Name: shared electrons in covalent molecules
Uploaded: 2020-09-03T13:38:28.000Z
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K12 Standards

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North Carolina K-12 Science Standard Course of Study

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Chemistry

Matter and its Interactions

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Covalent Bonding and Lewis Structures

Covalent bonds are formed between nonmetals by the sharing of valence electrons. But why do nonmetals prefer to share electrons rather than transfer them like in ionic bonds? Nonmetals have high ionization energies, which makes it difficult to transfer valence electrons from one atom to another. 

Consider a molecule of ammonia. The nitrogen atom requires three more electrons to reach an octet, and the hydrogen atom needs an electron to reach a duet. 

Therefore, the nitrogen atom bonds with three hydrogen atoms such that both nitrogen and hydrogen reach stable electron configurations. Since each of these bonds shares one pair of electrons, it is called a single bond. 

The shared pair of electrons in the covalent bond is called a bonding pair. The valence electrons not participating in bonding are called the lone pair or nonbonding electrons. 

With 6 valence electrons, oxygen atoms need two more electrons to reach an octet. Therefore, two oxygen atoms share two-electron pairs forming a double bond. Nitrogen, on the other hand, shares the three unpaired electrons in its diatomic form, creating a triple bond. 

Single and multiple bonds differ in terms of bond length and strength. Triple bonds are shorter than double bonds, which are shorter than single bonds. The bond strength increases with bond multiplicity. This is why it is hard to break the triple bond in nitrogen, making it relatively unreactive. 

The Lewis model helps to predict the structure and stability of molecules. According to the Lewis model, H2O is a stable molecule; because both oxygen and hydrogen have achieved stable electron configurations. 

If the oxygen shares electrons with only one hydrogen atom, the resulting OH molecule is not stable since oxygen has only 7 valence electrons and cannot reach the octet. However, if an extra electron is added to the oxygen to complete the octet, the resulting hydroxide ion becomes stable with a negative charge. 

Covalent bonds are directional because the shared electrons link two specific pairs of atoms. In contrast, ionic bonds are nondirectional and hold several ions in the lattice. Thus, in a covalent compound, the individual molecules are considered to be fundamental units.

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 attraction of atoms for a &#8220;shared&#8221; pair of electrons. 

Covalent bonds are formed between two atoms when both have similar tendencies to attract electrons to themselves (i.e., when both atoms have identical or fairly similar ionization energies and electron affinities). 

<h4>Physical Properties of Covalent Compounds</h4>

Compounds that contain covalent bonds exhibit different physical properties than ionic compounds. Because the attraction between molecules, which are electrically neutral, is weaker than that between electrically charged ions, covalent compounds generally have much lower melting and boiling points than ionic compounds. In fact, many covalent compounds are liquids or gases at room temperature, and, in their solid states, they are typically much softer than ionic solids. Furthermore, whereas ionic compounds are good conductors of electricity when dissolved in water, most covalent compounds are insoluble in water; since they are electrically neutral, they are poor conductors of electricity in any state.

<h4>Formation of Covalent Bonds</h4>

Nonmetal atoms frequently form covalent bonds with other nonmetal atoms. For example, the hydrogen molecule, H2, contains a covalent bond between its two hydrogen atoms. Two separate hydrogen atoms with particular potential energy approach each other, their valence orbitals (1s) begin to overlap. The single electrons on each hydrogen atom then interact with both atomic nuclei, occupying the space around both atoms. The strong attraction of each shared electron to both nuclei stabilizes the system, and the potential energy decreases as the bond distance decreases. If the atoms continue to approach each other, the positive charges in the two nuclei begin to repel each other, and the potential energy increases. The bond length is determined by the distance at which the lowest potential energy is achieved.

It is essential to remember that energy must be added to break chemical bonds (an endothermic process), whereas forming chemical bonds releases energy (an exothermic process). In the case of H2, the covalent bond is very strong; a large amount of energy, 436 kJ, must be added to break the bonds in one mole of hydrogen molecules and cause the atoms to separate:

<img alt="Hydrogen decomposition equation, H&#8322;(g) &#8594; 2H(g), &#916;H=436 kJ, thermodynamic process." src="/files/core/CDNSource/lessons/1059/1059_Equations-01.png" style="height: 40px;" data-altupdatedat="1754904545000" data-alt="Chemical reaction equation, H2(g) to 2H(g), with enthalpy &#916;H=436 kJ, indicating bond dissociation." />

Conversely, the same amount of energy is released when one mole of H2 molecules forms from two moles of H atoms:

<img alt="Chemical reaction equation 2H(g) &#8594; H&#8322;(g), &#916;H=-436 kJ, representing exothermic reaction." src="/files/core/CDNSource/lessons/1059/1059_Equations-02.png" style="height: 40px;" data-altupdatedat="1754904545000" data-alt="Chemical reaction diagram: 2H(g) &#8594; H&#8322;(g), enthalpy change &#916;H = -436 kJ, exothermic process." />

<h4>Lewis Structures</h4>

Lewis symbols can be used 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:

<img alt="Chlorine bonding diagram, Cl atoms form Cl&#8322; molecule, showcasing electron pair sharing." src="/files/core/CDNSource/lessons/1059/1059Image_1.svg" style="height: 90px;" data-altupdatedat="1754904545000" data-alt="Chlorine bond formation; Cl atoms pair to form Cl&#8322; molecule; chemical equation diagram." />

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: Cl&#8212;Cl.

