H Cation Lewis Structure

Alternate view of lewis dot structure of water: This arrangement of shared electrons between O and H results in the oxygen atom having an octet of electrons, and each H atom having two valence electrons. Multiple bonds can also form between elements when two or three pairs of electrons are shared to produce double or triple bonds, respectively. A complex ion consists of a central atom, typically a transition metal cation, surrounded by ions, or molecules called ligands. These ligands can be neutral molecules like H 2 O or NH 3, or ions such as CN – or OH –. Often, the ligands act as Lewis bases, donating a pair of electrons to the central atom. Trimethylammonium C3H10N+ CID 3782034 - structure, chemical names, physical and chemical properties, classification, patents, literature, biological activities.

Representing Valence Electrons in Lewis Symbols

Lewis symbols use dots to visually represent the valence electrons of an atom.

Learning Objectives

Recall the Lewis structure formalism for representing valance electrons

Key Takeaways

Key Points

• Electrons exist outside of an atom ‘s nucleus and are found in principal energy levels that contain only up to a specific number of electrons.
• The outermost principal energy level that contains electrons is called the valence level and contains valence electrons.
• Lewis symbols are diagrams that show the number of valence electrons of a particular element with dots that represent lone pairs.
• Lewis symbols do not visualize the electrons in the inner principal energy levels.

Key Terms

• principal energy levels: The different levels where electrons can be found and that occur at specific distances from the atom’s nucleus. Each level is associated with a particular energy value that electrons within it have.
• valence level: The outermost principal energy level, which is the level furthest away from the nucleus that still contains electrons.
• valence electrons: The electrons of atoms that participate in the formation of chemical bonds.
• Lewis symbols: Symbols of the elements with their number of valence electrons represented as dots

Lewis symbols (also known as Lewis dot diagrams or electron dot diagrams) are diagrams that represent the valence electrons of an atom. Lewis structures (also known as Lewis dot structures or electron dot structures) are diagrams that represent the valence electrons of atoms within a molecule. These Lewis symbols and Lewis structures help visualize the valence electrons of atoms and molecules, whether they exist as lone pairs or within bonds.

Principal Energy Levels

An atom consists of a positively charged nucleus and negatively charged electrons. The electrostatic attraction between them keeps electrons ‘bound’ to the nucleus so they stay within a certain distance of it. Careful investigations have shown that not all electrons within an atom have the same average position or energy. We say the electrons ‘reside’ in different principal energy levels, and these levels exist at different radii from the nucleus and have rules regarding how many electrons they can accommodate.

Principal energy levels of gold (Au): The figure shows the organization of the electrons around the nucleus of a gold (Au) atom. Notice that the first energy level (closest to the nucleus) can have only two electrons, while more electrons can ‘fit’ within a given level further out. The number of electrons in each level is listed on the upper right corner of the figure. Notice that the outermost level has only one electron.

As an example, a neutral atom of gold (Au) contains 79 protons in its nucleus and 79 electrons. The first principal energy level, which is the one closest to the nucleus, can hold a maximum of two electrons. The second principal energy level can have 8, the third can have 18, and so on, until all 79 electrons have been distributed.

The outermost principal energy level is of great interest in chemistry because the electrons it holds are the furthest away from the nucleus, and therefore are the ones most loosely held by its attractive force; the larger the distance between two charged objects, the smaller the force they exert on each other. Chemical reactivity of all of the different elements in the periodic table depends on the number of electrons in that last, outermost level, called the valence level or valence shell. In the case of gold, there is only one valence electron in its valence level.

Octet of Valence Electrons

Atoms gain, lose, or share electrons in their valence level in order to achieve greater stability, or a lower energy state. From this perspective, bonds between atoms form so that the bonded atoms are in a lower energy state compared to when they were by themselves. Atoms can achieve this more stable state by having a valence level which contains as many electrons as it can hold. For the first principal energy level, having two electrons in it is the most stable arrangement, while for all other levels outside of the first, eight electrons are necessary to achieve the most stable state.

Lewis Symbols

In the Lewis symbol for an atom, the chemical symbol of the element (as found on the periodic table) is written, and the valence electrons are represented as dots surrounding it. Only the electrons in the valence level are shown using this notation. For example, the Lewis symbol of carbon depicts a “C’ surrounded by 4 valence electrons because carbon has an electron configuration of 1s²2s²2p².

The Lewis symbol for carbon: Each of the four valence electrons is represented as a dot.

Electrons that are not in the valence level are not shown in the Lewis symbol. The reason for this is that the chemical reactivity of an atom of the element is solely determined by the number of its valence electrons, and not its inner electrons. Lewis symbols for atoms are combined to write Lewis structures for compounds or molecules with bonds between atoms.

