Ionic bonds usually occur between metal and nonmetal ions. 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. These substitution possibilities are shown in the above insert. Structural Formulas It is necessary to draw structural formulas for organic compounds because in most cases a molecular formula does not uniquely represent a single compound. Different compounds having the same molecular formula are called isomers , and the prevalence of organic isomers reflects the extraordinary versatility of carbon in forming strong bonds to itself and to other elements.
When the group of atoms that make up the molecules of different isomers are bonded together in fundamentally different ways, we refer to such compounds as constitutional isomers. There are seven constitutional isomers of C 4 H 10 O, and structural formulas for these are drawn in the following table. These formulas represent all known and possible C 4 H 10 O compounds, and display a common structural feature.
There are no double or triple bonds and no rings in any of these structures. Note that each of the carbon atoms is bonded to four other atoms, and is saturated with bonding partners.
Simplification of structural formulas may be achieved without any loss of the information they convey. In condensed structural formulas the bonds to each carbon are omitted, but each distinct structural unit group is written with subscript numbers designating multiple substituents, including the hydrogens.
Shorthand line formulas omit the symbols for carbon and hydrogen entirely. Each straight line segment represents a bond, the ends and intersections of the lines are carbon atoms, and the correct number of hydrogens is calculated from the tetravalency of carbon. Non-bonding valence shell electrons are omitted in these formulas. Developing the ability to visualize a three-dimensional structure from two-dimensional formulas requires practice, and in most cases the aid of molecular models.
As noted earlier, many kinds of model kits are available to students and professional chemists, and the beginning student is encouraged to obtain one. Constitutional isomers have the same molecular formula, but their physical and chemical properties may be very different. For an example Click Here. Distinguishing Carbon Atoms When discussing structural formulas, it is often useful to distinguish different groups of carbon atoms by their structural characteristics.
The three C 5 H 12 isomers shown below illustrate these terms. Structural differences may occur within these four groups, depending on the molecular constitution. A consideration of molecular symmetry helps to distinguish structurally equivalent from nonequivalent atoms and groups. The ability to distinguish structural differences of this kind is an essential part of mastering organic chemistry. It will come with practice and experience. Our ability to draw structural formulas for molecules is remarkable.
To see how this is done Click Here. Formula Analysis. Although structural formulas are essential to the unique description of organic compounds, it is interesting and instructive to evaluate the information that may be obtained from a molecular formula alone. Three useful rules may be listed: The number of hydrogen atoms that can be bonded to a given number of carbon atoms is limited by the valence of carbon.
The origin of this formula is evident by considering a hydrocarbon made up of a chain of carbon atoms. Here the middle carbons will each have two hydrogens and the two end carbons have three hydrogens each. Thus, when even-valenced atoms such as carbon and oxygen are bonded together in any number and in any manner, the number of remaining unoccupied bonding sites must be even. If these sites are occupied by univalent atoms such as H, F, Cl, etc.
If the four carbon atoms form a ring, two hydrogens must be lost. Similarly, the introduction of a double bond entails the loss of two hydrogens, and a triple bond the loss of four hydrogens. By rule 2 m must be an even number, so if m The presence of one or more nitrogen atoms or halogen substituents requires a modified analysis. The above formula may be extended to such compounds by a few simple principles: The presence of oxygen does not alter the relationship.
All halogens present in the molecular formula must be replaced by hydrogen. Each nitrogen in the formula must be replaced by a CH moiety. However, the structures of some compounds and ions cannot be represented by a single formula. For clarity the two ambiguous bonds to oxygen are given different colors in these formulas. If only one formula for sulfur dioxide was correct and accurate, then the double bond to oxygen would be shorter and stronger than the single bond.
This averaging of electron distribution over two or more hypothetical contributing structures canonical forms to produce a hybrid electronic structure is called resonance. The energy change associated with this bond depends on three main processes: the ionization of Na; the acceptance of the electron from a Na atom by a Cl atom; and Coulomb attraction of the resulting ions. If the ions get too close, they repel due to the exclusion principle 0. The equilibrium separation distance is.