<ul>
	<li>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.</li>
	<li>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 CH2O (formaldehyde) and between the two carbon atoms in C2H4 (ethylene). 
	<img alt="Lewis structures of formaldehyde and ethylene, molecular geometry and bonding concepts diagram." src="/files/core/CDNSource/lessons/1059/1059Image_2.svg" style="height: 120px;" data-altupdatedat="1754904545000" data-alt="Formaldehyde and ethylene Lewis structures diagram; chemical bonding, electron pair representation." /></li>
	<li>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&#8211;). 
	<img alt="Lewis structures of carbon monoxide and cyanide ion; electron arrangement diagram." src="/files/core/CDNSource/lessons/1059/1059Image_3.svg" style="height: 90px;" data-altupdatedat="1754904545000" data-alt="Lewis structures of carbon monoxide and cyanide ion; chemical bonding diagram." /></li>
</ul>

The periodic table can be used to predict the number of valence electrons in an atom and the number of bonds that will be formed to reach an octet. Group 18 elements such as Argon and Helium have filled electron configuration and thus rarely participate in chemical bonding. However, atoms from group 17, such as bromine or iodine, need only one electron to reach octet. Hence atoms belonging to group 17 can form a single covalent bond. The atoms of group 16 need 2 electrons to reach an octet; hence they can form two covalent bonds. Similarly, carbon that belongs to group 14, needs 4 more electrons to reach an octet; thus carbon can form four covalent bonds. 

This text is adapted from <a href="https://openstax.org/books/chemistry-2e/pages/7-2-covalent-bonding">Openstax, Chemistry 2e, Section 7.2: Covalent Bonds</a> and <a href="https://openstax.org/books/chemistry-2e/pages/7-3-lewis-symbols-and-structures">Openstax, Chemistry 2e, Section 7.3: Lewis Symbols and Structures.</a>

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

Education

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Chemical Bonding: Basic Concepts

KEY TERMS AND CONCEPTS

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 have lower melting and boiling points, and some are liquids or gases at room temperature. When they are solids, they are usually softer than ionic solids. Most covalent compounds are also insoluble in water and are poor conductors of electricity in any state.

Hydrogen gas shows how a covalent bond forms. In H2, two hydrogen atoms approach each other and their 1s valence orbitals overlap. Each electron interacts with both nuclei, which lowers the potential energy and stabilizes the molecule. If the atoms move too close, the positive nuclei repel each other, and the potential energy rises. The bond length is the distance where the potential energy is lowest. Forming a bond releases energy, so it is an exothermic process. Breaking a bond requires energy, so it is endothermic. For H2, 436 kJ must be added to break one mole of molecules, and the same amount is released when one mole forms from hydrogen atoms.

Lewis structures show covalent bonding with Lewis symbols. In chlorine gas, Cl2, the two chlorine atoms share one pair of electrons. Each chlorine atom also has three lone pairs, which are pairs not used in bonding. A shared pair can be shown with a dash, so Cl—Cl represents a single bond. Each chlorine atom then has eight valence electrons around it, including lone pairs and the bonding pair.

Some atoms need more than one shared pair to reach an octet, which is a full outer shell of eight electrons. A double bond forms when two pairs of electrons are shared, as in CH2O and C2H4. A triple bond forms when three pairs are shared, as in CO and CN−. The periodic table helps predict bonding patterns from valence electrons. Group 18 elements such as argon and helium already have filled electron configurations and rarely bond. Group 17 atoms, such as bromine and iodine, usually need one electron and can form one covalent bond. Group 16 atoms need two electrons and can form two covalent bonds. Carbon in group 14 needs four more electrons, so it can form four covalent bonds.

Shared Electrons in Covalent Molecules

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...

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...

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...

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...

trends in lattice energy: ion size and charge

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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 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...

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...

bond polarity, dipole moment, and percent ionic character

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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

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Drawing Lewis Structures with Valence Electron Counts

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

resonance

Resonance Forms in Lewis Structures

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

formal charges

Lewis Structures with Formal Charge

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

exceptions to the octet rule

Octet Rule Exceptions in Lewis Structures

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

bond 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

Explore how nonmetal atoms share electrons to form covalent bonds, from H2 and Cl2 to double and triple bonds, Lewis structures, and bonding patterns.

Shared Electrons in Covalent Molecules | JoVE K12

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Shared Electrons in Covalent Molecules

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