Writing Lewis Symbols for Atoms

H Cation Lewis Structure Definition

The Lewis symbol for an atom depicts its valence electrons as dots around the symbol for the element.

Key Takeaways

Key Points

• The columns, or groups, in the periodic table are used to determine the number of valence electrons for each element.
• The noble/ inert gases are chemically stable and have a full valence level of electrons.
• Other elements react in order to achieve the same stability as the noble gases.
• Lewis symbols represent the valence electrons as dots surrounding the elemental symbol for the atom.

Key Terms

• group: A column in the periodic table that consists of elements with similar chemical reactivity, because they have the same number of valence electrons.
• Noble Gases: Inert, or unreactive, elements in the last group in the periodic table which are typically found in the gaseous form.
• Lewis symbol: Formalism in which the valence electrons of an atom are represented as dots.

Determining the Number of Valence Electrons

In order to write the Lewis symbol for an atom, you must first determine the number of valence electrons for that element. The arrangement of the periodic table can help you figure out this information. Since we have established that the number of valence electrons determines the chemical reactivity of an element, the table orders the elements by number of valence electrons.

Each column (or group) of the periodic table contains elements that have the same number of valence electrons. Furthermore, the number of columns (or groups) from the left edge of the table tells us the exact number of valence electrons for that element. Recall that any valence level can have up to eight electrons, except for the first principal energy level, which can only have two.

Periodic table of the elements: Group numbers shown by Roman numerals (above the table) tell us how many valence electrons there are for each element.

Some periodic tables list the group numbers in Arabic numbers instead of Roman numerals. In that case, the transition metal groups are included in the counting and the groups indicated at the top of the periodic table have numbers 1, 2, 13, 14, 15, 16, 17, 18. The corresponding roman numerals used are I, II, III, IV, V, VI, VII, VIII.

Survey of the Groups in the Periodic Table

Take the first column or group of the periodic table (labeled ‘I’): hydrogen (H), lithium (Li), sodium (Na), potassium (K), etc. Each of these elements has one valence electron. The second column or group (labeled ‘II’) means that beryllium (Be), magnesium (Mg), calcium (Ca), etc., all have two valence electrons.

The middle part of the periodic table that contains the transition metals is skipped in this process for reasons having to do with the electronic configuration of these elements.

Proceeding to the column labeled ‘III’, we find that those elements (B, Al, Ga, In,…) have three valence electrons in their outermost or valence level.

We can continue this inspection of the groups until we reach the eighth and final column, in which the most stable elements are listed. These are all gaseous under normal conditions of temperature and pressure, and are called ‘noble gases.’ Neon (Ne), argon (Ar), krypton (Kr), etc., each contain eight electrons in their valence level. Therefore, these elements have a full valence level that has the maximum number of electrons possible. Helium (He), at the very top of this column is an exception because it has two valence electrons; its valence level is the first principal energy level which can only have two electrons, so it has the maximum number of electrons in its valence level as well.

The Lewis symbol for helium: Helium is one of the noble gases and contains a full valence shell. Unlike the other noble gases in Group 8, Helium only contains two valence electrons. In the Lewis symbol, the electrons are depicted as two lone pair dots.

The noble gases represent elements of such stability that they are not chemically reactive, so they can be called inert. In other words, they don’t need to bond with any other elements in order to attain a lower energy configuration. We explain this phenomenon by attributing their stability to having a ‘full’ valence level.

The significance in understanding the nature of the stability of noble gases is that it guides us in predicting how other elements will react in order to achieve the same electronic configuration as the noble gases by having a full valence level.

Writing Lewis Symbols for Atoms

Lewis symbols for the elements depict the number of valence electrons as dots. In accordance with what we discussed above, here are the Lewis symbols for the first twenty elements in the periodic table. The heavier elements will follow the same trends depending on their group.

Once you can draw a Lewis symbol for an atom, you can use the knowledge of Lewis symbols to create Lewis structures for molecules.

Valence Electrons and the Periodic Table: Electrons can inhabit a number of energy shells. Different shells are different distances from the nucleus. The electrons in the outermost electron shell are called valence electrons, and are responsible for many of the chemical properties of an atom. This video will look at how to find the number of valence electrons in an atom depending on its column in the periodic table.

Introduction to Lewis Structures for Covalent Molecules

In covalent molecules, atoms share pairs of electrons in order to achieve a full valence level.

Learning Objectives

Predict and draw the Lewis structure of simple covalent molecules and compounds

Key Takeaways

Key Points

• The octet rule says that the noble gas electronic configuration is a particularly favorable one that can be achieved through formation of electron pair bonds between atoms.
• In many atoms, not all of the electron pairs comprising the octet are shared between atoms. These unshared, non-bonding electrons are called ‘ lone pairs ‘ of electrons.
• Although lone pairs are not directly involved in bond formation, they should always be shown in Lewis structures.
• There is a logical procedure that can be followed to draw the Lewis structure of a molecule or compound.

