Volume 2

Why Are the Noble Gases so Stable? • Every Element Wants to Be Happy. The Octet Rule. Ions • Ionic Bond: "One Man's Trash is Another Man's Treasure" • Metals, Nonmetals, and Metalloids • Metallic Bond: What Metal Atoms Do When There Is No Electron Acceptor Around • Covalent Bond: What Nonmetal Atoms Do When There Is No Electron Donor Around • Lewis Dot Diagrams • Nonpolar and Polar Covalent Bonds. Electronegativity. • Coordinate (Donor-Acceptor) Bond • Hydrogen Bond • A Summary of Chemical Bonds • Exercises
2.4.1. Why Are the Noble Gases so Stable? The noble gases exhibit an extraordinary stability and lack of reactivity. Just 60 years ago, it was widely accepted that no noble gas could react with any other substance and therefore cannot form chemical compounds. In 1962, Neil Bartlett, a British chemist working at the University of British Columbia in Vancouver, Canada, broke the rule by making the first well-defined noble gas compound, XePtF6. Although many more noble gas compounds have been prepared since then, the extreme reluctance of the noble gases to react with other substances remains remarkable.

The reason for the exceptional chemical inertness of the noble gases is simple: their atoms are "happy". The "happiness" of noble gas atoms stems from their outermost (valence) electron shell being completely filled with electrons (Figure 2-39).
Figure 2-39. Electron configuration of the outermost shell of the noble gases.

2.4.2. Every Element Wants to Be Happy. The Octet Rule. Ions. If the ns2np6 arrangement (Figure 2-39) makes a chemical element stable ("happy"), can atoms of elements other than the noble gases attain the same valence electron configuration to gain the desired stability? They can and they do. For instance, an atom of fluorine (F; 1s2 2s2 2p5) readily accepts one electron in its half-filled 2p orbital to get the desired 2s2 2p6 outermost shell configuration of the next element, the noble gas neon (Ne). As a result of this process (Figure 2-40), a highly stable negatively charged species is produced, which is called a fluoride ion. An ion is an atom or molecule that has a net electric charge.
Figure 2-40. A fluorine atom (F) readily accepts one electron to transform into a much more stable fluoride ion (F-) that has the electron configuration of the noble gas neon (Ne).

The stable electron configuration of a noble gas can be achieved by atoms not only by acquiring electrons, but also through loss of electrons. For example, an atom of sodium easily loses one electron to become a much more stable, positively charged sodium ion, whose outermost shell has the electron configuration of neon (Figure 2-41).
igure 2-41. A sodium atom (Na) willingly loses one electron to become a much more stable positively charged sodium ion (Na+) with the electron configuration of neon (Ne).

Negatively charged ions are called anions. Ions bearing a positive charge are called cations. If it is not entirely clear why the fluoride anion is negatively charged and the sodium cation is positively charged, recall that atoms are always electroneutral and that an electron has a negative charge of -1. Adding an electron to an electroneutral atom (or molecule) gives rise to an anion. Removing an electron from an electroneutral atom (or molecule) produces a cation.

Formulas of anions and cations are written exactly like those of non-charged atoms and molecules, except the charge value and its sign are added as a superscript after the formula of the ion. If the charge of an ion is +1 or -1, the "1" is omitted. For example, formulas of the cations produced upon removal of one electron from Li, Na, and K are written as Li+, Na+, and K+. For the dications derived from Mg, Ca, and Zn, the formulas are presented as Mg+2 or Mg2+, Ca+2 or Ca2+, and Zn+2 or Zn2+, respectively. The formula of the triply charged aluminum cation is Al+3 or Al3+. Cations are easy to name: just say the name of the element and add "ion" or "cation" afterwards. For instance, Na+ is "sodium ion" or "sodium cation" and Al3+ is "aluminum ion" or "aluminum cation".

