot structures

VIDEO

  1. Oxidation Reduction Part 2 based on cee name institute

  2. Intro to REDOX reactions: Assigning Oxidation Numbers

  3. Chemistry for NEET/ JEE -Class-11- Equilibrium-Conjugate acid & based+Lewis acid & base Concept-L-18

  4. Assigning Oxidation numbers

  5. Assigning Oxidation Numbers (Chemistry 30 Lesson 13-7)

  6. Assigning Oxidation Numbers

COMMENTS

  1. Oxidation Numbers and Lewis Structures

    A method to assign oxidation numbers to evey atom in a molecule based on the Lewis dot structure of the molecule.

  2. How to Determine the Oxidation Number from Lewis Structure ...

    🎯 Want to ace chemistry? Access the best chemistry resource at http://www.conquerchemistry.com/masterclass📗 Need help with chemistry? Download 12 Secrets t...

  3. How do I use Lewis structures to determine the oxidation numbers of

    1 Answer Ernest Z. Oct 14, 2016 Warning! Long answer. You count the valence electrons around N according to a set of rules and then assign the oxidation number. Explanation: The Rules Lone pair electrons (LPs) belong entirely to the atom on which they reside.

  4. 4.3: Formal Charge and Oxidation State

    Subtract this number from the number of valence electrons for the neutral atom: I: 7 - 8 = -1. Cl: 7 - 7 = 0. The sum of the formal charges of all the atoms equals -1, which is identical to the charge of the ion (-1). Exercise 4.3.1 4.3. 1. Calculate the formal charge for each atom in the carbon monoxide molecule:

  5. PDF Oxidation Numbers in Lewis Structures

    First step: Draw a Lewis Structure; Oxalate Ion = C2O4 -2 Count valence electrons. (2x4 for carbon) + (4x6 for oxygen) + (2 extra for the "-2" charge) = 34 valence electrons Draw a skeleton with all single bonds Fill octets of peripheral atoms (oxygens in this case) Assigning Electrons

  6. 22.6: Assigning Oxidation Numbers

    In the chlorate ion (ClO−3) ( ClO 3 −), the oxidation number of Cl Cl is +5 + 5, and the oxidation number of O O is −2 − 2. In a neutral atom or molecule, the sum of the oxidation numbers must be 0. In a polyatomic ion, the sum of the oxidation numbers of all the atoms in the ion must be equal to the charge on the ion. Example 22.6.1 22 ...

  7. Formal Charges in Lewis Structures

    To find formal charges in a Lewis structure, for each atom, you should count how many electrons it "owns". Count all of its lone pair electrons, and half of its bonding electrons. The difference between the atom's number of valence electrons and the number it owns is the formal charge. For example, in NH 3, N has 1 lone pair (2 electrons) and 3 ...

  8. 7.2 Lewis Dot Structures

    A Lewis dot structure is a representation of the valence electrons of an atom that uses dots around the symbol of the element. The number of dots equals the number of valence electrons in the atom. These dots are arranged to the right and left and above and below the symbol, with no more than two dots on a side.

  9. Week 12

    There's more than one way to calculate oxidation numbers! Instead of learning a set of rules (and exceptions), you can calculate them directly from knowledge...

  10. PDF Practice Worksheet for Lewis Structures (Mahaffy Ch. 10.1

    e. Assigning Oxidation number using Lewis Structures i. Oxidation number is 0 for atoms in an element. ii. The sum of all oxidation numbers in a molecule or ion must add up to the total charge. iii. In compounds, alkalis (group 1) have oxidation number +1; alkaline earths (group 2) have oxidation number +2. iv. In compounds, fluorine (F) always ...

  11. 7.3 Lewis Symbols and Structures

    We use Lewis symbols to describe valence electron configurations of atoms and monatomic ions. A Lewis symbol consists of an elemental symbol surrounded by one dot for each of its valence electrons: Figure 7.9 shows the Lewis symbols for the elements of the third period of the periodic table. Figure 7.9 Lewis symbols illustrating the number of ...