Key Terms

• octet rule: Atoms try to achieve the electronic configuration of the noble gas nearest to them in the periodic table by achieving a full valence level with eight electrons.
• exceptions to the octet rule: Hydrogen (H) and helium (He) only need two electrons to have a full valence level.
• covalent bond: Two atoms share valence electrons in order to achieve a noble gas electronic configuration.
• Lewis structure: Formalism used to show the structure of a molecule or compound, in which shared electrons pairs between atoms are indicated by dashes. Non-bonding, lone pairs of electrons must also be shown.

The Octet Rule

Noble gases like He, Ne, Ar, Kr, etc., are stable because their valence level is filled with as many electrons as possible. Eight electrons fill the valence level for all noble gases, except helium, which has two electrons in its full valence level. Other elements in the periodic table react to form bonds in which valence electrons are exchanged or shared in order to achieve a valence level which is filled, just like in the noble gases. We refer to this chemical tendency of atoms as ‘the octet rule,’ and it guides us in predicting how atoms combine to form molecules and compounds.

Covalent Bonds and Lewis Diagrams of Simple Molecules

The simplest example to consider is hydrogen (H), which is the smallest element in the periodic table with one proton and one electron. Hydrogen can become stable if it achieves a full valence level like the noble gas that is closest to it in the periodic table, helium (He). These are exceptions to the octet rule because they only require 2 electrons to have a full valence level.

Two H atoms can come together and share each of their electrons to create a ‘ covalent bond.’ The shared pair of electrons can be thought of as belonging to either atom, and thus each atom now has two electrons in its valence level, like He. The molecule that results is H₂, and it is the most abundant molecule in the universe.

Lewis structure of diatomic hydrogen: This is the process through which the H₂ molecule is formed. Two H atoms, each contributing an electron, share a pair of electrons. This is known as a ‘single covalent bond.’ Notice how the two electrons can be found in a region of space between the two atomic nuclei.

Lewis Structure Ions

The Lewis formalism used for the H₂ molecule is H:H or H—H. The former, known as a ‘Lewis dot diagram,’ indicates a pair of shared electrons between the atomic symbols, while the latter, known as a ‘Lewis structure,’ uses a dash to indicate the pair of shared electrons that form a covalent bond. More complicated molecules are depicted this way as well.

Lewis dot dragram for methane: Methane, with molecular formula CH₄, is shown. The electrons are color-coded to indicate which atoms they belonged to before the covalent bonds formed, with red representing hydrogen and blue representing carbon. Four covalent bonds are formed so that C has an octet of valence electrons, and each H has two valence electrons—one from the carbon atom and one from one of the hydrogen atoms.

Now consider the case of fluorine (F), which is found in group VII (or 17) of the periodic table. It therefore has 7 valence electrons and only needs 1 more in order to have an octet. One way that this can happen is if two F atoms make a bond, in which each atom provides one electron that can be shared between the two atoms. The resulting molecule that is formed is F₂, and its Lewis structure is F—F.

Achieving an octet of valence electrons: Two fluorine atoms are able to share an electron pair, which becomes a covalent bond. Notice that only the outer (valence level) electrons are involved, and that in each F atom, 6 valence electrons do not participate in bonding. These are ‘lone pairs’ of electrons.

After a bond has formed, each F atom has 6 electrons in its valence level which are not used to form a bond. These non-bonding valence electrons are called ‘lone pairs’ of electrons and should always be indicated in Lewis diagrams.

Lewis structure of acetic acid: Acetic acid, CH₃COOH, can be written out with dots indicating the shared electrons, or, preferably, with dashes representing covalent bonds. Notice the lone pairs of electrons on the oxygen atoms are still shown. The methyl group carbon atom has six valence electrons from its bonds to the hydrogen atoms because carbon is more electronegative than hydrogen. Also, one electron is gained from its bond with the other carbon atom because the electron pair in the C−C bond is split equally.

Procedure for Drawing Simple Lewis Structures

We have looked at how to determine Lewis structures for simple molecules. The procedure is as follows:

1. Write a structural diagram of the molecule to clearly show which atom is connected to which (although many possibilities exist, we usually pick the element with the most number of possible bonds to be the central atom).
2. Draw Lewis symbols of the individual atoms in the molecule.
3. Bring the atoms together in a way that places eight electrons around each atom (or two electrons for H, hydrogen) wherever possible.
4. Each pair of shared electrons is a covalent bond which can be represented by a dash.

Alternate view of lewis dot structure of water: This arrangement of shared electrons between O and H results in the oxygen atom having an octet of electrons, and each H atom having two valence electrons.

