Covalent Molecules and Compounds

As was stated in section 4.1, each covalent compound is represented by a molecular formulaA representation of a covalent compound that consists of the atomic symbol for each component element (in a prescribed order) accompanied by a subscript indicating the number of atoms of that element in the molecule. The subscript is written only if the number is greater than 1., which gives the atomic symbol for each component element, in a prescribed order, accompanied by a subscript indicating the number of atoms of that element in the molecule. The subscript is written only if the number of atoms is greater than 1. For example, water, with two hydrogen atoms and one oxygen atom per molecule, is written as H₂O. Similarly, carbon dioxide, which contains one carbon atom and two oxygen atoms in each molecule, is written as CO₂.

Covalent compounds that contain predominantly carbon and hydrogen are called organic compoundsA covalent compound that contains predominantly carbon and hydrogen.. The convention for representing the formulas of organic compounds is to write carbon first, followed by hydrogen and then any other elements in alphabetical order (e.g., CH₄O is methyl alcohol, a fuel). Compounds that consist primarily of elements other than carbon and hydrogen are called inorganic compoundsAn ionic or covalent compound that consists primarily of elements other than carbon and hydrogen.; they include both covalent and ionic compounds. In inorganic compounds, the component elements are listed beginning with the one farthest to the left in the periodic table (see "Appendix H: Periodic Table of Elements" or click on periodic table button in the navigation bar at the top of this page), such as we see in CO₂ or SF₆.

Representations of Molecular Structures

Molecular formulas give only the elemental composition of molecules. In contrast, structural formulasA representation of a molecule that shows which atoms are bonded to one another and, in some cases, the approximate arrangement of atoms in space. show which atoms are bonded to one another and, in some cases, the approximate arrangement of the atoms in space. Knowing the structural formula of a compound enables chemists to create a three-dimensional model, which provides information about how that compound will behave physically and chemically.

The structural formula for water can be drawn as follows:

H-O-H or

Because the latter approximates the experimentally determined shape of the water molecule, it is more informative. Similarly, ammonia (NH₃) and methane (CH₄) are often written as planar molecules:

As shown in Figure 4.2(a) , however, the actual three-dimensional structure of NH₃ looks like a pyramid with a triangular base of three hydrogen atoms. The structure of CH₄, with four hydrogen atoms arranged around a central carbon atom as shown in Figure 4.2(a), is tetrahedral. That is, the hydrogen atoms are positioned at every other vertex of a cube. Many compounds-carbon compounds, in particular-have four bonded atoms arranged around a central atom to form a tetrahedron.

Figure 4.2(a) The Three-Dimensional Structures of Water, Ammonia, and Methane

(a) Water is a V-shaped molecule, in which all three atoms lie in a plane. (b) In contrast, ammonia has a pyramidal structure, in which the three hydrogen atoms form the base of the pyramid and the nitrogen atom is at the vertex. (c) The four hydrogen atoms of methane form a tetrahedron; the carbon atom lies in the center.

Figure 4.2(a) illustrates different ways to represent the structures of molecules. It should be clear that there is no single "best" way to draw the structure of a molecule; the method you use depends on which aspect of the structure you want to emphasize and how much time and effort you want to spend.

Ionic Compounds

The substances described in the preceding discussion are composed of molecules that are electrically neutral; that is, the number of positively charged protons in the nucleus is equal to the number of negatively charged electrons. In contrast, ions are atoms or assemblies of atoms that have a net electrical charge. Ions that contain fewer electrons than protons have a net positive charge and are called cationsAn ion that has fewer electrons than protons, resulting in a net positive charge.. Conversely, ions that contain more electrons than protons have a net negative charge and are called anionsAn ion that has fewer protons than electrons, resulting in a net negative charge.. Ionic compounds contain both cations and anions in a ratio that results in no net electrical charge.

Note the Pattern

Ionic compounds contain both cations and anions in a ratio that results in zero electrical charge.