Monoatomic anions are written in a similar manner (for example, F-, Cl-, Br-, S2-), but are named slightly differently. To name a monoatomic anion, just add the suffix "ide" to the stem of the name of the element. Some examples follow.

Fluorine (F) – fluoride (F-)
Chlorine (Cl) – chloride (Cl-)
Oxygen (O) – oxide (O2-)
Sulfur (S) – sulfide (S2-)
Nitrogen (N) – nitride (N3-)

As we know now, the enhanced stability of Na+ vs. Na, F- vs. F, Ca2+ vs. Ca, etc. comes from the exceptionally stable ns2np6 electron configuration of the nth outermost shell. Note that this most stable configuration involves a total of eight electrons (two s and six p), and is consequently often referred to as the octet (from German oktett and Italian ottetto = 8). From this number takes its name the octet rule, stating that atoms tend to pursue, by participation in chemical transformations, the stable eight-electron configuration of the outermost shell.

There is one element that the octet rule does not apply to. This element is hydrogen, of course, because the 1st electron shell can house only up to two electrons, not eight. Sometimes a hydrogen atom accepts one electron to become H-, the hydride ion, which has the electron configuration of helium. Much more often, however, a hydrogen atom loses its only electron to become just a proton. Just a proton? Yes, because an atom of hydrogen (protium, 1H) is made up of only one electron and one proton in the nucleus. Remove the electron from it, and what is left is just the nucleus, which consists of just one proton. That is why chemists prefer to call the H+ "proton" rather than the longer names "hydrogen cation" or "hydrogen ion".

2.4.3. Ionic Bond: "One Man's Trash is Another Man's Treasure". As shown in Figure 2-41, a sodium atom willingly discards its sole electron in the 3s orbital to attain the stable octet configuration 2s2 2p6. Where does this electron go after it leaves the sodium atom? As the electron cannot just disappear completely, it has to go somewhere to find a new home.

Our sodium atom wants to get rid of one electron (Figure 2-41). Meanwhile, a fluorine atom somewhere seeks exactly the opposite, an opportunity to acquire an extra electron in order to gain stability (Figure 2-40). The electron that a sodium atom wants to give away, its "trash", is a "treasure" for the fluorine atom (Figure 2-42). Unsurprisingly, the two enter into a symbiosis, a chemical reaction involving the transfer of one electron from the Na to the F. This reaction produces two highly stable ions, Na+ and F-, which, being oppositely charged, experience mutual electrostatic attraction. This electrostatic attraction holds the ions together, thereby constituting an ionic bond between the Na+ and F-. The process shown in Figure 2-42 is also nicely illustrated by a GIF presented in Video 2-4.
Figure 2-42. To be happy, an atom of Na needs to discard one electron, whereas an atom of F needs to acquire one electron. After the transfer of one electron from the Na atom to the F atom, the Na+ cation and the F- anion produced are electrostatically attracted to each other to form sodium fluoride, NaF, an ionic compound.
Video 2-4. Transfer of one electron from the Na atom to the F atom, followed by the formation of an ionic bond between the resultant Na+ cation and F- anion (source).

Note that the electron transferred from the Na to the F is denoted in Figure 2-42 as e-, a common symbol for an electron in science. Now that we started talking about transfer of electrons in chemical reactions, we will be using this symbol (e-) frequently.

We conclude that an ionic bond is a chemical bond resulting from the mutual electrostatic attraction of oppositely charged ions. Substances featuring this type of bond, called ionic compounds, are ubiquitous in chemistry and are all around us. For example, most salts, including sodium fluoride (NaF) considered above (Figure 2-42) and table salt (sodium chloride, NaCl) are typical ionic compounds.

2.4.4. Metals, Nonmetals, and Metalloids. In Volume 1, we defined metals, nonmetals, and metalloids as simple substances on the basis of their physical properties and visual appearance. It is now time to define metals, nonmetals, and metalloids not only as simple substances, but also as chemical elements (types of atoms).