Multiple bonds can also form between elements when two or three pairs of electrons are shared to produce double or triple bonds, respectively. The Lewis structure for carbon dioxide, CO₂, is a good example of this.

Lewis structure of carbon dioxide: This figure explains the bonding in a CO₂ molecule. Each O atom starts out with six (red) electrons and C with four (black) electrons, and each bond behind an O atom and the C atom consists of two electrons from the O and two of the four electrons from the C.

In order to achieve an octet for all three atoms in CO₂, two pairs of electrons must be shared between the carbon and each oxygen. Since four electrons are involved in each bond, a double covalent bond is formed. You can see that this is how the octet rule is satisfied for all atoms in this case. When a double bond is formed, you still need to show all electrons, so double dashes between the atoms show that four electrons are shared.

Final Lewis structure for carbon dioxide: Covalent bonds are indicated as dashes and lone pairs of electrons are shown as pairs of dots. in carbon dioxide, each oxygen atom has two lone pairs of electrons remaining; the covalent bonds between the oxygen and carbon atoms each use two electrons from the oxygen atom and two from the carbon.

Lewis Structures for Polyatomic Ions

The Lewis structure of an ion is placed in brackets and its charge is written as a superscript outside of the brackets, on the upper right.

Learning Objectives

Apply the rules for drawing Lewis structures to polyatomic ions

Key Takeaways

Key Points

• Ions are treated almost the same way as a molecule with no charge. However, the number of electrons must be adjusted to account for the net electric charge of the ion.
• When counting electrons, negative ions should have extra electrons placed in their Lewis structures, while positive ions should have fewer electrons than an uncharged molecule.

Key Terms

• polyatomic ion: A charged species composed of two or more atoms covalently bonded, or of a metal complex that acts as a single unit in acid-base chemistry or in the formation of salts. Also known as a molecular ion.

The total number of electrons represented in a Lewis structure is equal to the sum of the numbers of valence electrons in each individual atom. Non-valence electrons are not represented in Lewis structures. After the total number of available electrons has been determined, electrons must be placed into the structure.

Lewis structures for polyatomic ions are drawn by the same methods that we have already learned. When counting electrons, negative ions should have extra electrons placed in their Lewis structures; positive ions should have fewer electrons than an uncharged molecule. When the Lewis structure of an ion is written, the entire structure is placed in brackets, and the charge is written as a superscript on the upper right, outside of the brackets. For example, consider the ammonium ion, NH₄⁺, which contains 9 (5 from N and 1 from each of the four H atoms) –1 = 8 electrons. One electron is subtracted because the entire molecule has a +1 charge.

Coordinate covalent bonding: The ammonium ion, NH4⁺, contains 9–1 = 8 electrons.

Negative ions follow the same procedure. The chlorite ion, ClO₂^–, contains 19 (7 from the Cl and 6 from each of the two O atoms) +1 = 20 electrons. One electron is added because the entire molecule has a -1 charge.

Hypochlorite ion Lewis structure: The hypochlorite ion, ClO⁻, contains 13 + 1 = 14 electrons.

Learning Objectives

• To use Lewis dot symbols to explain the stoichiometry of a compound

We begin our discussion of the relationship between structure and bonding in covalent compounds by describing the interaction between two identical neutral atoms—for example, the H₂ molecule, which contains a purely covalent bond. Each hydrogen atom in H₂ contains one electron and one proton, with the electron attracted to the proton by electrostatic forces. As the two hydrogen atoms are brought together, additional interactions must be considered (Figure (PageIndex{1})):

• The electrons in the two atoms repel each other because they have the same charge (E > 0).
• Similarly, the protons in adjacent atoms repel each other (E > 0).
• The electron in one atom is attracted to the oppositely charged proton in the other atom and vice versa (E < 0). Recall that it is impossible to specify precisely the position of the electron in either hydrogen atom. Hence the quantum mechanical probability distributions must be used.

A plot of the potential energy of the system as a function of the internuclear distance (Figure (PageIndex{2})) shows that the system becomes more stable (the energy of the system decreases) as two hydrogen atoms move toward each other from r = ∞, until the energy reaches a minimum at r = r₀ (the observed internuclear distance in H₂ is 74 pm). Thus at intermediate distances, proton–electron attractive interactions dominate, but as the distance becomes very short, electron–electron and proton–proton repulsive interactions cause the energy of the system to increase rapidly. Notice the similarity between Figures (PageIndex{1}) and (PageIndex{2}), which described a system containing two oppositely charged ions. The shapes of the energy versus distance curves in the two figures are similar because they both result from attractive and repulsive forces between charged entities.

