(Generally, the least electronegative element should be placed in the center.) Connect each atom to the central atom with a single bond (one electron pair).

Draw a skeleton structure of the molecule or ion, arranging the atoms around a central atom.

For anions, add one electron for each negative charge. For cations, subtract one electron for each positive charge. Determine the total number of valence (outer shell) electrons.See these examples:įor more complicated molecules and molecular ions, it is helpful to follow the step-by-step procedure outlined here: A double bond forms when two pairs of electrons are shared between a pair of atoms, as between the carbon and oxygen atoms in CH 2O (formaldehyde) and between the two carbon atoms in C 2H 4 (ethylene):Ī triple bond forms when three electron pairs are shared by a pair of atoms, as in carbon monoxide (CO) and the cyanide ion (CN –): Writing Lewis Structures with the Octet Ruleįor very simple molecules and molecular ions, we can write the Lewis structures by merely pairing up the unpaired electrons on the constituent atoms. However, a pair of atoms may need to share more than one pair of electrons in order to achieve the requisite octet. Oxygen and other atoms in group 16 obtain an octet by forming two covalent bonds: Double and Triple BondsĪs previously mentioned, when a pair of atoms shares one pair of electrons, we call this a single bond. To obtain an octet, these atoms form three covalent bonds, as in NH 3 (ammonia).

Group 15 elements such as nitrogen have five valence electrons in the atomic Lewis symbol: one lone pair and three unpaired electrons. The transition elements and inner transition elements also do not follow the octet rule: Because hydrogen only needs two electrons to fill its valence shell, it is an exception to the octet rule. These four electrons can be gained by forming four covalent bonds, as illustrated here for carbon in CCl 4 (carbon tetrachloride) and silicon in SiH 4 (silane). For example, each atom of a group 14 element has four electrons in its outermost shell and therefore requires four more electrons to reach an octet. The number of bonds that an atom can form can often be predicted from the number of electrons needed to reach an octet (eight valence electrons) this is especially true of the nonmetals of the second period of the periodic table (C, N, O, and F). The tendency of main group atoms to form enough bonds to obtain eight valence electrons is known as the octet rule. This allows each halogen atom to have a noble gas electron configuration. The other halogen molecules (F 2, Br 2, I 2, and At 2) form bonds like those in the chlorine molecule: one single bond between atoms and three lone pairs of electrons per atom. Each Cl atom interacts with eight valence electrons: the six in the lone pairs and the two in the single bond. A dash (or line) is sometimes used to indicate a shared pair of electrons:Ī single shared pair of electrons is called a single bond. The Lewis structure indicates that each Cl atom has three pairs of electrons that are not used in bonding (called lone pairs) and one shared pair of electrons (written between the atoms). For example, when two chlorine atoms form a chlorine molecule, they share one pair of electrons: We also use Lewis symbols to indicate the formation of covalent bonds, which are shown in Lewis structures, drawings that describe the bonding in molecules and polyatomic ions. The total number of electrons does not change. Figure 7.10 Cations are formed when atoms lose electrons, represented by fewer Lewis dots, whereas anions are formed by atoms gaining electrons.

Periodic table and valence electrons