Covalent bonding generally happens between nonmetals. Covalent bonding is the type of bond that holds together the atoms within a polyatomic ion. It takes two electrons to make a covalent bond, one from each bonding atom. Lewis dot structures are one way to represent how atoms form covalent bonds.
A table of Lewis dot symbols of nonmetal elements that form covalent bonds is shown in Fig. There can be up to eight dots, for eight valence electrons. The first four electrons are placed as single electrons, then the remaining four are paired. The number of bonds that each element is able to form is usually equal to the number of unpaired electrons.
In order to form a covalent bond, each element has to share one unpaired electron. First, determine how many atoms of each element are needed to satisfy the octet rule for each atom. In the formation of water, an oxygen atom has two unpaired electrons, and each hydrogen atom has one Fig. To fill its valence shell, oxygen needs two additional electrons, and hydrogen needs one.
One oxygen atom can share its unpaired electrons with two hydrogen atoms, each of which need only one additional electron. The single electrons match up to make pairs Fig. The oxygen atom forms two bonds, one with each of two hydrogen atoms; therefore, the formula for water is H 2 O.
When an electron, or dot, from one element is paired with an electron, or dot, from another element, this makes a bond, which is represented by a line Fig. The number of bonds that an element can form is determined by the number of electrons in its valence shell Fig.
Similarly, the number of electrons in the valence shell also determines ion formation. The octet rule applies for covalent bonding, with a total of eight electrons the most desirable number of unshared or shared electrons in the outer valence shell.
For example, carbon has an atomic number of six, with two electrons in shell 1 and four electrons in shell 2, its valence shell see Fig. This means that carbon needs four electrons to achieve an octet. Carbon is represented with four unpaired electrons see Fig. If carbon can share four electrons with other atoms, its valence shell will be full.
Most elements involved in covalent bonding need eight electrons to have a complete valence shell. One notable exception is hydrogen H. Hydrogen can be considered to be in Group 1 or Group 17 because it has properties similar to both groups. Hydrogen can participate in both ionic and covalent bonding. When participating in covalent bonding, hydrogen only needs two electrons to have a full valence shell.
As it has only one electron to start with, it can only make one bond. Hydrogen is shown in Fig 2. In the formation of a covalent hydrogen molecule, therefore, each hydrogen atom forms a single bond, producing a molecule with the formula H 2.
A single bond is defined as one covalent bond, or two shared electrons, between two atoms. A molecule can have multiple single bonds. For example, water, H 2 O, has two single bonds, one between each hydrogen atom and the oxygen atom Fig.
Figure 2. Sometimes two covalent bonds are formed between two atoms by each atom sharing two electrons, for a total of four shared electrons. For example, in the formation of the oxygen molecule, each atom of oxygen forms two bonds to the other oxygen atom, producing the molecule O 2. Since hydrogen is the simplest of atoms, having only one electron, diatomic hydrogen is the simplest of molecules, having only a single covalent bond. When more complex atoms form covalent bonds, the molecules they form are also more complex, involving numerous covalent bonds.
In some instances, an atom will have valence electrons that are not involved in bonding. These valence electrons are known as lone pairs. The Lewis structures of some common atoms are shown below. Notice that each structure satisfies the octet rule for all its atoms. When two atoms share a single pair of electrons, the bond is referred to as a single bond.
Atoms can also share two or three pairs of electrons in the aptly named double and triple bonds. In Lewis structures, multiple bonds are depicted by two or three lines between the bonded atoms.
The bond order of a covalent interaction between two atoms is the number of electron pairs that are shared between them. Single bonds have a bond order of 1, double bonds 2, and triple bonds 3.
Bond order is directly related to bond strength and bond length. Higher order bonds are stronger and shorter, while lower order bonds are weaker and longer. Not all covalent bonds are fit for Sesame Street: some covalent bonds are shared unequally. Atoms with equal or similar electronegativity form covalent bonds, in which the valence electron density is shared between the two atoms.
The electron density resides between the atoms and is attracted to both nuclei. This type of bond forms most frequently between two non- metals. When there is a greater electronegativity difference than between covalently bonded atoms, the pair of atoms usually forms a polar covalent bond.
The electrons are still shared between the atoms, but the electrons are not equally attracted to both elements. As a result, the electrons tend to be found near one particular atom most of the time. Again, polar covalent bonds tend to occur between non-metals. Finally, for atoms with the largest electronegativity differences such as metals bonding with nonmetals , the bonding interaction is called ionic, and the valence electrons are typically represented as being transferred from the metal atom to the nonmetal.
Once the electrons have been transferred to the non-metal, both the metal and the non-metal are considered to be ions. The two oppositely charged ions attract each other to form an ionic compound.
Covalent interactions are directional and depend on orbital overlap, while ionic interactions have no particular directionality. Each of these interactions allows the atoms involved to gain eight electrons in their valence shell, satisfying the octet rule and making the atoms more stable. These atomic properties help describe the macroscopic properties of compounds. For example, smaller covalent compounds that are held together by weaker bonds are frequently soft and malleable.
On the other hand, longer-range covalent interactions can be quite strong, making their compounds very durable. Ionic compounds, though composed of strong bonding interactions, tend to form brittle crystalline lattices.
Ionic bonds are a subset of chemical bonds that result from the transfer of valence electrons, typically between a metal and a nonmetal. Ionic bonds are a class of chemical bonds that result from the exchange of one or more valence electrons from one atom, typically a metal, to another, typically a nonmetal.
This electron exchange results in an electrostatic attraction between the two atoms called an ionic bond. An atom that loses one or more valence electrons to become a positively charged ion is known as a cation, while an atom that gains electrons and becomes negatively charged is known as an anion.
This exchange of valence electrons allows ions to achieve electron configurations that mimic those of the noble gases, satisfying the octet rule. The octet rule states that an atom is most stable when there are eight electrons in its valence shell.
Atoms with less than eight electrons tend to satisfy the duet rule, having two electrons in their valence shell. By satisfying the duet rule or the octet rule, ions are more stable. An anion is indicated by a negative superscript charge - something to the right of the atom.
Similarly, if a chlorine atom gains an extra electron, it becomes the chloride ion, Cl —. Both ions form because the ion is more stable than the atom due to the octet rule. Once the oppositely charged ions form, they are attracted by their positive and negative charges and form an ionic compound.
Ionic bonds are also formed when there is a large electronegativity difference between two atoms. This difference causes an unequal sharing of electrons such that one atom completely loses one or more electrons and the other atom gains one or more electrons, such as in the creation of an ionic bond between a metal atom sodium and a nonmetal fluorine.
Formation of sodium fluoride : The transfer of electrons and subsequent attraction of oppositely charged ions.
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