No2+ Lewis Structure

At long distances, both attractive and repulsive interactions are small. As the distance between the atoms decreases, the attractive electron–proton interactions dominate, and the energy of the system decreases. At the observed bond distance, the repulsive electron–electron and proton–proton interactions just balance the attractive interactions, preventing a further decrease in the internuclear distance. At very short internuclear distances, the repulsive interactions dominate, making the system less stable than the isolated atoms.

Using Lewis Dot Symbols to Describe Covalent Bonding

The valence electron configurations of the constituent atoms of a covalent compound are important factors in determining its structure, stoichiometry, and properties. For example, chlorine, with seven valence electrons, is one electron short of an octet. If two chlorine atoms share their unpaired electrons by making a covalent bond and forming Cl₂, they can each complete their valence shell:

Each chlorine atom now has an octet. The electron pair being shared by the atoms is called a bonding pair; the other three pairs of electrons on each chlorine atom are called lone pairs. Lone pairs are not involved in covalent bonding. If both electrons in a covalent bond come from the same atom, the bond is called a coordinate covalent bond. Examples of this type of bonding are presented in Section 8.6 when we discuss atoms with less than an octet of electrons.

We can illustrate the formation of a water molecule from two hydrogen atoms and an oxygen atom using Lewis dot symbols:

The structure on the right is the Lewis electron structure, or Lewis structure, for H₂O. With two bonding pairs and two lone pairs, the oxygen atom has now completed its octet. Moreover, by sharing a bonding pair with oxygen, each hydrogen atom now has a full valence shell of two electrons. Chemists usually indicate a bonding pair by a single line, as shown here for our two examples:

The following procedure can be used to construct Lewis electron structures for more complex molecules and ions:

1. Arrange the atoms to show specific connections. When there is a central atom, it is usually the least electronegative element in the compound. Chemists usually list this central atom first in the chemical formula (as in CCl₄ and CO₃²⁻, which both have C as the central atom), which is another clue to the compound’s structure. Hydrogen and the halogens are almost always connected to only one other atom, so they are usually terminal rather than central.
2. Determine the total number of valence electrons in the molecule or ion. Add together the valence electrons from each atom. (Recall that the number of valence electrons is indicated by the position of the element in the periodic table.) If the species is a polyatomic ion, remember to add or subtract the number of electrons necessary to give the total charge on the ion. For CO₃²⁻, for example, we add two electrons to the total because of the −2 charge.
3. Place a bonding pair of electrons between each pair of adjacent atoms to give a single bond. In H₂O, for example, there is a bonding pair of electrons between oxygen and each hydrogen.
4. Beginning with the terminal atoms, add enough electrons to each atom to give each atom an octet (two for hydrogen). These electrons will usually be lone pairs.
5. If any electrons are left over, place them on the central atom. We will explain later that some atoms are able to accommodate more than eight electrons.
6. If the central atom has fewer electrons than an octet, use lone pairs from terminal atoms to form multiple (double or triple) bonds to the central atom to achieve an octet. This will not change the number of electrons on the terminal atoms.

Now let’s apply this procedure to some particular compounds, beginning with one we have already discussed.

The central atom is usually the least electronegative element in the molecule or ion; hydrogen and the halogens are usually terminal.

The (H_2O) Molecule

1. Because H atoms are almost always terminal, the arrangement within the molecule must be HOH.
2. Each H atom (group 1) has 1 valence electron, and the O atom (group 16) has 6 valence electrons, for a total of 8 valence electrons.
3. Placing one bonding pair of electrons between the O atom and each H atom gives H:O:H, with 4 electrons left over.
4. Each H atom has a full valence shell of 2 electrons.
5. Adding the remaining 4 electrons to the oxygen (as two lone pairs) gives the following structure:

This is the Lewis structure we drew earlier. Because it gives oxygen an octet and each hydrogen two electrons, we do not need to use step 6.

The (OCl^−) Ion

1. With only two atoms in the molecule, there is no central atom.
2. Oxygen (group 16) has 6 valence electrons, and chlorine (group 17) has 7 valence electrons; we must add one more for the negative charge on the ion, giving a total of 14 valence electrons.
3. Placing a bonding pair of electrons between O and Cl gives O:Cl, with 12 electrons left over.
4. If we place six electrons (as three lone pairs) on each atom, we obtain the following structure:

Each atom now has an octet of electrons, so steps 5 and 6 are not needed. The Lewis electron structure is drawn within brackets as is customary for an ion, with the overall charge indicated outside the brackets, and the bonding pair of electrons is indicated by a solid line. OCl⁻ is the hypochlorite ion, the active ingredient in chlorine laundry bleach and swimming pool disinfectant.

The (CH_2O) Molecule

1. Because carbon is less electronegative than oxygen and hydrogen is normally terminal, C must be the central atom. One possible arrangement is as follows:

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.

