Hydrogen bonding occurs only in molecules where hydrogen is covalently bonded to one of three elements: fluorine, oxygen, or nitrogen. These three elements are so electronegative that they withdraw the majority of the electron density in the covalent bond with hydrogen, leaving the H atom very electron-deficient.
The H atom nearly acts as a bare proton, leaving it very attracted to lone pair electrons on a nearby atom. The hydrogen bonding that occurs in water leads to some unusual, but very important properties. Most molecular compounds that have a mass similar to water are gases at room temperature.
Because of the strong hydrogen bonds, water molecules are able to stay condensed in the liquid state. In the liquid state, the hydrogen bonds of water can break and reform as the molecules flow from one place to another.
When water is cooled, the molecules begin to slow down. When water is liquid, the molecules are more motile and don't produce this rigid structure. Electrons are transferred between atoms. An ion will give one or more electrons to another ion. Table salt, sodium chloride NaCl , is a common example of an ionic compound. Note that sodium is on the left side of the periodic table and that chlorine is on the right side of the periodic table.
This happens to full fill their outermost valence shell. These ions bond because they experience an attractive force due to the difference in sign of their charges. Chemical Bonding When you think of bonding, you may not think of ions. Why Bonds Form To understand why chemical bonds form, consider the common compound known as water, or H 2 O. 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. To determine the chemical formulas of ionic compounds, the following two conditions must be satisfied:.
This is because Mg has two valence electrons and it would like to get rid of those two ions to obey the octet rule. Fluorine has seven valence electrons and usually forms the F — ion because it gains one electron to satisfy the octet rule. Therefore, the formula of the compound is MgF 2. The subscript two indicates that there are two fluorines that are ionically bonded to magnesium.
On the macroscopic scale, ionic compounds form crystalline lattice structures that are characterized by high melting and boiling points and good electrical conductivity when melted or solubilized. Fluorine has seven valence electrons and as such, usually forms the F — ion because it gains one electron to satisfy the octet rule.
Covalent bonds are a class of chemical bonds where valence electrons are shared between two atoms, typically two nonmetals. The formation of a covalent bond allows the nonmetals to obey the octet rule and thus become more stable. For example:. Covalent bonding requires a specific orientation between atoms in order to achieve the overlap between bonding orbitals. For example, sodium Na , a metal, and chloride Cl , a nonmetal, form an ionic bond to make NaCl.
In a covalent bond, the atoms bond by sharing electrons. Covalent bonds usually occur between nonmetals. For example, in water H 2 O each hydrogen H and oxygen O share a pair of electrons to make a molecule of two hydrogen atoms single bonded to a single oxygen atom. In general, ionic bonds occur between elements that are far apart on the periodic table.
Covalent bonds occur between elements that are close together on the periodic table. Ionic compounds tend to be brittle in their solid form and have very high melting temperatures. Covalent compounds tend to be soft, and have relatively low melting and boiling points. Water, a liquid composed of covalently bonded molecules, can also be used as a test substance for other ionic and covalently compounds.
Ionic compounds tend to dissolve in water e. Properties of ionic and covalent compounds are listed in Table 2. The properties listed in Table 2. Like other ionic compounds, sodium chloride Fig. Chlorine gas Fig. Ionic and covalent compounds also differ in what happens when they are placed in water, a common solvent. For example, when a crystal of sodium chloride is put into water, it may seem as though the crystal simply disappears.
Three things are actually happening. Ionic compounds like sodium chloride dissolve, dissociate, and diffuse. Covalent compounds, like sugar and food coloring, can dissolve and diffuse, but they do not dissociate.
Without stirring, the food coloring will mix into the water through only the movement of the water and food coloring molecules. As water evaporates, the salt solution becomes more and more concentrated. Eventually, there is not enough water left to keep the sodium and chloride ions from interacting and joining together, so salt crystals form.
This occurs naturally in places like salt evaporation ponds Fig. Salt crystals can also be formed by evaporating seawater in a shallow dish, as in the Recovering Salts from Seawater Activity. This document may be freely reproduced and distributed for non-profit educational purposes. Skip to main content. Search form Search. The application of resonance to this case requires a weighted averaging of these canonical structures.
The double bonded structure is regarded as the major contributor, the middle structure a minor contributor and the right hand structure a non-contributor. Since the middle, charge-separated contributor has an electron deficient carbon atom, this explains the tendency of electron donors nucleophiles to bond at this site. The basic principles of the resonance method may now be summarized.
These are the canonical forms to be considered, and all must have the same number of paired and unpaired electrons. The following factors are important in evaluating the contribution each of these canonical structures makes to the actual molecule. The stability of a resonance hybrid is always greater than the stability of any canonical contributor.
Consequently, if one canonical form has a much greater stability than all others, the hybrid will closely resemble it electronically and energetically. This is the case for the carbonyl group eq. On the other hand, if two or more canonical forms have identical low energy structures, the resonance hybrid will have exceptional stabilization and unique properties. This is the case for sulfur dioxide eq.
To illustrate these principles we shall consider carbon monoxide eq. In each case the most stable canonical form is on the left. For carbon monoxide, the additional bonding is more important than charge separation. Furthermore, the double bonded structure has an electron deficient carbon atom valence shell sextet. A similar destabilizing factor is present in the two azide canonical forms on the top row of the bracket three bonds vs.
The bottom row pair of structures have four bonds, but are destabilized by the high charge density on a single nitrogen atom. All the examples on this page demonstrate an important restriction that must be remembered when using resonance. No atoms change their positions within the common structural framework. Only electrons are moved. A more detailed model of covalent bonding requires a consideration of valence shell atomic orbitals.
The spatial distribution of electrons occupying each of these orbitals is shown in the diagram below. Very nice displays of orbitals may be found at the following sites: J. Gutow, Univ. Wisconsin Oshkosh R. Spinney, Ohio State M. Winter, Sheffield University. If this were the configuration used in covalent bonding, carbon would only be able to form two bonds.
In this case, the valence shell would have six electrons- two shy of an octet. However, the tetrahedral structures of methane and carbon tetrachloride demonstrate that carbon can form four equivalent bonds, leading to the desired octet. In order to explain this covalent bonding, Linus Pauling proposed an orbital hybridization model in which all the valence shell electrons of carbon are reorganized. These hybrid orbitals have a specific orientation, and the four are naturally oriented in a tetrahedral fashion.
Thus, the four covalent bonds of methane consist of shared electron pairs with four hydrogen atoms in a tetrahedral configuration, as predicted by VSEPR theory. Molecular Orbitals Just as the valence electrons of atoms occupy atomic orbitals AO , the shared electron pairs of covalently bonded atoms may be thought of as occupying molecular orbitals MO.
It is convenient to approximate molecular orbitals by combining or mixing two or more atomic orbitals. In general, this mixing of n atomic orbitals always generates n molecular orbitals. The hydrogen molecule provides a simple example of MO formation. The bonding MO is occupied by two electrons of opposite spin, the result being a covalent bond. The notation used for molecular orbitals parallels that used for atomic orbitals. In the case of bonds between second period elements, p-orbitals or hybrid atomic orbitals having p-orbital character are used to form molecular orbitals.
For example, the sigma molecular orbital that serves to bond two fluorine atoms together is generated by the overlap of p-orbitals part A below , and two sp 3 hybrid orbitals of carbon may combine to give a similar sigma orbital.
When these bonding orbitals are occupied by a pair of electrons, a covalent bond, the sigma bond results.
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