Metals as simple substances are excellent conductors of electricity and heat, lustrous, ductile, and malleable. A metal as an element is the type of atom that readily forms positively charged ions by willingly giving away its valence electron(s) to attain the stable ns2np6 (octet) electron configuration. A metal atom usually has one, two, or sometimes three (at the most) electrons in its valence shell.

Nonmetals as simple substances may be defined as those that lack "such characteristic properties of metals as hardness, mechanical adaptability, or the ability to conduct electricity." (While in general this is a reasonable definition, we should keep in mind that the alkali metals can be easily cut with a knife, that diamond is harder than any metal, that plastic sulfur is very mechanically adaptable, and that graphite is an excellent conductor of electricity. However, diamond and graphite are allotropes of carbon and plastic sulfur is an allotropic form of sulfur. Both carbon and sulfur are undoubtedly nonmetals.) A nonmetal as an element is the type of atom that is not prone to form cations by giving away its electrons. Atoms of nonmetallic elements have more than three valence electrons.

Metalloids as simple substances often exhibit some features of typical metals such as luster, while ultimately lacking in most characteristic properties of metals, such as high electrical conductivity. There are only a handful of commonly recognized metalloids: boron (B), silicon (Si), germanium (Ge), arsenic (As), antimony (Sb), and tellurium (Te). A metalloid as a chemical element is undefined. "There is no standard definition of a metalloid and no complete agreement on which elements are metalloids." In our course, we will categorize metalloid elements as nonmetals because, like nonmetals and unlike metals, metalloids have three or more electrons in their valence shells and are not inclined to form cations by losing their electrons.

2.4.5. Metallic Bond: What Metal Atoms Do When There Is No Electron Acceptor Around. We know that a metal atom is prone to get rid of its outermost shell electron(s) to become a cation with the stable octet configuration. Where do these electrons go after they leave the metal atoms?

Electrons cannot disappear into nowhere. If there are some nonmetal atoms around, they are happy to accept the electrons from the metal to become anions; it is a win-win situation. But, what happens if there are no species that can accommodate the "extra" electrons that a metal atom wants to discard?

In their "pursuit of happiness", metal atoms have to take care of themselves, and they do. Metals in bulk exist as a lattice of cations immersed in the so-called electron gas that acts as a "glue" to hold the metal ions together. The electron gas is the totality of electrons that have been discarded by the metal atoms and are now evenly spread throughout the entire body of the metallic substance. Figure 2-43 illustrates this metallic bonding model using potassium (K) and calcium (Ca) as two examples.
Figure 2-43. Electron gas models of potassium (K; top) and calcium (Ca; bottom) metals in bulk.

In Figure 2-43, the number of potassium ions (top) and calcium ions (bottom) is the same (twelve). Yet the picture at the bottom (Ca) displays twice as many electrons as the one at the top (K). Why? To answer this question, take a look at the electron configuration of K and Ca atoms (not ions). For potassium, it is 1s2 2s2 2p6 3s2 3p6 4s1 and for calcium 1s2 2s2 2p6 3s2 3p6 4s2. To get to the stable octet of electrons 1s2 2s2 2p6 3s2 3p6, a potassium atom needs to lose only one electron from its 4s orbital, whereas an atom of calcium has to discard two. Consequently, each of the Ca atoms gives rise to a doubly charged cation (Ca2+) and two electrons, whereas each K atom generates the singly charged K+ and one electron (Figure 2-44).
Figure 2-44. To attain the stable octet configuration, a potassium atom discards one electron, whereas a calcium atom discards two.

The metallic bond model accounts for the high mobility of electrons in metals, which makes them excellent conductors.

2.4.6. Covalent Bond: What Nonmetal Atoms Do When There Is No Electron Donor Around. A metal has only one or two valence electrons, which it can easily discard in order to attain the stable electron configuration of the noble gas (of the previous period). A nonmetal cannot gain stabilization in the same way because it has too many (four or more) valence electrons to get rid of. Consequently, a nonmetallic atom has to acquire external electrons from a donor (such as a metal atom) to fill up its vacant orbitals to the valence octet of the noble gas (of the same period).