4. Adding all 6 remaining electrons to oxygen (as three lone pairs) gives the following:

Although oxygen now has an octet and each hydrogen has 2 electrons, carbon has only 6 electrons.

5. There are no electrons left to place on the central atom.

6. To give carbon an octet of electrons, we use one of the lone pairs of electrons on oxygen to form a carbon–oxygen double bond:

Both the oxygen and the carbon now have an octet of electrons, so this is an acceptable Lewis electron structure. The O has two bonding pairs and two lone pairs, and C has four bonding pairs. This is the structure of formaldehyde, which is used in embalming fluid.

An alternative structure can be drawn with one H bonded to O. Formal charges, discussed later in this section, suggest that such a structure is less stable than that shown previously.

Example (PageIndex{1})

Write the Lewis electron structure for each species.

1. NCl₃
2. S₂²⁻
3. NOCl

Given: chemical species

Asked for: Lewis electron structures

Strategy:

Use the six-step procedure to write the Lewis electron structure for each species.

Solution:

1. Nitrogen is less electronegative than chlorine, and halogen atoms are usually terminal, so nitrogen is the central atom. The nitrogen atom (group 15) has 5 valence electrons and each chlorine atom (group 17) has 7 valence electrons, for a total of 26 valence electrons. Using 2 electrons for each N–Cl bond and adding three lone pairs to each Cl account for (3 × 2) + (3 × 2 × 3) = 24 electrons. Rule 5 leads us to place the remaining 2 electrons on the central N:

   Nitrogen trichloride is an unstable oily liquid once used to bleach flour; this use is now prohibited in the United States.

2. In a diatomic molecule or ion, we do not need to worry about a central atom. Each sulfur atom (group 16) contains 6 valence electrons, and we need to add 2 electrons for the −2 charge, giving a total of 14 valence electrons. Using 2 electrons for the S–S bond, we arrange the remaining 12 electrons as three lone pairs on each sulfur, giving each S atom an octet of electrons:
3. Because nitrogen is less electronegative than oxygen or chlorine, it is the central atom. The N atom (group 15) has 5 valence electrons, the O atom (group 16) has 6 valence electrons, and the Cl atom (group 17) has 7 valence electrons, giving a total of 18 valence electrons. Placing one bonding pair of electrons between each pair of bonded atoms uses 4 electrons and gives the following:

   Adding three lone pairs each to oxygen and to chlorine uses 12 more electrons, leaving 2 electrons to place as a lone pair on nitrogen:

   Because this Lewis structure has only 6 electrons around the central nitrogen, a lone pair of electrons on a terminal atom must be used to form a bonding pair. We could use a lone pair on either O or Cl. Because we have seen many structures in which O forms a double bond but none with a double bond to Cl, it is reasonable to select a lone pair from O to give the following:

   All atoms now have octet configurations. This is the Lewis electron structure of nitrosyl chloride, a highly corrosive, reddish-orange gas.

Exercise (PageIndex{1})

Write Lewis electron structures for CO₂ and SCl₂, a vile-smelling, unstable red liquid that is used in the manufacture of rubber.

Answer

Using Lewis Electron Structures to Explain Stoichiometry

Lewis dot symbols provide a simple rationalization of why elements form compounds with the observed stoichiometries. In the Lewis model, the number of bonds formed by an element in a neutral compound is the same as the number of unpaired electrons it must share with other atoms to complete its octet of electrons. For the elements of Group 17 (the halogens), this number is one; for the elements of Group 16 (the chalcogens), it is two; for Group 15 elements, three; and for Group 14 elements four. These requirements are illustrated by the following Lewis structures for the hydrides of the lightest members of each group:

Elements may form multiple bonds to complete an octet. In ethylene, for example, each carbon contributes two electrons to the double bond, giving each carbon an octet (two electrons/bond × four bonds = eight electrons). Neutral structures with fewer or more bonds exist, but they are unusual and violate the octet rule.

Allotropes of an element can have very different physical and chemical properties because of different three-dimensional arrangements of the atoms; the number of bonds formed by the component atoms, however, is always the same. As noted at the beginning of the chapter, diamond is a hard, transparent solid; graphite is a soft, black solid; and the fullerenes have open cage structures. Despite these differences, the carbon atoms in all three allotropes form four bonds, in accordance with the octet rule.

Lewis structures explain why the elements of groups 14–17 form neutral compounds with four, three, two, and one bonded atom(s), respectively.

Elemental phosphorus also exists in three forms: white phosphorus, a toxic, waxy substance that initially glows and then spontaneously ignites on contact with air; red phosphorus, an amorphous substance that is used commercially in safety matches, fireworks, and smoke bombs; and black phosphorus, an unreactive crystalline solid with a texture similar to graphite (Figure (PageIndex{3})). Nonetheless, the phosphorus atoms in all three forms obey the octet rule and form three bonds per phosphorus atom.

