Bonding valence electrons with dots

A structural formula that shows all valence electrons, with the bonds symbolized by dashes ( —) or by pairs of dots, and nonbonding electrons symbolized by dots. (p. 7)... 

The American chemist G. N. Lewis introduced a useful model that describes the electronic structure of the atom and provides a starting point for describing chemical bonds. The Lewis model represents the valence electrons as dots arranged around the chemical symbol for the atom the core electrons are not shown. The first four dots are displayed singly around the four sides of the elemental symbol. If the atom has more than four valence electrons, their dots are then paired with those already present. The result is a Lewis dot symbol for that atom. The Lewis notation for the elements of the first two periods is... 

The electron-dot symbol or electron-dot structure of an element shows the valence electrons as dots. Electrons that are paired in an orbital are shown as a pair of dots, and unpaired electrons are shown as single dots. The paired valence electrons are called lone pairs (because they do not participate in bonding). In an electron-dot symbol, the lone pairs and the single dots are arranged to the right, left, top, and bottom of the element s symbol. The electron-dot symbol for fluorine can be drawn with the single dot in any of the four positions ... 

Hydrocarbons, as the name suggests, are organic compounds that consist of only carbon and hydrogen. In the simplest hydrocarbon, methane (CH4), the carbon atom forms an octet by sharing four valence electrons with four hydrogen atoms. In the electron-dot formula, each shared pair of electrons represents a single covalent bond. In organic... 

Covalent bonds form between atoms with similar electronegativities. In these reactions, electrons do not migrate from one atom to another as they do in ionic bonds they are shared by the atoms in the molecule. A good way to visualize this was proposed by Gilbert Lewis, a chemist at the University of California, Berkeley. His representations of molecular bonds are called Lewis dot structures. These structures use dots to denote the valence electrons of an element or molecule. 

These are extensions of Lewis dot structures, where bonding electrons associated with each bond are shown as dots. In our simple structures, bonding is associated with eight electrons in the valence shell of the atom, unless it is hydrogen, when two electrons are required for bonding. Whilst we have almost completely abandoned putting in electron dots for bonds, we still routinely show some pairs of electrons not involved in bonding (lone pairs) because these help in our mechanistic rationalizations of chemical reactions. 

The first thing you must be able to do in order to predict molecular shapes is to draw an electron-dot formula, so we ll tackle that subject first Including H, there are 16 active nonmetals for which you should know the numbers of valence electrons in the uncombined atoms Except for H (which has only one s electron), these elements are all found to the right of the diagonal in the p block of the periodic table (see inside front cover) Each atom has two v electrons in its valence shell, the number ofp electrons is different for different atoms (Basically, we are uninterested in metals here, metals rarely form predominantly covalent bonds, but tend to form ionic bonds ignore the noble gases, with an already filled s-yi6 unreactive )... 

The splitting of a Cl2 molecule is an initiation step that produces two highly reactive chlorine atoms. A chlorine atom is an example of a reactive intermediate, a short-lived species that is never present in high concentration because it reacts as quickly as it is formed. Each Cl- atom has an odd number of valence electrons (seven), one of which is unpaired. The unpaired electron is called the odd electron or the radical electron. Species with unpaired electrons are called radicals or free radicals. Radicals are electron-deficient because they lack an octet. The odd electron readily combines with an electron in another atom to complete an octet and form a bond. Figure 4-1 shows the Lewis structures of some free radicals. Radicals are often represented by a structure with a single dot representing the unpaired odd electron. 

Resonance Structures Molecules with two or more valid Lewis dot structures are said to be resonant. The actual structure is neither of the alternatives but rather a lower-energy molecule with delocalized valence electrons. Benzene with its alternating double and single bonds is an example of a resonant structure. Benzene actually has no single or double bonds. Its real structure lies somewhere between the two possibilities. 