Solution The energy change associated with the transfer of an electron from Na to Cl is 1. At equilibrium separation, the atoms are apart. The electrostatic potential energy of the atoms is.
The total energy difference associated with the formation of a NaCl formula unit is. Significance The formation of a NaCl formula unit by ionic bonding is energetically favorable. The dissociation energy, or energy required to separate the NaCl unit into ions is 4. Check Your Understanding Why is the potential energy associated with the exclusion principle positive in Figure? For a sodium ion in an ionic NaCl crystal, the expression for Coulomb potential energy must be modified by a factor known as the Madelung constant.
This factor takes into account the interaction of the sodium ion with all nearby chloride and sodium ions. The Madelung constant for a NaCl crystal is about 1. This value implies an equilibrium separation distance between ions of 0. We will return to this point again later.
In an ionic bond, an electron transfers from one atom to another. However, in a covalent bond, an electron is shared between two atoms. The ionic bonding mechanism cannot explain the existence of such molecules as and CO, since no separation distance exists for which the negative potential energy of attraction is greater in magnitude than the energy needed to create ions. Understanding precisely how such molecules are covalently bonded relies on a deeper understanding of quantum mechanics that goes beyond the coverage of this book, but we will qualitatively describe the mechanisms in the following section.
Covalent bonds can be understood using the simple example of a molecule, which consists of one electron in the electric field of two protons. This system can be modeled by an electron in a double square well Figure. The electron is equally likely to be found in each well, so the wave function is either symmetric or antisymmetric about a point midway between the wells.
Now imagine that the two wells are separated by a large distance. These states have the same energy. However, when the wells are brought together, the symmetric wave function becomes the ground state and the antisymmetric state becomes the first excited state—in other words, the energy level of the electron is split. Notice, the space-symmetric state becomes the energetically favorable lower energy state. The same analysis is appropriate for an electron bound to two hydrogen atoms.
Here, the shapes of the ground-state wave functions have the form or in one dimension. The energetically favorable, space-symmetric state implies a high charge density midway between the protons where the electrons are likely to pull the positively charged protons together.
If a second electron is added to this system to form a molecule, the wave function must describe both particles, including their spatial relationship and relative spins. This wave function must also respect the indistinguishability of electrons. Exchange symmetry can be symmetric , producing no change in the wave function, or antisymmetric , producing an overall change in the sign of the wave function—neither of which is observable. As we discuss later, the total wave function of two electrons must be antisymmetric on exchange.
For example, two electrons bound to a hydrogen molecule can be in a space-symmetric state with antiparallel spins or space-antisymmetric state with parallel spins.
The state with antiparallel spins is energetically favorable and therefore used in covalent bonding. If the protons are drawn too closely together, however, repulsion between the protons becomes important.
In other molecules, this effect is supplied by the exclusion principle. As a result, reaches an equilibrium separation of about 0. Because the electron bones in our analogy have a negative charge, the puppy thief becomes negatively charged due to the additional bone. The puppy that lost its electron bone becomes positively charged. Because the puppy who lost his bone has the opposite charge of the thief puppy, the puppies are held together by electrostatic forces, just like sodium and chloride ions!
In our analogy, each puppy again starts out with an electron bone. Some covalently bonded molecules, like chlorine gas Cl2 , equally share their electrons like two equally strong puppies each holding both bones. Other covalently bonded molecules, like hydrogen fluoride gas HF , do not share electrons equally.
The fluorine atom acts as a slightly stronger puppy that pulls a bit harder on the shared electrons see Fig. Even though the electrons in hydrogen fluoride are shared, the fluorine side of a water molecule pulls harder on the negatively charged shared electrons and becomes negatively charged.
The hydrogen atom has a slightly positively charge because it cannot hold as tightly to the negative electron bones. Covalent molecules with this type of uneven charge distribution are polar. Molecules with polar covalent bonds have a positive and negative side. In this analogy, each puppy represents an atom and each bone represents an electron. Water H2O , like hydrogen fluoride HF , is a polar covalent molecule.
When you look at a diagram of water see Fig.
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