Valence electrons as dots

Why are some substances chemically bonded molecules and others are an association of

ions

? The answer to this question depends upon the electronic structures of the

atoms

and nature of the chemical forces within the compounds. Although there are no sharply defined boundaries,

chemical bonds

are typically classified into three main types: ionic bonds, covalent bonds, and metallic bonds. In this chapter, each type of bond wil be discussed and the general properties found in typical substances in which the bond type occurs

1. Ionic bonds results from electrostatic forces that exist between

   ions

   of opposite charge. These bonds typically involves a metal with a nonmetal

2. Covalent bonds result from the sharing of electrons between two

   atoms

   . The bonds typically involves one nonmetallic element with another

3. Metallic bonds These bonds are found in solid metals (copper, iron, aluminum) with each metal bonded to several neighboring groups and bonding electrons free to move throughout the 3-dimensional structure.

Each bond classification is discussed in detail in subsequent sections of the chapter. Let's look at the preferred arrangements of electrons in

atoms

when they form chemical compounds.

Figure 9.3.19.3.1: G. N. Lewis and the

Octet Rule

. (a) Lewis is working in the laboratory. (b) In Lewis’s original sketch for the

octet rule

, he initially placed the electrons at the corners of a cube rather than placing them as we do now.

Lewis Symbols

At the beginning of the 20^th century, the American chemist G. N. Lewis (1875–1946) devised a system of symbols—now called

Lewis electron dot symbols

(often shortened to Lewis dot symbols) that can be used for predicting the number of bonds formed by most elements in their compounds. Each Lewis dot symbol consists of the chemical symbol for an element surrounded by dots that represent its

valence electrons

.

Lewis Dot symbols:

• convenient representation of

  valence electrons

• allows you to keep track of

  valence electrons

  during bond formation

• consists of the chemical symbol for the element plus a dot for each valence electron

To write an element’s Lewis dot symbol, we place dots representing its

valence electrons

, one at a time, around the element’s chemical symbol. Up to four dots are placed above, below, to the left, and to the right of the symbol (in any order, as long as elements with four or fewer

valence electrons

have no more than one dot in each position). The next dots, for elements with more than four

valence electrons

, are again distributed one at a time, each paired with one of the first four. For example, the

electron configuration

for atomic sulfur is [Ne]3s²3p⁴, thus there are six

valence electrons

. Its Lewis symbol would therefore be:

Fluorine, for example, with the

electron configuration

[He]2s²2p⁵, has seven

valence electrons

, so its Lewis dot symbol is constructed as follows:

Figure 8.1.2.

Lewis used the unpaired dots to predict the number of bonds that an element will form in a compound. Consider the symbol for nitrogen in Figure 8.1.2. The Lewis dot symbol explains why nitrogen, with three unpaired

valence electrons

, tends to form compounds in which it shares the unpaired electrons to form three bonds. Boron, which also has three unpaired

valence electrons

in its Lewis dot symbol, also tends to form compounds with three bonds, whereas carbon, with four unpaired

valence electrons

in its Lewis dot symbol, tends to share all of its unpaired

valence electrons

by forming compounds in which it has four bonds.

Figure 9.3.29.3.2: Lewis Dot Symbols for the Elements in Period 2L i has one dot on the right, B e has one dot on the left and right, B has one dot on the left, right, and above. C has one dot above, below, on the left and on the right. N has one dot above, below, on the left and two dots on the right. O has two dots on the left and right and one dot above and below. F has two dots above, on the left, on the right and on dot below. N e has two dots above, below, on the right, and on the left.

The

Octet Rule

In 1904, Richard Abegg formulated what is now known as Abegg's rule, which states that the difference between the maximum positive and negative valences of an element is frequently eight. This rule was used later in 1916 when Gilbert N. Lewis formulated the "

octet rule

" in his cubical atom

theory

. The

octet rule

refers to the tendency of

atoms

to prefer to have eight electrons in the valence shell. When

atoms

have fewer than eight electrons, they tend to react and form more stable compounds.

Atoms

will react to get in the most stable state possible. A complete octet is very stable because all orbitals will be full.

Atoms

with greater stability have less energy, so a reaction that increases the stability of the

atoms

will release energy in the form of heat or light ;reactions that decrease stability must absorb energy, getting colder.