  12. Lewis structures, formal charge, and oxidation numbers: A more user

    This paper presents a set of rules for writing Lewis structures requiring only the ability to add, subtract, count, and know the number of valence electrons of neutral atoms. KEYWORDS (Audience): High School / Introductory Chemistry KEYWORDS (Domain): Inorganic Chemistry KEYWORDS (Pedagogy): Mnemonics / Rote Learning KEYWORDS (Subject):

  13. PDF Oxidation Numbers: Rules

    7) The oxidation number of Group 1A elements is always +1 and the oxidation number of Group 2A elements is always +2. 8) The oxidation number of oxygen in most compounds is -2. 9) Oxidation numbers for other elements are usually determined by the number of electrons they need to gain or lose to attain the electron configuration of a noble gas.

  14. 9.2: Interpreting Lewis Structures

    Interpreting Lewis Structures. A Lewis structure contains symbols for the elements in a molecule, connected by lines and surrounded by pairs of dots. For example, here is the Lewis structure for water, H 2 O. Each symbol represents the nucleus and the core electrons of the atom. Here, each "H" represents the nucleus of a hydrogen atom, and ...

  15. Assigning Oxidation States: Rules & Lewis Structures

    In summary, there are two main methods for assigning oxidation states in compounds - drawing Lewis structures and following specific rules. However, as a beginner in chemistry, these methods may not always work for all compounds. As you advance in chemistry, you will learn more rules and methods for assigning oxidation numbers.

  16. PDF Practice worksheet for Lewis Structures (Mahaffy, 2e section: 10.1 10.5

    f. Oxidation number i. Oxidation number is 0 for atoms in an element. ii. The sum of all oxidation numbers in a molecule or ion must add up to the total charge. iii. In compounds, alkalis (group 1) have oxidation number +1; alkaline earths (group 2) have oxidation number +2. iv. In compounds, fluorine (F) always has oxidation number -1.

  17. Assign oxidation numbers to each element in the following ...

    Lewis Dot Structures: Sigma & Pi Bonds 14m. Lewis Dot Structures: Ions 15m. Lewis Dot Structures: Exceptions 13m. Lewis Dot Structures: Acids 15m. Resonance Structures 19m. Average Bond Order 4m. ... Assign oxidation numbers to each element in the following compounds. (b) SO3. Verified Solution.

  18. 9.3: Drawing Lewis Structures

    Step 1: Figure out how many electrons the molecule must have, based on the number of valence electrons in each atom. When drawing the structure of an ion, be sure to add/subtract electrons to account for the charge. Step 2: Connect the atoms to each other with single bonds to form a "skeleton structure.".

  19. Assign oxidation numbers to each element in the following ...

    Lewis Dot Structures: Sigma & Pi Bonds 4m. Lewis Dot Structures: Ions 9m. Lewis Dot Structures: Exceptions 6m. Lewis Dot Structures: Acids 3m. Resonance Structures 12m. Average Bond Order 3m. ... Assign oxidation numbers to each element in the following compounds. (e) OsO4. Verified Solution. 6m.

  20. CHEM101: Formal Charge and Dot Structures

    Formal Charge and Dot Structures. Sometimes we can draw more than one Lewis structure for a given molecule or polyatomic ion that are not equivalent. In this case, we use a concept called formal charge to determine, which Lewis structure is best. Formal charge is not actually a charge; rather, it is just a system to keep track of electrons in a ...

  21. Solved Assign oxidation numbers to the following elements ...

    Based on the Lewis Dot Structure and additional doodle below, determine the oxidation numbers of the following atoms: The RED carbon atom has an oxidation number of The GREEN carbon atom has an oxidation number of The PURPLE carbon atom has an oxidation number of All hydrogen atoms have an oxidation number of in this molecule.

  22. Chapter 5.3: Lewis Structures

    2. Each hydrogen atom (group 1) has one valence electron, carbon (group 14) has 4 valence electrons, and oxygen (group 16) has 6 valence electrons, for a total of [ (2) (1) + 4 + 6] = 12 valence electrons. 3. Placing a bonding pair of electrons between each pair of bonded atoms gives the following: Six electrons are used, and 6 are left over.

  23. 4.2: Lewis Structures (Problems)

    PROBLEM 4.2.4 4.2. 4. Methanol, H 3 COH, is used as the fuel in some race cars. Ethanol, C 2 H 5 OH, is used extensively as motor fuel in Brazil. Both methanol and ethanol produce CO 2 and H 2 O when they burn. Write the chemical equations for these combustion reactions using Lewis structures instead of chemical formulas.