In covalent compounds, electrons are shared between bonded atoms and are simultaneously attracted to more than one nucleus. In contrast, ionic compounds contain cations and anions rather than discrete neutral molecules. Ionic compounds are held together by the attractive electrostatic interactions between cations and anions. In an ionic compound, the cations and anions are arranged in space to form an extended three-dimensional array that maximizes the number of attractive electrostatic interactions and minimizes the number of repulsive electrostatic interactions (Figure 4.2(b) ). As shown in Equation 4.2(e1), the electrostatic energy of the interaction between two charged particles is proportional to the product of the charges on the particles and inversely proportional to the distance between them:

Equation 4.2(e1)

$electrostaticenergy\propto \frac{\left(Q_1\right)\left(Q_2\right)}{r}$

where Q₁ and Q₂ are the electrical charges on particles 1 and 2, and r is the distance between them. When Q₁ and Q₂ are both positive, corresponding to the charges on cations, the cations repel each other and the electrostatic energy is positive. When Q₁ and Q₂ are both negative, corresponding to the charges on anions, the anions repel each other and the electrostatic energy is again positive. The electrostatic energy is negative only when the charges have opposite signs; that is, positively charged species are attracted to negatively charged species and vice versa. As shown in Figure 4.2(c) , the strength of the interaction is proportional to the magnitude of the charges and decreases as the distance between the particles increases.

Note the Pattern

If the electrostatic energy is positive, the particles repel each other; if the electrostatic energy is negative, the particles are attracted to each other.

Figure 4.2(b) Covalent and Ionic Bonding

(a) In molecular hydrogen (H₂), two hydrogen atoms share two electrons to form a covalent bond. (b) The ionic compound NaCl forms when electrons from sodium atoms are transferred to chlorine atoms. The resulting Na⁺ and Cl^- ions form a three-dimensional solid that is held together by attractive electrostatic interactions.

Figure 4.2(c) The Effect of Charge and Distance on the Strength of Electrostatic Interactions

As the charge on ions increases or the distance between ions decreases, so does the strength of the attractive (-)…(+) or repulsive (-)…(-) or (+)…(+) interactions. The strength of these interactions is represented by the thickness of the arrows.

One example of an ionic compound is sodium chloride (NaCl; Figure 4.2(d) ), formed from sodium and chlorine. In forming chemical compounds, many elements have a tendency to gain or lose enough electrons to attain the same number of electrons as the noble gas closest to them in the periodic table. When sodium and chlorine come into contact, each sodium atom gives up an electron to become a Na⁺ ion, with 11 protons in its nucleus but only 10 electrons (like neon), and each chlorine atom gains an electron to become a Cl^- ion, with 17 protons in its nucleus and 18 electrons (like argon), as shown in part (b) in Figure 4.2(b) . Solid sodium chloride contains equal numbers of cations (Na⁺) and anions (Cl^-), thus maintaining electrical neutrality. Each Na⁺ ion is surrounded by 6 Cl^- ions, and each Cl^- ion is surrounded by 6 Na⁺ ions. Because of the large number of attractive Na⁺Cl^- interactions, the total attractive electrostatic energy in NaCl is great.

Figure 4.2(d) Sodium Chloride: an Ionic Solid

The planes of an NaCl crystal reflect the regular three-dimensional arrangement of its Na⁺ (purple) and Cl^- (green) ions. Looking for all the world like a snowflake, this is actually a close up view of sodium chloride crystals. The crystals are in a water bubble within a 50-millimetre metal loop that was part of an experiment in the Destiny laboratory aboard the International Space Station.Image Credit: By Photograph by the NASA Expedition 6 crew (NASA Image of the Day) [Public domain], via Wikimedia Commons; Inset: Public domain by Benjah_bmm27 via Wikimedia Commons.

Video: A shiny ribbon of magnesium metal is reacted with oxygen from the air. The physical properties of the magnesium metal are contrasted with the physical properties of the ionic salt (magnesium oxide). Video Credit: Part of NCSSM CORE collection: http://www.dlt.ncssm. Please attribute this work as being created by the North Carolina School of Science and Mathematics. This work is licensed under Creative Commons CC-BY https://creativecommons.org/licenses/by/3.0/ via YouTube

Ionization Energies

Because atoms do not spontaneously lose electrons, energy is required to remove an electron from an atom to form a cation. Chemists define the ionization energy(I)The minimum amount of energy needed to remove an electron from the gaseous atom in its ground state: $\text{E(g)}+I\to {\text{E}}^{+}\text{(g)}+{\text{e}}^{-}\text{.}$ of an element as the amount of energy needed to remove an electron from the gaseous atom E in its ground state. I is therefore the energy required for the reaction

Equation 4.2(e2)

$$\text{E}(\text{g})\to {\text{E}}^{\text{+}}(\text{g})+{\text{e}}^{-}\text{energy required}=I$$

Because an input of energy is required, the ionization energy is always positive (I > 0) for the reaction as written in Equation 4.2(e2). Larger values of I mean that the electron is more tightly bound to the atom and harder to remove.