But, what do atoms of nonmetals do in the absence of an external source of electrons? They share their own electrons in a smart manner to form covalent bonds. Figure 2-45 shows the formation of a covalent bond between two fluorine atoms that come together to give a molecule of fluorine (F2).
Figure 2-45. Orbital overlap and electron sharing between two fluorine atoms. Only the 2p orbitals are shown. The inner 1s orbital and the 2s orbital are omitted for clarity to avoid an overcrowded image.

Each fluorine atom has 7 valence electrons (two in the 2s orbital and five in the three 2p orbitals; 2s2 2p5). The spherical 2s orbital (not shown in Figure 2-45) and two of the three 2p orbitals are completely filled with two electrons each. The third 2p orbital, however, is only half-filled (has only one electron). To attain the desired stable octet, 2s2 2p6, a fluorine atom needs one more electron. When there is no way to get that one much-needed extra electron from anywhere, two fluorine atoms approach each other until their half-filled 2p atomic orbitals overlap to form a new molecular orbital (Figure 2-45). This newly formed orbital has two electrons (one from each of the two original half-filled 2p orbitals) that are now shared by the two fluorine atoms.

After the sharing, each of the two F atoms has the sought-after stable octet of valence electrons:

- Two in the spherical 2s orbital (not shown in Figure 2-45);

- Four in the two 2p orbitals, which are not shared with the other atom; and

- Two that are shared with the other F atom.

The two electrons that belong to both atoms are called a shared electron pair. By sharing the electrons, the two F atoms establish a chemical bond between them. This type of bond involving the formation of a shared electron pair between atoms is called a covalent bond.

A pair of valence electrons in a single orbital that are not shared with another atom is referred to as a lone electron pair. A molecule of F2 has six lone electron pairs (three on each of the two F atoms) and one shared electron pair.

A simpler molecule of hydrogen (H2) is formed from two H atoms in the same way (Figure 2-46). Each of the two hydrogen atoms has one electron in the 1s orbital (1s1). To gain stability by attaining the configuration of He (1s2), the two H atoms form a shared electron pair by overlapping their 1s orbitals.
Figure 2-46. Covalent bond formation between two H atoms.

Figure 2-46 displays two slightly different shape representations of the molecular orbital formed upon the overlap of the two 1s atomic orbitals. Either one is fine. However, to depict a covalent bond, chemists prefer to use so-called Lewis dot diagrams, also known as Lewis structures. We will learn how to draw and use Lewis dot diagrams in the next subsection.

2.4.7. Lewis Dot Diagrams. The idea of a covalent bond has been around for about 100 years since it was introduced by Gilbert N. Lewis (1875-1946), an outstanding American scientist. Lewis also came up with a remarkably simple and clear graphical way to illustrate how the atoms in a molecule are connected by means of covalent bonds and how valence electrons are distributed around atoms.

The rules for drawing a Lewis structure are as follows.

- Atoms are denoted by their chemical symbols;

- Valence electrons of an atom are drawn as dots around its symbol;

- No distinction is made between s and p valence electrons;

- All inner shell (non-valence) electrons are omitted;

- Each covalent bond is drawn as two dots symbolizing the two electrons of the shared electron pair that provides the bonding. Alternatively, each covalent bond can be drawn as a line or, sometimes, as both a line and two dots; and

- Lone electron pairs may be omitted if irrelevant to the context.

As an example, the formation of a molecule of H2 from two hydrogen atoms using Lewis structures is presented in Figure 2-47.
Figure 2-47. The formation of a molecule of H2 from two H atoms, as represented by Lewis dot diagrams.