The existence of multiple bonds between atoms has been known for some considerable time. Thus with the nitrogen molecule each atom uses fully its three valency bonds and the link is therefore a triple one. Such a bond was represented by Lewis as three pairs of dots to correspond to the three pairs of bonding p electrons,... 

Both ionic and covalent bonds involve valence electrons, the electrons in the outermost energy level of an atom. In 1920, G. N. Lewis, the American chemist shown in Figure 9, came up with a system to represent the valence electrons of an atom. This system—known as electron-dot diagrams or Lewis structures —uses dots to represent valence electrons. Lewis s system is a valuable model for covalent bonding. However, these diagrams do not show the actual locations of the valence electrons. They are models that help you to keep track of valence electrons. 

The Lewis model for covalent bonding starts with the recognition that electrons are not transferred from one atom to another in a nonionic compound, but rather are shared between atoms to form covalent bonds. Hydrogen and chlorine combine, for example, to form the covalent compound hydrogen chloride. This result can be indicated with a Lewis diagram for the molecule of the product, in which the valence electrons from each atom are redistributed so that one electron from the hydrogen atom and one from the chlorine atom are now shared by the two atoms. The two dots that represent this electron pair are placed between the symbols for the two elements ... 

Molecular structure in which the valence electrons are shown as dots placed between the bonded atoms, with one pair of dots representing two electrons or one (single) covalent bond, for example... 

Electron flow paths are written in the language of Lewis dot structures and curved arrows. Lewis dot structures are used to keep track of all electrons, and curved arrows are used to symbolize electron movement. You must be able to draw a proper Lewis structure complete with formal charges accurately and quickly. Your command of curved arrows must also be automatic. These two points cannot be overemphasized, since all explanations of reactions will be expressed in the language of Lewis structures and curved arrows. A Lewis structure contains the proper number of electrons, the correct distribution of those electrons over the atoms, and the correct formal charge. We will show all valence electrons lone pairs are shown as darkened dots and bonds by lines. 

One way to deal with the specific geometry of a molecule is to return to Coulomb s law, that is, to look at electron-electron repulsion. VSEPR is such an electrostatic theory of bonding. As with Lewis dot structures, it ignores specific orbitals. The observed geometry reflects the attempt to minimize electron-electron repulsion by maximizing the distance between electrons. Bond angles are determined solely by the number of valence electrons around a central atom. It is instructive to use examples ... 

Two sets of atomic symbols are shown on the facing page the gray symbols, with dots representing electrons, and the red symbols, with dashes representing valence bonds. Both sets are commonly used by chemists, the choice being made as indicated by convenience or habit. Often the two sets are combined, bonds being represented by dashes and unshared electrons by dots. 

To write electron dot diagrams for molecules that contain several atoms, first determine the number of valence electrons available in each atom. Second, determine the number of electrons necessary to satisfy the octet mle with no sharing. The difference between the numbers in these first two steps is the number of bonding electrons. Place the atoms as symmetrically as possible. Place the number of electrons to be shared between the atoms, one pair at a time, at first one pair between each pair of atoms. Use as many pairs as remain to make double or triple bonds. Add the remainder of available electrons to complete the octets of all the atoms. There should be just enough if the molecule or ion follows the octet rule. (There are exceptions, which will not be covered here.)... 

Can you arrange the 16 valence electrons from these three atoms to produce a molecule in which aU three atoms have a stable configuration You know that at least one bond must exist between the carbon and each oxygen, so start there. Here s an approach to the puzzle. Have each oxygen share an electron with carbon as in the following dot structures. 

The rules and procedures for drawing structural isomers are the same used for drawing electron dot formulas. Every atom in the molecular formula must be used and each atom must have its valence satisfied. To draw a structure, bond all atoms with a valence greater than one with single bonds. Attach monovalent atoms to the polyvalent ones until all valences have been satisfied. If there are insufficient monovalent atoms in the formula to accomplish this, insert double bonds, triple bonds or draw cyclic structures until it is possible to satisfy all valences. To draw isomers, vary the arrangements of atoms and bonds to form different molecules. 

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