Formal Charges

It is sometimes possible to write more than one Lewis structure for a substance that does not violate the octet rule, as we saw for CH₂O, but not every Lewis structure may be equally reasonable. In these situations, we can choose the most stable Lewis structure by considering the formal charge on the atoms, which is the difference between the number of valence electrons in the free atom and the number assigned to it in the Lewis electron structure. The formal charge is a way of computing the charge distribution within a Lewis structure; the sum of the formal charges on the atoms within a molecule or an ion must equal the overall charge on the molecule or ion. A formal charge does not represent a true charge on an atom in a covalent bond but is simply used to predict the most likely structure when a compound has more than one valid Lewis structure.

To calculate formal charges, we assign electrons in the molecule to individual atoms according to these rules:

• Nonbonding electrons are assigned to the atom on which they are located.
• Bonding electrons are divided equally between the bonded atoms.

For each atom, we then compute a formal charge:

( begin{matrix}
formal; charge= & valence; e^{-}- & left ( non-bonding; e^{-}+frac{bonding;e^{-}}{2} right )
& ^{left ( free; atom right )} & ^{left ( atom; in; Lewis; structure right )}
end{matrix} label{8.5.1} ) (atom in Lewis structure)

To illustrate this method, let’s calculate the formal charge on the atoms in ammonia (NH₃) whose Lewis electron structure is as follows:

A neutral nitrogen atom has five valence electrons (it is in group 15). From its Lewis electron structure, the nitrogen atom in ammonia has one lone pair and shares three bonding pairs with hydrogen atoms, so nitrogen itself is assigned a total of five electrons [2 nonbonding e⁻ + (6 bonding e⁻/2)]. Substituting into Equation (ref{8.5.2}), we obtain

[ formal; chargeleft ( N right )=5; valence; e^{-}-left ( 2; non-bonding; e^{-} +dfrac{6; bonding; e^{-}}{2} right )=0 label{8.5.2}]

A neutral hydrogen atom has one valence electron. Each hydrogen atom in the molecule shares one pair of bonding electrons and is therefore assigned one electron [0 nonbonding e⁻ + (2 bonding e⁻/2)]. Using Equation (ref{8.5.2}) to calculate the formal charge on hydrogen, we obtain

[ formal; chargeleft ( H right )=1; valence; e^{-}-left ( 0; non-bonding; e^{-} +dfrac{2; bonding; e^{-}}{2} right )=0 label{8.5.3}]

The hydrogen atoms in ammonia have the same number of electrons as neutral hydrogen atoms, and so their formal charge is also zero. Adding together the formal charges should give us the overall charge on the molecule or ion. In this example, the nitrogen and each hydrogen has a formal charge of zero. When summed the overall charge is zero, which is consistent with the overall charge on the NH₃ molecule.

An atom, molecule, or ion has a formal charge of zero if it has the number of bonds that is typical for that species.

Typically, the structure with the most charges on the atoms closest to zero is the more stable Lewis structure. In cases where there are positive or negative formal charges on various atoms, stable structures generally have negative formal charges on the more electronegative atoms and positive formal charges on the less electronegative atoms. The next example further demonstrates how to calculate formal charges.

Example (PageIndex{2}): The Ammonium Ion

Calculate the formal charges on each atom in the NH₄⁺ ion.

Given: chemical species

Asked for: formal charges

Strategy:

Nitride Ion Lewis Structure

Identify the number of valence electrons in each atom in the NH₄⁺ ion. Use the Lewis electron structure of NH₄⁺ to identify the number of bonding and nonbonding electrons associated with each atom and then use Equation (ref{8.5.2}) to calculate the formal charge on each atom.

Solution:

H Cation Lewis Structure Worksheet

The Lewis electron structure for the NH₄⁺ ion is as follows:

The nitrogen atom shares four bonding pairs of electrons, and a neutral nitrogen atom has five valence electrons. Using Equation (ref{8.5.1}), the formal charge on the nitrogen atom is therefore

[ formal; chargeleft ( N right )=5-left ( 0+dfrac{8}{2} right )=0 ]

Azide Ion Lewis Structure

Each hydrogen atom in has one bonding pair. The formal charge on each hydrogen atom is therefore

[ formal; chargeleft ( H right )=1-left ( 0+dfrac{2}{2} right )=0 ]

The formal charges on the atoms in the NH₄⁺ ion are thus

Adding together the formal charges on the atoms should give us the total charge on the molecule or ion. In this case, the sum of the formal charges is 0 + 1 + 0 + 0 + 0 = +1.

Exercise (PageIndex{2})

Write the formal charges on all atoms in BH₄⁻.