The formation of a molecule of F2 from two F atoms using Lewis structures can be represented similarly (Figure 2-48). Note that in the structure of F2 on the right in Figure 2-48, all of the lone electron pairs are omitted and the covalent bond between the two F atoms is represented by a line. Compare Figures 2-45 and 2-48 to realize that Lewis structures are at least as informative as orbital shape diagrams, while clearly being much easier and faster to draw.
Figure 2-48. The formation of a molecule of F2 from two F atoms, as represented by Lewis dot diagrams.

Our next example deals with the formation of a molecule of nitrogen (N2) from two nitrogen atoms (N). First, we draw a Lewis dot diagram for a nitrogen atom. The valence shell of a nitrogen atom (2s2 2p3) contains 5 electrons, of which two are paired (2s) and three are single, one in each of the three 2p orbitals (Figure 2-49). Between the two N atoms, there are 6 unpaired 2p electrons, which are used to form three shared electron pairs. The three covalent bonds thereby formed between the two N atoms are called a triple bond.
Figure 2-49. The formation of a molecule of N2 from two N atoms, as represented by Lewis dot diagrams.

2.4.8. Nonpolar and Polar Covalent Bonds. Electronegativity. Covalent bonds are formed not only between atoms of the same type, but also between atoms of different type. For example, a hydrogen atom and a fluorine atom readily form a covalent bond to give a molecule of hydrogen fluoride, HF (Figure 2-50).
Figure 2-50. An atom of hydrogen and an atom of fluorine share their electrons to form a molecule of hydrogen fluoride.

A molecule of ammonia, NH3, featuring three covalent N-H bonds is formed in the same manner from one nitrogen atom and three hydrogen atoms (Figure 2-51).
Figure 2-51. An atom of N and three atoms of H form a molecule of ammonia, NH3, by sharing their electrons.

In terms of the way covalent bonds are formed, there is no difference between those connecting two identical atoms (Figures 2-47, 2-48, and 2-49) or two different atoms (Figures 2-50 and 2-51). It is the sharing of electrons that governs the formation of covalent bonds in all cases. But, is the electron sharing always equal for both atoms forming the bond?

If two equally hungry and considerate individuals are sharing a pizza, each person will get exactly one-half of the meal. But, if one of the two is more selfish while being stronger and hungrier than the other, who do you think will get more? Likewise, the electron pair effecting a covalent bond between two identical atoms such as in H2, N2, and F2 is shared equally. Different elements, however, are "hungry" for electrons to a different degree. Of two covalently bonded different atoms, the one that has a higher affinity (hungrier) for electrons attracts the shared pair more strongly. The tendency of an atom to pull a shared electron pair toward itself is measured by electronegativity. The electronegativity of an element is expressed in numbers that can be found in a special periodic table like the one presented in Figure 2-52. The higher the number, the more strongly the atom attracts electrons toward itself.
Figure 2-52. Electronegativity periodic table (source).

Let us consider three molecules, H2, F2, and HF, whose Lewis structures we have already touched on (Figures 2-47, 2-48, 2-50). Figure 2-53 displays all three, with the covalent bonds depicted as both a line and two dots. While in both the H2 and F2 molecules the shared electron pair is right in between the atoms, in the HF molecule the shared electrons are shifted toward the more electronegative fluorine atom. Since electrons are negatively charged, a partial negative charge appears on the F atom and a partial positive charge on the H atom of the HF. In chemistry, partial charges on atoms are denoted as δ- and δ+ (pronounced delta-minus and delta-plus).
Figure 2-53. Covalent bond electrons are shared equally by identical atoms (H-H and F-F) but not equally by different atoms (H-F).

A covalent bond between identical atoms is referred to as "nonpolar" to indicate that there is no polarization (shifting) of the shared electron pair toward either atom. A covalent bond between atoms of different type and, consequently, of different electronegativity, is referred to as "polar" because the asymmetrical distribution of the shared electrons results in the induction of partial positive and negative charges on the atoms involved in the bonding.

The degree of bond polarization, also known as bond polarity, varies depending on the difference in the electronegativities of the two atoms involved in the bonding. The greater the difference, the more polar the bond is. To evaluate the polarity of a covalent bond, we calculate the difference in the electronegativity values from the table in Figure 2-52.

