Chemistry: A Self-Teaching Guide - Post R., Snyder C., Houk C.C. 2020

Periodic Properties and Chemical Bonding

In Chapter 1 you learned that the elements in a horizontal row of the periodic table show regular variation in properties from left to right. The elements are arranged in the table in order by increasing atomic number (reading the table left to right, line by line, in the way you are reading this paragraph). The reason this arrangement works so well is that all atoms consist of electrons, protons, and neutrons. The neutrons and protons are in the nucleus with electrons arranged in “shells” around the nucleus.

Why did we consider the electronic arrangement of atoms in such detail? Because the chemical properties of an element depend upon the number of electrons in its outermost shell, the energy levels of its outermost electrons, and the size of the atom. These details of atomic structure determine what kinds and how many chemical bonds can be formed by an atom.

In this chapter we discuss several properties not mentioned in Chapter 1 that depend upon the outermost shell electronic structure. We will review electron configuration and introduce new “dot” symbols. We will then discuss whether atoms gain, lose, or share electrons, and how many, when they combine to form new substances. The major portion of the chapter is devoted to the types of chemical bonds (ionic, covalent, polar covalent) formed between atoms in chemical compounds. Finally, we look briefly at the shapes of compounds and ionization energy.

OBJECTIVES

After completing this chapter, you will be able to

· recognize and apply or illustrate: ion, ionic, covalent, coordinate covalent, molecule, compound, electron dot symbol, formula, electronegativity, polar, ionization energy, metallic character, the octet rule, nonmetallic character, and valence shell electrons;

· write electron dot symbols that represent the number of valence shell electrons for the Group A elements and noble gases;

· write the symbol for the ion formed by any of the Group A elements;

· write equations using dot symbols to show the formation of ionic, covalent, and coordinate covalent compounds;

· predict the type of bond (ionic, covalent, polar covalent) formed when any two elements combine when given the electronegativity of the elements;

· use the octet rule to choose a correct dot structure for a compound;

· predict whether a molecule containing three or more atoms would be polar or nonpolar if given its shape;

· predict whether an element, according to its position in the periodic table: (1) is metallic or nonmetallic, (2) has a high or low electronegativity, (3) has a high or low ionization energy, and (4) will gain or lose electrons and how many.

OUTER SHELL ELECTRONS

The following chart represents the first 20 elements of the periodic table. Only the symbols and atomic numbers have been included.

The atomic number increases by one each time as the elements are viewed from left to right (11 to 12 to 13 to 14 to 15, and so on). Each time the atomic number increases by one, how many electrons are added?_______

Answer: one (Remember that the atomic number represents the number of protons. The symbols in the periodic table represent atoms of neutral elements. In a neutral atom, the number of protons equals the number of electrons.)

Periodic tables can be made to include different kinds of information. The following table includes the electronic structure for the outermost shell for each of the first 20 elements.

There are similarities in the structure of elements within groups. Notice the structure of each outermost shell in Group IA. How many electrons are present in the outer shell of each element in Group IA (H, Li, Na, K, and others)? ________

Answer: one (Each element in Group IA has only one electron in its outermost shell.)

Use the periodic table in frame 2 for the following statements.

1. Group IIA elements each have two electrons in their outermost shell. Group IIIA elements each have ________ electrons in their outermost shell.

2. Group IVA elements each have ________ electrons in their outermost shell.

3. Group VA elements each have five, Group VIA elements each have ________, and Group VIIA elements each have ________ electrons in their outermost shell.

4. Group VIIIA elements (the noble gases) each have eight electrons in their outermost shell with the exception of helium, which has only ________electrons.

Answer: (a) three; (b) four; (c) six, seven; (d) two

Nitrogen (N), a Group VA element, has ________ electrons in its outermost shell. Aluminum (Al), a Group IIIA element, has ____ electrons in its outermost shell.

Answer: five; three

The outer shell electrons are also known as valence electrons. The periodic table on the next page includes all of the symbols for the first 20 elements. The numbers of valence electrons are listed for lithium (Li), carbon (C), and argon (Ar). Fill in the number of valence electrons for each remaining element in this periodic table.

Remember that helium (He) is an exception in Group VIIIA, since it only has a total of two electrons in its entire atom, both of which are valence electrons.

Answer:

The outer shell or valence electrons are especially important because those electrons are involved when atoms unite chemically to form compounds.

When an atom of magnesium unites with another different atom to form a compound, what electrons of the magnesium atom are primarily involved?

Answer: the two valence or outer shell electrons

ELECTRON DOT SYMBOLS (LEWIS SYMBOLS)

The outer shell or valence electrons may also be represented by a series of dots. Beryllium, with two valence electrons, can be represented as Be: (each dot represents one valence electron).

represents a chlorine atom with how many valence electrons? ____

Answer: seven

The first 20 elements and their electron dot symbols are as follows.

Note that some books represent boron as , aluminum as , carbon as , and silicon as . Such a representation is used to simplify the explanation of how many bonds a given element may have or may form with other elements, but it does not agree with the quantum mechanical concept of the atom (Chapter 1).

We will discuss the number of bonds an element has in a compound in Chapter 14, Organic Chemistry. Later in this chapter we discuss a type of bonding that requires agreement with what is called the octet rule and involves pairs of electrons, so we have chosen to use the scheme shown in this table.

Which elements in the table have five valence electrons? _______

Answer: and

Notice in the table that, as we go from left to right, the first two electrons pair up, while all additional electrons remain single until all four sides of a symbol are occupied. The symbol :Mg is the same as or Mg: or . The placement of the dots is not critical as long as the proper number of electron pairs and unpaired electrons are represented. The symbol is the same as or any arrangement of one pair and one single electron. Note that is not correct since one pair and a single electron are called for, not three single electrons.