Answer

If an atom in a molecule or ion has the number of bonds that is typical for that atom (e.g., four bonds for carbon), its formal charge is zero.

Using Formal Charges to Distinguish Viable Lewis Structures

As an example of how formal charges can be used to determine the most stable Lewis structure for a substance, we can compare two possible structures for CO₂. Both structures conform to the rules for Lewis electron structures.

CO₂

Thiocyanate Ion Lewis Structure

1. C is less electronegative than O, so it is the central atom.
2. C has 4 valence electrons and each O has 6 valence electrons, for a total of 16 valence electrons.
3. Placing one electron pair between the C and each O gives O–C–O, with 12 electrons left over.
4. Dividing the remaining electrons between the O atoms gives three lone pairs on each atom:

This structure has an octet of electrons around each O atom but only 4 electrons around the C atom.

5. No electrons are left for the central atom.
6. To give the carbon atom an octet of electrons, we can convert two of the lone pairs on the oxygen atoms to bonding electron pairs. There are, however, two ways to do this. We can either take one electron pair from each oxygen to form a symmetrical structure or take both electron pairs from a single oxygen atom to give an asymmetrical structure:

Both Lewis electron structures give all three atoms an octet. How do we decide between these two possibilities? The formal charges for the two Lewis electron structures of CO₂ are as follows:

Both Lewis structures have a net formal charge of zero, but the structure on the right has a +1 charge on the more electronegative atom (O). Thus the symmetrical Lewis structure on the left is predicted to be more stable, and it is, in fact, the structure observed experimentally. Remember, though, that formal charges do not represent the actual charges on atoms in a molecule or ion. They are used simply as a bookkeeping method for predicting the most stable Lewis structure for a compound.

The Lewis structure with the set of formal charges closest to zero is usually (but not always) the most stable.

Example (PageIndex{3}): The Thiocyanate Ion

The thiocyanate ion (SCN⁻), which is used in printing and as a corrosion inhibitor against acidic gases, has at least two possible Lewis electron structures. Draw two possible structures, assign formal charges on all atoms in both, and decide which is the preferred arrangement of electrons.

Given: chemical species

Asked for: Lewis electron structures, formal charges, and preferred arrangement

Strategy:

1. Use the step-by-step procedure to write two plausible Lewis electron structures for SCN⁻.
2. Calculate the formal charge on each atom using Equation (ref{8.5.1}).
3. Predict which structure is preferred based on the formal charge on each atom and its electronegativity relative to the other atoms present.

Solution:

A Possible Lewis structures for the SCN⁻ ion are as follows:

B We must calculate the formal charges on each atom to identify the more stable structure. If we begin with carbon, we notice that the carbon atom in each of these structures shares four bonding pairs, the number of bonds typical for carbon, so it has a formal charge of zero. Continuing with sulfur, we observe that in (a) the sulfur atom shares one bonding pair and has three lone pairs and has a total of six valence electrons. The formal charge on the sulfur atom is therefore ( 6-left ( 6+frac{2}{2} right )=-1.5-left ( 4+frac{4}{2} right )=-1 ) In (c), nitrogen has a formal charge of −2.

C Which structure is preferred? Structure (b) is preferred because the negative charge is on the more electronegative atom (N), and it has lower formal charges on each atom as compared to structure (c): 0, −1 versus +1, −2.

Exercise (PageIndex{1}): The Fulminate Ion

Salts containing the fulminate ion (CNO⁻) are used in explosive detonators. Draw three Lewis electron structures for CNO⁻ and use formal charges to predict which is more stable. (Note: N is the central atom.)

Answer

The second structure is predicted to be more stable.

Summary

Lewis dot symbols provide a simple rationalization of why elements form compounds with the observed stoichiometries. A plot of the overall energy of a covalent bond as a function of internuclear distance is identical to a plot of an ionic pair because both result from attractive and repulsive forces between charged entities. In Lewis electron structures, we encounter bonding pairs, which are shared by two atoms, and lone pairs, which are not shared between atoms. If both electrons in a covalent bond come from the same atom, the bond is called a coordinate covalent bond. Lewis structures are an attempt to rationalize why certain stoichiometries are commonly observed for the elements of particular families. Neutral compounds of group 14 elements typically contain four bonds around each atom (a double bond counts as two, a triple bond as three), whereas neutral compounds of group 15 elements typically contain three bonds. In cases where it is possible to write more than one Lewis electron structure with octets around all the nonhydrogen atoms of a compound, the formal charge on each atom in alternative structures must be considered to decide which of the valid structures can be excluded and which is the most reasonable. The formal charge is the difference between the number of valence electrons of the free atom and the number of electrons assigned to it in the compound, where bonding electrons are divided equally between the bonded atoms. The Lewis structure with the lowest formal charges on the atoms is almost always the most stable one.

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