- If the electronegativity difference is less than 0.5, the bond is weakly polar or almost nonpolar. As an example, the carbon-hydrogen bond in methane, CH4, is weakly polar because the difference in the electronegativity values for carbon (2.5) and hydrogen (2.1) is: 2.5 - 2.1 = 0.4, less than 0.5.

- If the electronegativity difference is greater than 0.5, the bond is polar. For example, the H-F bond in the molecule of HF is polar because the difference in the electronegativity values for fluorine (4.0) and hydrogen (2.1) is: 4.0 - 2.1 = 1.9. In fact, the value of 1.9 is very high, indicating that the H-F bond is very polar, as polar as a covalent bond can get.

- If the electronegativity difference is 2.0 or greater, then the bond is usually ionic. In such a case, one of the two atoms is more electronegative than the other to the extent that it gets the shared electron pair in its full possession to become an anion. Accordingly, the other atom, deprived of its valence electron(s), becomes a cation. This is the case, for example, with sodium fluoride, NaF. The fact that NaF is an ionic compond (Figure 2-42) is consistent with the huge electronegativity difference of 3.1, as calculated using the values of 4.0 for fluorine and 0.9 for sodium (Figure 2-52).

2.4.9. Coordinate (Donor-Acceptor) Bond. This bond is actually a type of covalent bond. The special feature of any coordinate bond is the way it is formed. In the examples of covalent bond formation above, each of the two bonding atoms contributed one electron to the shared electron pair. In the assembly of a coordinate bond, however, both electrons for the shared electron pair come from one of the two atoms. A classic example illustrating this process (Figure 2-54) is the formation of an ammonium cation (NH4+) from a molecule of ammonia (NH3) and a hydrogen ion (proton).
Figure 2-54. The formation of an ammonium cation (NH4+) from a molecule of ammonia (NH3) and a proton (H+).

The nitrogen atom of NH3 has three electron pairs that it shares with the H atoms and one lone electron pair. This lone electron pair is offered for sharing with, and is accepted by, the H+. The newly formed N-H bond is indistinguishable from the other three N-H bonds; after the N atom has shared its lone electron pair with the H+, all four N-H bonds in the ammonium ion (NH4+) become equivalent. These four equivalent N-H bonds of NH4+, however, differ, by a number of parameters, from the original three equivalent N-H bonds of NH3.

A coordinate bond is sometimes called a donor-acceptor bond because of the two bonding atoms one donates its lone electron pair for the covalent bond formation and the other accepts it.

2.4.10. Hydrogen Bond. When a hydrogen (H) atom is bonded to a much more electronegative atom such as fluorine (F) or oxygen (O), particularly strong polarization of the F-H or O-H bond takes place. Consequently, a large partial positive charge (δ+) is induced on the H atom and a partial negative charge (δ-) on the atom of the more electronegative element (Figure 2-55, top). The interaction between the partially positively charged H atom and a lone electron pair on the partially negatively charged atom of the F or O atoms of another molecule of HF or H2O represents a hydrogen bond (Figure 2-55, bottom).
Figure 2-55. Strong bond polarization in HF and H2O (top) leads to hydrogen bond formation (bottom).

Hydrogen bonds are very weak bonds, roughly 10 times weaker than most covalent and ionic bonds. In spite of their low strength, however, hydrogen bonds are crucial for the existence of life. The formula of water, H2O, is arguably the most frequently encountered formula of a chemical compound. However, there are no single H2O molecules in liquid water. At room temperature, water exists as an extended network of hydrogen bonds between individual molecules (Figure 2-56). Experimental studies have demonstrated that hydrogen bonds exist not only in liquid water and ice, but even in steam, above the boiling point of H2O (100 oC). If there were no hydrogen bonds, water would be a gas at room temperature and atmospheric pressure, and there would be no life on Earth.
Figure 2-56. Multiple hydrogen bonds between water molecules.