Nitrogen is pictured in electron dot symbols as with one pair and three single electrons. Which of the following are also correct symbols for nitrogen?

________

Answer: and (since they show three single electrons and one pair)

The dot symbols showing the valence electrons for the third period elements of the periodic table are as follows.

Remember that the first two electrons pair up while all additional electrons remain single until all four sides of a symbol are occupied. The first pair of electrons represent an s subshell while the next six electrons represent a p subshell. By the Principle of Maximum Multiplicity, electrons in a subshell prefer to remain unpaired until each orbital contains one electron. The dot diagrams reflect this principle.

The element arsenic (As) has five valence electrons in its outermost shell. The outermost shell consists of two electrons in an s subshell and three electrons in a p subshell. Draw an electron dot symbol for As. _______

Answer: (or any arrangement with one pair and three single electrons)

The element tellurium (Te) has six valence electrons. Its outermost shell consists of a completed s subshell and four electrons in a p subshell. Draw an electron dot symbol for Te. _______

Answer: (or any arrangement with two single electrons and two pairs of electrons)

At this point, you may be wondering how these electron dot symbols can be used. You already know that the dots represent valence electrons and that the valence electrons are especially important because they are involved when two or more atoms combine chemically to form compounds. The electron dot symbol arrangement changes when atoms combine to form compounds.

Before you can use the dot symbols to show compounds, however, you must learn to distinguish between two major categories of chemical compounds.

These two categories, ionic and covalent, will be introduced next. The valence electron arrangements for these two kinds of compounds are different, but both arrangements can be shown through the use of dot symbols.

IONS

An ion is an atom or group of atoms that is no longer neutral. The numbers of electrons and protons in an ion are not equal. You have learned that a proton has a positive (+) charge and an electron has a negative (−) charge. A neutral atom, in contrast, has an equal number of + and − charges, with the net result that the charges cancel each other. The number of protons (atomic number) remains unchanged when an atom becomes an ion, but one or more electrons may be gained or lost. This results in an ion with either fewer or more electrons than protons.

1. Oxygen, with an atomic number of 8, was found to have 10 electrons. Is it an ion or a neutral atom? _______

2. Aluminum, with an atomic number of 13, also was found to have 10 electrons. Is it an ion or a neutral atom? _______

3. Magnesium, with an atomic number of 12, was found to have 12 electrons. Is it an ion or a neutral atom? _______

Answer: (a) ion; (b) ion; (c) neutral atom

The symbol for an ion is determined as follows, using two of the examples from frame 12.

An oxygen ion with 10 electrons and an atomic number of 8 has gained two extra electrons. Since each electron has a negative (−) charge, the oxygen ion has a 2− charge. The 2− is written to the upper right of the (O) symbol to indicate two extra electrons, O²⁻. The aluminum ion has 10 electrons and an atomic number of 13. It has lost three electrons, and its symbol is Al³⁺. The 3+ to the upper right of Al indicates that the protons outnumber the electrons by three.

The calcium ion (Ca²⁺) has 20 protons (same as atomic number) and _______ electrons.

Answer: 18 (It has lost two electrons; therefore, the protons outnumber the electrons by two.)

A bromide ion has gained one electron. It is symbolized by Br⁻. A sodium ion has lost one electron. It is symbolized by Na⁺. (Note that the number 1 is implied and is not written as part of the symbol for an ion.)

A lithium ion is symbolized as Li⁺. It has (lost, gained) _____ electrons. How many? ______

Answer: lost; one

If copper (Cu) has lost two electrons, write the symbol for its ion.   _____

If sulfur (S) has gained two electrons, write the symbol for its ion.   _____

If fluorine (F) has gained one electron, write the symbol for its ion. _____

Answer: Cu²⁺, S²⁻, F⁻

An ion is formed when a neutral atom gains or loses one or more electrons. You may be wondering why an atom would gain or lose electrons to form an ion. As a general rule, atoms tend to form sets of eight electrons in their outermost shells (valence electrons) when they combine to form compounds. Atoms with six or seven valence electrons have a great attraction for one or two extra electrons, and atoms with only one, two, or three valence electrons have only a weak hold on those electrons. An atom with six or seven valence electrons will tend to gain electrons from another atom with one, two, or three valence electrons when those atoms combine chemically.

Calcium (Ca) has two valence electrons. Sulfur (S) has six valence electrons. If an atom of calcium and an atom of sulfur are combined chemically, which atom would gain electrons and which atom would lose electrons? ________________

Answer: Sulfur would gain electrons while calcium would lose electrons.

By gaining two electrons, the sulfur atom, which had six valence (outer shell) electrons, becomes an ion with a total of eight outer shell electrons. The ion symbol is S²⁻. By losing two valence electrons, the calcium atom becomes a positive ion symbolized by Ca²⁺. The general tendency of atoms to form sets of eight electrons in their outermost shells (valence electrons) when they combine to form compounds is commonly called the octet rule.

1. When a calcium ion and a sulfur ion combine to form a compound, the sulfur ion (S²⁻) has a total of how many electrons in its outermost shell? _______

2. According to what rule? _______

Answer: (a) eight; (b) the octet rule

The noble gases (Group VIIIA in the periodic table) already have eight electrons in their outermost shells (with the exception of helium). When atoms combine to form compounds, the atoms generally tend to form electron configurations that are similar to those of the noble gases.