When discussing the

octet rule

, we do not consider d or f electrons. Only the s and p electrons are involved in the

octet rule

, making it a useful rule for the main group elements (elements not in the transition metal or inner-transition metal blocks); an octet in these

atoms

corresponds to an electron configurations ending with s²p⁶.

Definition:

Octet Rule

A stable arrangement is attended when the atom is surrounded by eight electrons. This octet can be made up by own electrons and some electrons which are shared. Thus, an atom continues to form bonds until an octet of electrons is made. This is known as

octet rule

by Lewis.

1. Normally two electrons pairs up and forms a bond, e.g., H2H2

2. For most

   atoms

   there will be a maximum of eight electrons in the valence shell (octet structure), e.g., CH4CH4

Figure 1: Bonding in H2𝐻2 and methane (CH4𝐶𝐻4)A

hydrogen bonds

with each of the four

valence electrons

of a carbon meaning that four hydrogens can bond with one carbon.

The other tendency of

atoms

is to maintain a neutral charge. Only the

noble gases

(the elements on the right-most column of the

periodic table

) have zero charge with filled valence octets. All of the other elements have a charge when they have eight electrons all to themselves. The result of these two guiding principles is the explanation for much of the reactivity and bonding that is observed within

atoms

:

atoms

seek to share electrons in a way that minimizes charge while fulfilling an octet in the valence shell.

The

noble gases

rarely form compounds. They have the most stable configuration (full octet, no charge), so they have no reason to react and change their configuration. All other elements attempt to gain, lose, or share electrons to achieve a noble gas configuration.

Example 9.3.19.3.1:

Salt

The formula for table

salt

is NaCl. It is the result of Na⁺

ions

and Cl^-

ions

bonding together. If sodium metal and chlorine gas mix under the right conditions, they will form

salt

. The sodium loses an electron, and the chlorine gains that electron. In the process, a great amount of light and heat is released. The resulting

salt

is mostly unreactive — it is stable. It will not undergo any explosive reactions, unlike the sodium and chlorine that it is made of. Why?

Solution

Referring to the

octet rule

,

atoms

attempt to get a noble gas

electron configuration

, which is eight

valence electrons

. Sodium has one valence electron, so giving it up would result in the same

electron configuration

as neon. Chlorine has seven

valence electrons

, so if it takes one it will have eight (an octet). Chlorine has the

electron configuration

of argon when it gains an electron.

The

octet rule

could have been satisfied if chlorine gave up all seven of its

valence electrons

and sodium took them. In that case, both would have the electron configurations of noble gasses, with a full valence shell. However, their charges would be much higher. It would be Na^7- and Cl⁷⁺, which is much less stable than Na⁺ and Cl^-.

Atoms

are more stable when they have no charge, or a small charge.

Lewis dot symbols can also be used to represent the

ions

in ionic compounds. The reaction of cesium with fluorine, for example, to produce the ionic compound CsF can be written as follows:

No dots are shown on Cs⁺ in the product because cesium has lost its single valence electron to fluorine. The transfer of this electron produces the Cs⁺ ion, which has the valence

electron configuration

of Xe, and the F⁻ ion, which has a total of eight

valence electrons

(an octet) and the Ne

electron configuration

. This description is consistent with the statement that among the main group elements,

ions

in simple binary ionic compounds generally have the electron configurations of the nearest noble gas. The charge of each ion is written in the product, and the anion and its electrons are enclosed in brackets. This notation emphasizes that the

ions

are associated electrostatically; no electrons are shared between the two elements.

As you might expect for such a qualitative approach to bonding, there are exceptions to the

octet rule

, which we describe elsewhere. These include molecules in which one or more

atoms

contain fewer or more than eight electrons.

Summary

Lewis dot symbols can be used to predict the number of bonds formed by most elements in their compounds. One convenient way to predict the number and basic arrangement of bonds in compounds is by using

Lewis electron dot symbols

, which consist of the chemical symbol for an element surrounded by dots that represent its

valence electrons

, grouped into pairs often placed above, below, and to the left and right of the symbol. The structures reflect the fact that the elements in period 2 and beyond tend to gain, lose, or share electrons to reach a total of eight

valence electrons

in their compounds, the so-called

octet rule

. Hydrogen, with only two

valence electrons

, does not obey the

octet rule

.