Equation 4.2(e3)

$$\text{E}(\text{g})\to {\text{E}}^{+}(\text{g})+{\text{e}}^{-}{I}_1=1\text{st ionization energy}$$

The first ionization energies of the elements in the first six rows of the periodic table are plotted in Figure 4.2(e) . This figure illustrates three important trends:

1. The changes seen in the second (Li to Ne), fourth (K to Kr), fifth (Rb to Xe), and sixth (Cs to Rn) rows of the s and p blocks follow a pattern similar to the pattern described for the third row of the periodic table. The transition metals are included in the fourth, fifth, and sixth rows, however, and the lanthanides are included in the sixth row. The first ionization energies of the transition metals are somewhat similar to one another, as are those of the lanthanides. Ionization energies increase from left to right across each row, with discrepancies occurring at ns²np¹ (group 13), ns²np⁴ (group 16), and ns²(n − 1)d¹⁰ (group 12) electron configurations.
2. First ionization energies generally decrease down a column. Although the principal quantum number n increases down a column, filled inner shells are effective at screening the valence electrons, so there is a relatively small increase in the effective nuclear charge. Consequently, the atoms become larger as they acquire electrons. Valence electrons that are farther from the nucleus are less tightly bound, making them easier to remove, which causes ionization energies to decrease. A larger radius corresponds to a lower ionization energy.
3. Because of the first two trends, the elements that form positive ions most easily (have the lowest ionization energies) lie in the lower left corner of the periodic table, whereas those that are hardest to ionize lie in the upper right corner of the periodic table. Consequently, ionization energies generally increase diagonally from lower left (Cs) to upper right (He).

Note the Pattern

Generally, I₁ increases diagonally from the lower left of the periodic table to the upper right.

Figure 4.2(e) A Plot of Periodic Variation of First Ionization Energy with Atomic Number for the First Six Rows of the Periodic Table

There is a decrease in ionization energy within a group (most easily seen here for groups 1 and 18).

Physical Properties of Ionic and Covalent Compounds

In general, ionic and covalent compounds have different physical properties. Ionic compounds usually form hard crystalline solids that melt at rather high temperatures and are very resistant to evaporation. These properties stem from the characteristic internal structure of an ionic solid, illustrated schematically in part (a) in Figure 4.2(f) , which shows the three-dimensional array of alternating positive and negative ions held together by strong electrostatic attractions. In contrast, as shown in part (b) in Figure 4.2(f) , most covalent compounds consist of discrete molecules held together by comparatively weak intermolecular forces (the forces between molecules), even though the atoms within each molecule are held together by strong intramolecular covalent bonds (the forces within the molecule). Covalent substances can be gases, liquids, or solids at room temperature and pressure, depending on the strength of the intermolecular interactions. Covalent molecular solids tend to form soft crystals that melt at rather low temperatures and evaporate relatively easily.Some covalent substances, however, are not molecular but consist of infinite three-dimensional arrays of covalently bonded atoms and include some of the hardest materials known, such as diamond. The covalent bonds that hold the atoms together in the molecules are unaffected when covalent substances melt or evaporate, so a liquid or vapor of discrete, independent molecules is formed. For example, at room temperature, methane, the major constituent of natural gas, is a gas that is composed of discrete CH₄ molecules. A comparison of the different physical properties of ionic compounds and covalent molecular substances is given in Table 4.2(1) .

Table 4.2(1) The Physical Properties of Typical Ionic Compounds and Covalent Molecular Substances

Ionic Compounds	Covalent Molecular Substances
hard solids	gases, liquids, or soft solids
high melting points	low melting points
nonvolatile	volatile

Figure 4.2(f) Interactions in Ionic and Covalent Solids

(a) The positively and negatively charged ions in an ionic solid such as sodium chloride (NaCl) are held together by strong electrostatic interactions. (b) In this representation of the packing of methane (CH₄) molecules in solid methane, a prototypical molecular solid, the methane molecules are held together in the solid only by relatively weak intermolecular forces, even though the atoms within each methane molecule are held together by strong covalent bonds.