A neutral bromine (Br) atom has 35 electrons in its electron configuration. Its total electron configuration is 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁵. Its outermost shell configuration (the fourth shell) is represented by 4s² 4p⁵, which adds up to 2 + 5 = 7 outer shell (valence) electrons. When combining chemically with another atom such as sodium, the bromine atom will gain one electron to form the bromide ion (Br⁻).The bromine atom (Br) has an electron configuration of 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁵. The bromide ion (Br⁻) has an electron configuration of 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁶.

The bromine atom has seven outer shell electrons and a total of 35 electrons. The bromide ion, however, has how many outer shell electrons? __________ How many total electrons? _________

Answer: eight; 36

Remember that when atoms combine to form compounds, the atoms generally tend to form electron configurations that are similar to those of the noble gases. When a bromine atom becomes an ion, it gains one electron. Instead of having 35 electrons in its total electronic structure, it gains one electron to have a total of 36 electrons. The resulting bromide ion has attained an electron configuration that is similar to that of a noble gas atom.

What neutral noble gas atom has a total of 36 electrons? (Hint: Use the periodic table for your answer.) _________

Answer: Krypton (Kr) has an atomic number of 36 and, therefore, has 36 electrons in its neutral atom.

When chlorine (Cl) and sodium (Na) combine to form NaCl, both atoms become ions. Chlorine gains one electron from sodium to become the Cl⁻ ion. The sodium atom loses its one outer shell electron to chlorine and becomes the Na⁺ ion.

The sodium atom originally had 11 electrons in its electron configuration (1s² 2s² 2p⁶ 3s¹). In becoming an ion, sodium loses its only outer shell electron (3s¹) and becomes the Na⁺ ion with a total of 10 electrons. What neutral atom has the same number of electrons as the sodium ion (Na⁺)? __________ To what family of elements does that neutral atom belong? (Use the periodic table.) _________

Answer: Ne (neon) (has 10 electrons); noble gas family

All of the elements in Group VIIA of the periodic table (F, Cl, Br, I, At and Ts) have outer shell structures with seven electrons. Each element in Group VIIA has just one less electron than the corresponding noble gas in each period and each, when forming an ion, tends to gain one electron. You are already familiar with the chloride ion, which is symbolized by Cl⁻. The single minus sign to the upper right of the Cl symbol indicates a gain of one electron.

Write the symbol for the ion of each of the following neutral atoms in Group VIIA. Assume that each ion has gained one electron.

Answer: Ions: F⁻, Cl⁻, Br⁻, I⁻, At⁻

An atom of each element in Group VIA has six electrons in its outermost shell. In order to become ions with eight electrons, these atoms must gain two electrons. You are already familiar with the symbol for an oxygen ion (O²⁻). The symbol represents two electrons gained.

Write the symbols for the ions of each of the following neutral atoms from Group VIA. Each has gained two electrons. (The first has been filled in.)

Answer: Ions: O²⁻, S²⁻, Se²⁻, Te²⁻, Po²⁻

An atom of each element in Group IA has only one electron in its outermost shell. An atom of each element in Group IIA has two electrons in its outermost shell. These atoms readily give up their outer shell electrons to become positive ions. You are already familiar with the sodium ion (Na⁺), which has lost one electron, and the calcium ion (Ca²⁺), which has lost two electrons.

Write the symbols for the ions of each of the following neutral atoms from Group IA and Group IIA. Use the periodic table to determine whether an atom is from Group IA or Group IIA.

Answer: Ions: Ca²⁺, K⁺, Ba²⁺, Li⁺, Mg²⁺, H⁺

Many of the elements in groups other than IA, IIA, VIA, and VIIA can commonly form ions with more than one value. In some cases the octet rule is followed, and in some cases it is not. For example, the element copper (Cu) can form Cu⁺ as well as Cu²⁺. In other words, an atom of copper can lose either one electron or two electrons to form two different kinds of ions. When such ions are encountered in this and later chapters, you will be informed of the number of electrons lost or gained or be given the proper symbol for the ion.

Use the periodic table to write the proper ion symbol for each of the following atoms from Groups IA, IIA, VIA, or VIIA.

Answer: Ions: Br⁻, S²⁻, Be²⁺, Rb⁺, Na⁺

A zinc (Zn) atom loses two electrons to become an ion. An iron (Fe) atom loses three electrons to become an ion. Write the proper symbols for these ions.

Answer: Zn²⁺, Fe³⁺

Metals generally form positive ions (ions with a + charge) while nonmetals generally form negative ions (ions with a − charge). Na⁺, Ca²⁺, and Fe³⁺ are some examples of metallic ions, and Cl⁻, O²⁻, and Br⁻ are examples of nonmetallic ions.

Metals, in forming ions, would generally be expected to (lose, gain) _______ electrons. Nonmetals, in forming ions, would generally be expected to (lose, gain) _______ electrons.

Answer: lose; gain

IONIC BONDS

Compounds can be formed from the combination of a negative ion and a positive ion. Such compounds are called ionic compounds. The negative and positive ions of the compound are held together by an ionic bond (sometimes called an electrostatic bond). You are already familiar with the ionic compound NaCl, which is made up of two ions with opposite charges. One ion is positive and metallic. The other ion is negative and nonmetallic. Write the appropriate symbols for the ions that make up NaCl. ___________________

Answer: Na⁺ and Cl⁻

An ionic (electrostatic) bond involves oppositely charged ions held together by the attraction from their opposite electrical charges. (Oppositely charged particles attract.) Circle those pairs of ions that could probably form ionic compounds.

Answer: K⁺ and F⁻; Mg²⁺ and O²⁻ (The other pairs of ions have like charges and will not form ionic bonds, since like charges repel.)