2.4.11. A Summary of Chemical Bonds. The driving force behind the formation of a chemical bond is the stability of the electron configuration of a noble gas, which atoms of other elements seek to attain. To have its outermost shell completely filled with electrons, an atom may discard electrons, add electrons, or share electrons with other atoms. The three main types of chemical bond emerging from these scenarios are exemplified in Figure 2-57.
Figure 2-57. Examples of covalent nonpolar, covalent polar, and ionic chemical bonds. Lone electron pairs are omitted.

The sharing of an electron pair gives rise to a covalent bond. When an electron pair is shared by two identical atoms, the covalent bond is nonpolar because atoms of the same element attract electrons equally (Figure 2-57, left). If two different atoms are connected by a covalent bond, the electron pair is shifted toward the more electronegative atom and away from the less electronegative atom (Figure 2-57, center). This induces a partial positive charge on the less electronegative atom of the two and a partial negative charge on its more electronegative partner. Finally, if the electronegativity difference is large enough, the electron pair is no longer shared but rather taken away by the more electronegative atom. As a result, an ionic bond is formed, which may be viewed as the farthest extreme of a covalent polar bond with the shared electron pair entirely shifted toward one of the two atoms.

2.4.12. Exercises.

1. The driving force behind the formation of a chemical bond is (a) to get as many electrons as possible for atoms of all of the elements involved; (b) to gain stability by attaining the valence electron configuration of a noble gas; (c) to balance the number of negatively charged electrons and positively charged protons to make the system electroneutral. Answer

2. Write the electron configuration for a fluorine atom, a neon atom, and a sodium atom. Which of these three atoms is most stable? How can the other two gain similar stability? Ditto for Cl, Ar, and K.

3. What is an ionic bond?

4. Describe the formation of MgCl2, an ionic compound, from one Mg atom and two Cl atoms. [Tip: See Figure 2-42]

5. Define a metal and a nonmetal element from the perspective of their electron configuration. [Answer: See subsection 2.4.4 above]

6. What is the structure of metals in the bulk? What is the electron gas and how is it formed? Sketch a diagram describing the structure of Na metal. [Tip: See subsection 2.4.5 above]

7. Define a covalent bond. What are the conditions for nonpolar and polar covalent bonds?

8. Draw Lewis dot diagrams for the following atoms, molecules, and ions: Li, C, O, Ne, P, Cl, Cl-, Cl2, Ca, Br, HCl, LiF, H2O, NF3, NH4+.

9. Using electronegativity values in Figure 2-52, identify the chemical bonds as weakly polar covalent, polar covalent, or ionic for the following compounds: BaCl2, CO2 (O=C=O), CS2 (S=C=S), CH4, BH3, BCl3, BaO, CsBr. Draw Lewis dot structures for the covalent compounds. Mark the atoms bearing partial negative and positive charges with the δ- and δ+ symbols.

10. Draw a scheme describing the formation of a hydronium cation, H3O+, from a molecule of water and a proton. [Tip: Refer to Figure 2-54. It is exactly the same idea]

11. A hydrogen bond is usually (a) slightly weaker than ionic and covalent bonds; (b) stronger than covalent bonds but weaker than ionic bond; (c) much weaker than covalent and ionic bonds; (d) about 10 times stronger than ionic and covalent bonds. Answer

12. It is sometimes said that without hydrogen bonds there would be no life on Earth. Why? Draw a sketch illustrating hydrogen bonds between water molecules.

13. A molecule of water is formed by overlap of each of the two singly occupied 2p orbitals of O with the 1s orbital of H, as shown below. Given that all three dumbbell-shaped p-orbitals are perpendicular to one another, the H-O-H angle in an individual water molecule would be expected to be 90o.
Experimental measurements, however, show that the H-O-H bond angle is more obtuse, 105o. Can you come up with an explanation for this fact? Answer