Summary

Ionic compounds contain positively and negatively charged ions in a ratio that results in an overall charge of zero. The ions are held together in a regular spatial arrangement by electrostatic forces. Most covalent compounds consist of molecules, groups of atoms in which one or more pairs of electrons are shared by at least two atoms to form a covalent bond. The atoms in molecules are held together by the electrostatic attraction between the positively charged nuclei of the bonded atoms and the negatively charged electrons shared by the nuclei. The molecular formula of a covalent compound gives the types and numbers of atoms present. Compounds that contain predominantly carbon and hydrogen are called organic compounds, whereas compounds that consist primarily of elements other than carbon and hydrogen are inorganic compounds. Diatomic molecules contain two atoms, and polyatomic molecules contain more than two. A structural formula indicates the composition and approximate structure and shape of a molecule. Single bonds, double bonds, and triple bonds are covalent bonds in which one, two, and three pairs of electrons, respectively, are shared between two bonded atoms. Atoms or groups of atoms that possess a net electrical charge are called ions; they can have either a positive charge (cations) or a negative charge (anions). Ions can consist of one atom (monatomic ions) or several (polyatomic ions).Periodic behavior is most evident for ionization energy (I), the energy required to remove an electron from a gaseous atom. The energy required to remove successive electrons from an atom increases steadily, with a substantial increase occurring with the removal of an electron from a filled inner shell. Consequently, only valence electrons can be removed in chemical reactions, leaving the filled inner shell intact. Ionization energies increase diagonally from the lower left of the periodic table to the upper right. Ionic compounds usually form hard crystalline solids with high melting points. Covalent molecular compounds, in contrast, consist of discrete molecules held together by weak intermolecular forces and can be gases, liquids, or solids at room temperature and pressure.

Key Takeaway

• There are two fundamentally different kinds of chemical bonds (covalent and ionic) that cause substances to have very different properties.

Conceptual Problems

1. Ionic and covalent compounds are held together by electrostatic attractions between oppositely charged particles. Describe the differences in the nature of the attractions in ionic and covalent compounds. Which class of compounds contains pairs of electrons shared between bonded atoms?

2. Which contains fewer electrons than the neutral atom--the corresponding cation or the anion?

3. What is the difference between an organic compound and an inorganic compound?

4. What is the advantage of writing a structural formula as a condensed formula?

5. The majority of elements that exist as diatomic molecules are found in one group of the periodic table. Identify the group.

6. Discuss the differences between covalent and ionic compounds with regard to

   1. the forces that hold the atoms together.
   2. melting points.
   3. physical states at room temperature and pressure.

7. Why do covalent compounds generally tend to have lower melting points than ionic compounds?

Answer

7. Covalent compounds generally melt at lower temperatures than ionic compounds because the intermolecular interactions that hold the molecules together in a molecular solid are weaker than the electrostatic attractions that hold oppositely charged ions together in an ionic solid.

Numerical Problems

1. What is the total number of electrons present in each ion?

   1. F^-
   2. Rb⁺
   3. Ce³⁺
   4. Zr⁴⁺
   5. Zn²⁺
   6. Kr²⁺
   7. B³⁺

2. What is the total number of electrons present in each ion?

   1. Ca²⁺
   2. Se^2-
   3. In³⁺
   4. Sr²⁺
   5. As³⁺
   6. N^3-
   7. Tl⁺

3. Predict how many electrons are in each ion.

   1. an oxygen ion with a -2 charge
   2. a beryllium ion with a +2 charge
   3. a silver ion with a +1 charge
   4. a selenium ion with a +4 charge
   5. an iron ion with a +2 charge
   6. a chlorine ion with a -1 charge

4. Predict how many electrons are in each ion.

   1. a copper ion with a +2 charge
   2. a molybdenum ion with a +4 charge
   3. an iodine ion with a -1 charge
   4. a gallium ion with a +3 charge
   5. an ytterbium ion with a +3 charge
   6. a scandium ion with a +3 charge

5. For each representation of a monatomic ion, identify the parent atom, write the formula of the ion using an appropriate superscript, and indicate the period and group of the periodic table in which the element is found.

   1. ${}_4^9{X}^{2+}$
   2. ${}_1^1{X}^{-}$
   3. ${}_8^{16}{X}^{2-}$

6. For each representation of a monatomic ion, identify the parent atom, write the formula of the ion using an appropriate superscript, and indicate the period and group of the periodic table in which the element is found.

   1. ${}_3^7{X}^{+}$
   2. ${}_9^{19}{X}^{-}$
   3. ${}_{13}^{27}{X}^{3+}$

Answers

1. 
4. 

   1. 27
   2. 38
   3. 54
   4. 28
   5. 67
   6. 18

6. 1. Li, Li⁺, 2nd period, group 1
   2. F, F^-, 2nd period, group 17
   3. Al, Al³⁺, 3nd period, group 13

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