In an ionic bond, the two ions are of opposite charge. When a potassium (K) atom and a fluorine (F) atom combine to form a compound, the F atom with only seven valence electrons gains one electron from the K atom, which has only one valence electron to lose. The result is the K⁺ ion is united in an ionic bond with the F⁻ ion. Both ions have achieved an electron configuration similar to that of the noble gases and are held together by their opposite electrical charges.

1. The metallic ion is (K⁺, F⁻)_______.

2. The nonmetallic ion is (K⁺, F⁻)______________.

3. The K⁺ ion has the same number of electrons as which noble gas atom? (Use the periodic table.) ___________

4. The F⁻ ion has the same number of electrons as which noble gas atom? (Use the periodic table.) ___________

Answer: (a) K⁺; (b) F⁻; (c) Ar (argon) (with 18 electrons); (d) Ne (neon) (with 10 electrons)

When the formula (set of symbols representing a compound) for an ionic compound is written, the most metallic element is written first. In the case of NaCl, it should be clear that Na is much more metallic than Cl (since Na is a metal and Cl is a nonmetal). As you may remember, the most metallic elements are located on the left side of the periodic table and the most nonmetallic elements are located on the right. Which of the following ionic compounds are correctly written with the most metallic element first?

MgO, FK, LiBr, OCa ___________

Answer: MgO and LiBr (The other two are incorrect. Since K is more metallic than Br, it should be written as KF, and since Ca is more metallic than O, the formula is written as CaO.)

Up to this point, you have dealt only with compounds made of ions with opposite but equal charges. Suppose we wish to make an ionic compound with Ca²⁺ and Cl⁻ ions. In this case, the calcium (Ca) atom has two electrons to lose to form an ion, but the chlorine (Cl) atom only needs to gain one electron to form an ion. We can form the compound by using two Cl⁻ ions for each Ca²⁺ ion. The result is an ionic compound written as CaCl₂. The number 2 written at the lower right side (subscript) of Cl in the formula indicates that there are two Cl⁻ ions for each Ca²⁺ ion. Write the formula for the compound made up of the ions Mg²⁺ and Br⁻. _______

Answer: MgBr₂ (This requires two Br⁻ ions for each Mg²⁺ ion.)

The formula Li₂O represents an ionic compound. Based upon what you have just learned, how many Li⁺ ions are necessary for each O²⁻ ion? _______

Answer: Two Li ⁺ ions are necessary for each O²⁻ ion.

Write the formulas for ionic compounds that are made up of the following sets of ions.

1. Ca²⁺, I⁻ ___________

2. O²⁻, K⁺ ___________

3. Na⁺, S²⁻ ___________

4. Fe³⁺, Cl⁻ ___________

Answer: (a) CaI₂; (b) K₂O; (c) Na₂S; (d) FeCl₃

The next few frames discuss a second type of bonding that occurs when atoms share electrons instead of actually exchanging them.

COVALENT BONDS

In the formation of an ionic (electrostatic) bond, one or more valence electrons are removed from one atom and taken by another atom, and in the process both atoms become ions. The resulting atoms are bonded together by the attraction of their opposite charges.

A different type of bond is the covalent bond. In a covalent bond, one or more valence electrons from each atom are shared. The resultant compound is made up of molecules, not ions.

Carbon monoxide exists as a molecule composed of carbon and oxygen with shared outer shell electrons. Carbon monoxide uses what kind of bonding, ionic or covalent? ___________

Answer: covalent

Covalent bonds are usually formed between nonmetals. Ionic bonds are usually formed between a metal and a nonmetal. The bond between nitrogen and oxygen in the compound NO would most likely be (ionic, covalent) ___________.

Answer: covalent (because both atoms are nonmetals, on the right side of the periodic table)

Carbon tetrachloride is composed of one carbon and four chlorine atoms. Both carbon and chlorine are nonmetals. What type of bonding is probably involved in this compound? ___________

Answer: covalent

Carbon tetrachloride, a compound with covalent bonding, is made up of (ions, molecules) ___________.

Answer: molecules

Which type of bonding (ionic or covalent) could more likely be expected in a compound whose atoms:

1. are all nonmetals? ___________

2. share valence electrons? ___________

3. have become ions? ___________

4. have transferred valence electrons from one atom to another? ___________

Answer: (a) covalent; (b) covalent; (c) ionic; (d) ionic

Electron dot symbols can be used to represent both ionic and covalent bonding. The dot symbols below represent the ionic bonding of KI (potassium iodide). A potassium (K) atom transfers its single valence electrons to an iodine (I) atom, and both become ions. The brackets, [ ], show that all eight electrons in the outer shell now belong to the negative ion. The comma placed between the negative and positive ions is there only to separate the two ionic symbols.

The following dot symbols represent covalent bonding. Two fluorine atoms share a pair of electrons to become one fluorine gas molecule.

What type of bonding do the following symbols represent?

___________

Answer: ionic bonding (The two valence electrons from Ca were transferred to the O atom, and both atoms become ions.)

Assume ionic bonding for the following reaction. Complete the dot symbols equation. The strontium atom gives up both valence electrons to the sulfur atom, and the atoms become ions.

Answer:

Ionic bonding for gallium iodide, GaI₃, is represented as follows.

Note that three iodine atoms each take one electron from the three available valence electrons of gallium. The number 3 placed in front of the brackets indicates three iodide ions each having a charge of 1−.

Complete the ionic bonding equation for MgCl₂ in which two chlorine atoms each take one electron from the two available valence electrons of magnesium.

Answer:

Look at the following ionic bonding equations.

In these equations, how many total outer shell electrons does each negative ion have? (Count the electrons around the negative ion in brackets.) ___________

What rule would lead us to expect this? ___________

Answer: eight; the octet rule

All noble gas atoms, with the exception of helium, have eight outer shell electrons. In following the octet rule, bonding atoms tend to form an outer shell electron configuration similar to that of the noble gases. Fluorine (F₂) is composed of two fluorine atoms and is covalently bonded, as indicated in the following bonding equation.

Including the shared electron pair, how many electrons are in the outer shell of each atom? ___________

Answer: eight (Note that one pair is shared by both atoms.)

Each circle contains eight electrons.

The correct bonding equation for N₂ must be written so that each N atom has eight electrons, following the octet rule. More than one pair of electrons can be shared.

Knowing that each N atom within the N₂ molecule must have eight electrons, which of the following is the correct outer shell electron structure for N₂? (You may wish to draw circles around each atom to decide your answer. Be sure to circle all of the shared electrons around each atom.)

Answer: (This structure showing three pairs of shared electrons — a triple bond — is the only structure that allows each atom access to eight electrons.)

Each circle contains eight electrons.

The others are not correct, as shown by the number of electrons in each circle.

Select the proper electron dot structure for the outer shell electrons of the O₂ molecule. Follow the octet rule.

Answer: (These structures give each atom eight electrons. Two pairs of electrons must be shared to give each atom eight electrons.)

Which of the following is the correct representation of the valence electrons of the carbon dioxide molecule? Four covalent bonds are included in the molecules.

Answer: (This is the only structure that allows each atom eight electrons, no more or less.)

The H₂ molecule electron structure shows that each hydrogen atom prefers the helium noble gas structure with two shared electrons.

Determine the electron outer shell structure of the HCl gas molecule. The hydrogen atom prefers the helium structure and the chlorine atom follows the octet rule.

Answer: (H has two electrons and Cl has eight electrons.)

Not all compounds follow the octet rule, but it serves as a useful rule for showing most electronic bonding structures. The general tendency in bonding is for atoms or ions to attain the noble gas outer shell structure.

With the exception of helium, the noble gases have how many electrons in their outer shells? ___________ Helium has how many electrons in its outermost (and only) shell? ___________

Answer: eight; two

In a regular covalent bond, a shared pair of electrons originates when each of two atoms supplies one electron. A special type of covalent bond is known as the coordinate covalent bond. In a coordinate covalent bond, both electrons in a shared pair of electrons come from the same atom. The bonding between the nitrogen atom in ammonia (NH₃) and the boron atom in boron trifluoride (BF3) is an example of a coordinate covalent bond.

In the above equation, which atom supplied both electrons to form the coordinate covalent bond, H, N, B, or F? ___________

Answer: The N atom supplied both electrons to form a coordinate covalent bond. (Notice that in NH₃ on the left side of the equation, the N is already shown with eight electrons.)

A coordinate covalent bond is like any other covalent bond except for the origin of the shared electron pair. Does the molecule below conform to the octet rule (with the exception of the hydrogen atoms)? ___________

Answer: Yes, all the atoms in the molecule (except the hydrogen atoms) have eight outer shell electrons.

Review your understanding of ionic, covalent, and coordinate covalent bonding.

1. A bond in which one or more electrons are removed from one atom and taken by another is called a(n) ____________ bond.

2. A bond in which two atoms share a pair of electrons with one electron coming from each atom is called a(n) ____________ bond.

3. A bond in which two atoms share a pair of electrons with both electrons coming from one atom is called a(n)____________ bond.

Answer: (a) ionic; (b) covalent; (c) coordinate covalent

Up to this point we have assumed that the electrons in a covalent bond are shared equally between the bonding atoms, that is, the electrons are midway between the bonding atoms. In the frames that follow we discuss what happens if the electrons are not equally shared, that is, if the electron pair is closer to one atom than the other.

POLAR BONDS

When all the atoms in a covalent molecule are identical, the attraction each atom has for shared electrons is equal. When the atoms in a covalent molecule are different, the attraction for the shared electrons is not equal. In a covalent molecule such as Cl₂, the attraction for the shared electron pair by both atoms within the molecule is equal. This is obvious since the two chlorine atoms are exactly the same. A molecule of HCl gas is also covalently bonded, but the attraction for electrons by the H and Cl atom is not the same. The Cl atom has a greater attraction for electrons than the H atom.

Which of the following molecules (all have covalent bonds) should have atoms with equal attraction for electrons H₂, F₂, HF, NH₃? ___________

Answer: H₂ and F₂ are molecules with an identical pair of atoms that equally attract electrons.

The attraction an atom has for the electrons in a covalent bond is called electronegativity. In the HCl molecule, the shared pair of electrons spend more time nearer the chlorine atom. In the HF molecule, the shared electrons spend more time nearer the F atom. Based on this information, which of the two atoms would be more electronegative, H or F? _________

Answer: F is more electronegative because it attracts the electron pair more than H.

A molecule that has atoms with differing electronegativities has a slight + charge on one side of the molecule and a slight − charge on the other side because the shared electron pair is located nearer the more electronegative atom. When this happens, the molecule has “poles” (like the earth) or regions of + and − charge and is said to be polar. Later, in Chapter 11, you will see how polarity affects the properties of a compound. HCl is a covalent molecule with a slight − charge on the chlorine side and a slight + charge on the hydrogen side of the molecule. Is HCl a polar molecule? ___________

Answer: yes (It has regions of slight + and − charges in the molecule.)

HCl is a polar molecule. The Cl atom has a greater electronegativity than the H atom. As a result, the bonding electron pair is shared unequally and spends more time near the Cl than the H. The HCl molecule therefore has a slight negative charge on the Cl side and a slight positive charge on the H side. The slight negative charge is equal but opposite the slight positive charge. Based upon what you have learned so far, which of the following covalent molecules cannot be polar molecules: HF, H2, F2, NH₃? ___________

Answer: H₂ and F₂ are definitely not polar molecules because the atoms are identical. HF is polar (frame 53), but at this time you have insufficient information to judge NH₃.

HBr is a polar covalent molecule. The shared bonding electrons are nearer the Br than the H. Since electrons are negative charges, which of the following correctly describes this polar molecule?

1. Slightly negative H and slightly positive Br ___________

2. Slightly positive H and slightly negative Br ___________

Answer: (b)

Iodine is more electronegative than hydrogen, so HI is a polar molecule. Which atom is slightly negative? _______ Which is slightly positive? _____

Answer: I; H

The small Greek letter δ (delta) is sometimes used to show a partial positive or negative charge on a polar covalent molecule. For example, the HCl molecule could be symbolized as:

The partial positive and negative charges are indicated near the proper atoms.

Fluorine is more electronegative than hydrogen. Which of the following symbols is correct?

Answer: (a)

Carbon monoxide (CO) is also a polar covalent molecule. The structure for CO is

Oxygen is more electronegative than carbon. Which of the following is correct?

Answer: (a)

Note that the partial + and − charges in a polar covalent molecule are much smaller than the unit + or − charges of a proton or electron. The polar charge is always much less than the charges that hold ions together in an ionic compound. Which type of bonding has the greater charge difference holding the atoms together: ionic, polar covalent, or covalent? _______

Answer: ionic

For compounds consisting of only two atoms, it is possible to classify the bond(s) as covalent, polar covalent, or ionic based upon the electronegativity difference between the atoms. Electronegativity is abbreviated “en,” and the electronegativity difference is shown as Δen. We will use this scheme of classification. The compound is:

1. covalent if Δen is less than 0.5 en units;

2. polar covalent if Δen is between 0.5 and 1.5 en units;

3. ionic if Δen is greater than 1.5 en units.

The purpose of such a classification is to provide a rough rule of thumb for predicting the type of bond that might form when two atoms combine. To be sure, any difference in electronegativities will result in a polar bond. However, it is the degree or extent of the difference that will permit you to tell whether one compound is more or less polar than other compounds, or whether a particular bond is more or less polar than another or more ionic than covalent. The only time we have a 100% covalent bond is when Δen = 0, which usually happens only when identical atoms combine. Using the table shown below and the scheme above, CO would be classified as a polar covalent compound.

The electronegativity value is the number below the symbol in the table, C = 2.5 and O = 3.5 units. The difference (Δen) is 3.5—2.5 = 1.0 unit, therefore CO is polar covalent.

Electronegativities of the First 20 Elements

How should KF be classified: ionic, polar covalent, or covalent?

Answer: Δen = 4.1 (fluorine) − 0.9 (potassium) = 3.2, thus it is an ionic bond

Using the table again, how should each of the following be classified?

1. HCl __________

2. LiCl __________

3. F₂ ____________

Answer: (a) polar covalent (Δen = 0.8 unit); (b) ionic (Δen = 1.9 units); (c) covalent (Δen = 0.0; they are identical atoms)

SHAPES OF MOLECULES

The classification scheme discussed in frames 61 and 62 works only with compounds composed of two atoms. If we have compounds of three or more atoms, the shape of the molecule becomes important in determining its polarity.

A discussion of the shapes of molecules requires an exhaustive look at atomic and molecular orbitals and their shapes and spatial orientation, along with a discussion of the valence bond electron repulsion theory. Such a presentation is beyond the scope of this book. We will limit our discussion to a few simple molecules that are known to be linear (all the atoms in the compound can be joined by a single straight line through their nuclei) or bent (nuclei cannot be joined by a straight line). Bent molecules may be planar (all the atoms fall within the same flat surface, such as this page) or nonplanar (one or more atoms fall outside the flat surface — say, above or below the plane of this page).

Carbon dioxide (CO) is a linear molecule with the structure shown below.

When a carbon atom and an oxygen atom form a bond or bonds between them, will the electrons be closer to the carbon atom or the oxygen atom?

Answer: oxygen (because it is more electronegative)

Carbon dioxide is a linear molecule. Oxygen (1) is pulling on the electrons between it and the carbon to the same extent as oxygen (2) is pulling on its electrons.

An analogy could be two tug-of-war teams pulling on a rope with neither team moving. Even though each bond within the molecule may be polar, the net result, because the molecule is linear, is that the molecule as a whole is nonpolar.

Carbon disulfide (CS₂) has a dot structure of:

Would you expect the molecule to be polar or nonpolar? _______

Answer: nonpolar

Water (H₂O) is an example of a bent planar molecule. Its line structure is:

The “tugs” in this case are not along the same straight line so they do not offset each other. Would you expect H₂O to be a polar or nonpolar molecule?

Answer: polar

H₂S has a structure similar to water.

1. What is its dot structure? _______

2. Would you expect it to be polar or nonpolar? _______

(Note: Sulfur and oxygen are both in Group VIA in the periodic table, so they have the same number of outer shell electrons.)

Answer:

1. (Or something similar. The molecule is bent and planar but not linear.)

2. polar

Recall that in frame 55 we could not judge whether ammonia, NH₃, would be polar or nonpolar. Ammonia is a bent molecule and is not planar. It looks like a pyramid with a triangular base. The nitrogen atom serves as the “peak” with the hydrogen atoms located at the corners of the base.

1. Would you expect the “tug” on the electrons to be equal and offset? __________

2. Would NH₃ be a polar or nonpolar molecule? ________

Answer: (a) no; (b) polar

Planar molecules are nonpolar. A planar molecule consists of a molecule with all of its atoms lying in the same plane. SO₃ is both planar and nonlinear.

1. Which of these structures could represent SO₃?

o

o (1) (2) (3)

2. Is SO₃ polar or nonpolar? ____________

Answer:

1. Structure 2 is the only one that is planar and nonlinear.

2. nonpolar

It is also possible to determine experimentally if compounds are polar covalent, nonpolar covalent, or ionic. A liquefied ionic compound, or an ionic compound in water solution, will conduct electricity and a covalent compound will not. A polar covalent compound will be attracted toward a charged rod. A nonpolar compound will not be attracted.

A compound that will not conduct electricity and will not be attracted by a charged rod should be which kind of compound? __________

Answer: nonpolar covalent

A stream of pure water is attracted by a charged rod, but the water does not conduct electricity.

Water must be what kind of compound? ____________________

Answer: polar covalent

A water solution of KCl is a good conductor of electricity. KCl must be what kind of compound? ___________________________

Answer: ionic

PERIODIC PROPERTIES

Use the periodic table for the following review questions.

1. Where are metals located in the periodic table: right side, left side, or both sides of the steplike line? __________________

2. Where are nonmetals located? __________________

Answer: (a) left side; (b) right side

The alkali metals (in Group IA) are the most metallic family in the periodic table. Metallic characteristics increase with increased atomic number. Given these statements, which element could be expected to be the most metallic? __________________

Answer: Fr (francium) would be the most metallic, since it is the element with the largest atomic number in the most metallic group (IA).

Of the alkaline earth family (IIA), which element could be expected to be the most metallic? __________________

Answer: Ra (radium) is the most metallic element in Group IIA, since it has the largest atomic number in the group.

In the same group, nonmetallic properties decrease as the atomic number increases. Of the halogens (Group VIIA), which is the most nonmetallic? _________________

Answer: F (fluorine) (Since nonmetallic properties decrease with increasing atomic number, the most nonmetallic halogen must be the one with the smallest atomic number.)

Which of the noble gases (Group VIIIA) could be expected to have the most non-metallic characteristics? __________

Answer: He (helium) (the noble gas with the smallest atomic number; therefore, it is the most nonmetallic of the noble gases)

A number of other properties also increase or decrease in relation to position in the periodic table. One such property is electronegativity, which we have just learned is a measure of the tendency of a covalently bonded atom to attract the bonding electron pair. Electronegativity generally increases from left to right across the periodic table and from the bottom to the top. The elements of least electronegativity would be found on the bottom left side of the periodic table. Where would the elements of greatest electronegativity be located? ___________________

Answer: top right side of the periodic table (The most electronegative element is fluorine. Note that the electronegativities of the noble gases are unknown.)

Ionization energy (the energy required to remove electrons from a neutral atom) increases and decreases in the same general way as electronegativity. Elements requiring the most ionization energy are found on the top right side of the periodic table. Helium, the element of lowest atomic number in the noble gases, requires the most ionization energy of any element.

Which element within the halogen family (Group VIIA) could be expected to require the most ionization energy? _______________

Answer: fluorine (Within a group or family, the element of lowest atomic number would require the most ionization energy. The halogen of lowest atomic number is fluorine.)

In Group IIA (the alkaline earth metals), which element could be expected to require the greatest ionization energy? _________ Which element in Group IIA has the greatest electronegativity? ___________

Answer: Be (beryllium) is the correct answer to both questions. It has the lowest atomic number within the group. Thus, it has the greatest electronegativity and requires the greatest ionization energy within Group IIA.

Within the third period (the elements from Na on the left to Ar on the right), which element would be expected to require the greatest ionization energy? _______________

Answer: argon (Ionization energy generally increases from left to right in the periodic table.)

Within the fourth period (potassium to krypton), which element would be expected to require the least ionization energy? ________________

Answer: potassium

You have just learned that the valence shell electron structure of an atom is the controlling factor of several properties of the atom. It determines whether an atom gains or loses electrons and how many. It also determines the type of chemical bond formed when atoms combine and whether atoms have metallic or nonmetallic properties.

In any vertical column of elements in the A groups of the periodic table, as you go from top to bottom:

· the elements become more metallic;

· the number of electron shells increases;

· the number of valence shell electrons remains the same;

· the ionization energy decreases;

· the electronegativity decreases.

In any horizontal row of the elements in the A groups of the periodic table, as you go from left to right:

· the elements become more nonmetallic;

· the number of electron shells remains the same;

· the number of valence shell electrons increases;

· the ionization energy increases;

· the electronegativity increases.

You have been introduced to ions and how they combine to form compounds. You will learn more about molecules, molecular weights, naming of compounds, and writing formulas of compounds in the chapters that follow.

A key point to remember is that even though we focused primarily on the first 20 elements in the periodic table, you can apply what you have learned to any of the elements in the A groups.

SELF-TEST

This self-test is designed to show how well you have mastered this chapter's objectives. Correct answers and review instructions follow the test. Use any of the tables presented in the chapter to answer the questions.

1. A covalent bond is formed when two atoms (share or exchange) ___________ two or more electrons.

2. When two atoms form a bond by the exchange of one or more valence electrons, a(n) (covalent or ionic) ___________ bond is formed.

3. How many electrons are present in the outermost shell of a neutral S atom? ___________

4. How many electrons are present in the outermost shell of a neutral Mg atom? ___________

5. Which of the following neutral atoms has three electrons in its outmost shell, silicon (Si), chlorine (Cl), or boron (B)? ___________

6. Circle the correct responses in the following questions.

1. Which has the lower ionization energy, Li or K?

2. Which would be more polar, HF or HBr?

3. Which is more nonmetallic, F or I?

4. Which is more electronegative, K or Rb?

5. Which has more outer shell electrons, Ca or C?

6. Which would you expect to be more ionic, LiF or HCI?

7. Circle the correct responses in the following questions.

1. Which has the higher ionization energy, Mg or Sr?

2. Which would be less polar, HCl or HI?

3. Which is more nonmetallic, Cl or Br?

4. Which is more electronegative, Cr or Y?

5. Which has fewer outer shell electrons, Li or F?

6. Which would you expect to be more ionic, CaO or H₂O?

8. Write equations, using electron dot symbols, showing the formation of MgCl₂ and H₂O using Mg, Cl, H, and O as starting materials. What type of bond is formed in each case?

9. Write equations, using electron dot symbols, showing the formation of LiF and H₂S using Li, F, H, and S as starting materials. What type of bond is formed in each case?

10. Here are the dot symbols of four fictitious elements:

1. In what groups would elements Q and Y be found in the periodic table?_______________________

2. Would a compound formed between Q and Y more likely be ionic or covalent? ___________________

3. What ion would you expect element X to form? _________

4. X and Y combine to form a molecule of XY₂, which is covalent. Which of the following could be the correct dot structure for the compound? __________

11. Which noble gas structure does Mg have when it becomes an ion? __________

12. Which noble gas structure does Cl have when it becomes an ion? __________

13. What is the line structure for HBr? __________

14. Nitrogen dioxide (NO₂) is a molecule with a bent structure. The NO₂ molecule is (polar, nonpolar) _______.

15. Phosphorus trichloride (PCl₃) is used in making compounds known as haloalkanes from alcohols. Phosphorus trichloride's phosphorus atom has a lone pair of electrons on its line structure. As a result PCl₃ is (polar, nonpolar) _______.

ANSWERS

Compare your answers to the self-test with those given below. If you answer all questions correctly, you are ready to proceed to the next chapter. If you miss any, review the frames indicated in parentheses following the answers. If you miss several questions, you should probably reread the chapter carefully.

1. share (frame 34)

2. Ionic (frames 27—33)

3. 6 (S is in Group VIA) (frames 2, 5, 8)

4. 2 (Mg is in Group IIA) (frames 2, 5, 8)

5. Boron is in Group IIIA so it will have three electrons in its outmost shell (frames 2, 5, 8)

6.

1. K (frames 78—81 and summary)

2. HF (frames 54—61)

3. F (frames 73—76)

4. K (frames 53, 61)

5. C (frames 2—6)

6. LiF (frames 54—61)

7.

1. Mg (frames 78—81 and summary)

2. HI (frames 54—61)

3. Cl (frames 73—76)

4. Cr (frames 53, 61)

5. Li (frames 2—6)

6. CaO (frames 54—61)

8. , ionic

, polar covalent (frames 39—48)

9.  , ionic

, polar covalent (frames 39—48)

10.

1. Q in Group IA, Y in Group VIIA (frames 2, 8)

2. ionic (frames 27—61)

3. 2— (frames 12—26)

4. structure 2 (It is the only one that obeys the octet rule for each atom.) (frames 17, 43—48)

11. neon (frames 18—20)

12. argon (frames 18—20)

13. linear (frame 63)

14. polar (frames 63, 67)

15. polar (frames 63, 67)

G.N. LEWIS AND CHEMISTRY

In this chapter you learned about Lewis symbols and how an understanding of these are used to make covalent bonds and molecules. This is very important concept in chemistry, especially in the field of organic chemistry. Lewis symbols helped pioneer our understanding for chemical bonding.

Gilbert Newton Lewis was born in Weymouth, Massachusetts, in 1875 and received his Ph.D. in chemistry from Harvard University. He was a prominent physical chemist who is most widely known for his work in covalent bonding.

G.N. Lewis

Lewis symbols and Lewis structures (also known as Lewis dot diagrams) all originate with G.N. Lewis's work in covalent bonding.

Figure of Lewis symbol and Lewis structures: (A) Lewis symbol for carbon, (B) Lewis structure for an organic compound, C₂H₅Cl, and (C) Lewis structure for C₂H₅Cl using bars between two atoms instead of two dots.

In this figure you first see the Lewis symbol for carbon. As you learned in this chapter, the Lewis symbol for carbon has four dots around it representing its four valence electrons. It has four valence electrons and is found in Group IV of the periodic table. Figure B shows a Lewis structure (aka Lewis dot structure) for an organic compound, C₂H₅Cl (chloroethane). The Lewis structure shows the arrangement of electrons within a molecule.

Two shared electrons between two atoms represent a bond. In Figure B you also notice that chlorine has three sets of electron pairs around it. Chlorine's three sets of electron pairs are not involved in bonding. Because they are not involved in bonding, electron pairs that reside on atoms are known as lone-pair electrons or nonbonding electron pairs. Finally, Figure C shows a shorthand version of Figure B, which uses a bar in place of two shared dots between two atoms within a bond. This bar, or bond, is often drawn between two atoms instead of two dots. Both Figure B and Figure C represent the same molecule with Figure C's structure more commonly drawn especially for organic compounds as you will see in Chapter 14.

G.N. Lewis not only helped our understanding of covalent bonding but was also instrumental in the fields of thermodynamics, photochemistry, and acid-base chemistry. It is no surprise that Lewis was nominated for the Nobel Prize in Chemistry. In fact, he was nominated 41 times, but unfortunately he never won it. Nonetheless, G.N. Lewis's work remains foundational to much of our understanding in chemistry.