Which molecule obeys the octet rule

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PCl5 Which of these molecules show resonance? Draw a Lewis structure for CH4 and answer the following Draw a Lewis structure for SF6 and answer the following questions based on your drawing. For the central carbon atom: The number of lone pairs The number of single bonds The number of double Lewis symbols can be used to illustrate the formation of cations from atoms, as shown here for sodium and calcium:.

Likewise, they can be used to show the formation of anions from atoms, as shown here for chlorine and sulfur:. The total number of electrons does not change.

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. For example, when two chlorine atoms form a chlorine molecule, they share one pair of electrons:. 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.

A dash or line is usually used to indicate a shared pair of electrons:. In the Lewis model, a single shared pair of electrons constitutes a single bond. Each Cl atom interacts with eight valence electrons total: the six in the lone pairs and the two in the single bond. 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.

This allows each halogen atom to have a noble gas electron configuration, which corresponds to eight valence electrons. The tendency of main group atoms to form enough bonds to obtain eight valence electrons is known as the octet rule. 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.

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. 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.

Because hydrogen only needs two electrons to fill its valence shell, it is an exception to the octet rule and only needs to form one bond. The transition elements and inner transition elements also do not follow the octet rule since they have d and f electrons involved in their valence shells. Group 15 elements such as nitrogen have five valence electrons in the atomic Lewis symbol: one lone pair and three unpaired electrons.

To obtain an octet, these atoms form three covalent bonds, as in NH 3 ammonia. Oxygen and other atoms in group 16 obtain an octet by forming two covalent bonds:. As previously mentioned, when a pair of atoms shares one pair of electrons, we call this a single bond.

However, a pair of atoms may need to share more than one pair of electrons in order to achieve the requisite octet. 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 2 O formaldehyde and between the two carbon atoms in C 2 H 4 ethylene :. A triple bond forms when three electron pairs are shared by a pair of atoms, as in carbon monoxide CO and the cyanide ion CN — :.

For very simple molecules and molecular ions, we can write the Lewis structures by merely pairing up the unpaired electrons on the constituent atoms. See these examples:. For more complicated molecules and molecular ions, it is helpful to follow the step-by-step procedure outlined here:.

Draw a skeleton structure of the molecule or ion, arranging the atoms around a central atom and connecting each atom to the central atom with a single one electron pair bond. Note that we denote ions with brackets around the structure, indicating the charge outside the brackets:.

Rearrange the electrons of the outer atoms to make multiple bonds with the central atom in order to obtain octets wherever possible. What are the Lewis structures of these molecules? Draw a skeleton and connect the atoms with single bonds.

Remember that H is never a central atom:. Where needed, place remaining electrons on the central atom:. Where needed, rearrange electrons to form multiple bonds in order to obtain an octet on each atom:. Both carbon monoxide, CO, and carbon dioxide, CO 2 , are products of the combustion of fossil fuels. Both of these gases also cause problems: CO is toxic and CO 2 has been implicated in global climate change.

What are the Lewis structures of these two molecules? Carbon soot has been known to man since prehistoric times, but it was not until fairly recently that the molecular structure of the main component of soot was discovered. In , the Nobel Prize in Chemistry was awarded to Richard Smalley , Robert Curl, and Harold Kroto for their work in discovering a new form of carbon, the C 60 buckminsterfullerene molecule. An entire class of compounds, including spheres and tubes of various shapes, were discovered based on C This type of molecule, called a fullerene, consists of a complex network of single- and double-bonded carbon atoms arranged in such a way that each carbon atom obtains a full octet of electrons.

However, many atoms below atomic number 20 often form compounds that do not follow the octet rule. For example, with the duet rule of the first principal energy level, the noble gas helium, He, has two electrons in its outer level. Since there is no 1p subshell, 1s is followed immediately by 2s, and thus level 1 can only have at most two valence electrons.

Hydrogen only needs one additional electron to attain this stable configuration, through either covalent sharing of electrons or by becoming the hydride ion :H — , while lithium needs to lose one by combining ionically with other elements.

This leads to hydrogen and lithium both having two electrons in their valence shell—the same electronic configuration as helium—when they form molecules by bonding to other elements. There are also a variety of molecules in which there are too few electrons to provide an octet for every atom. Boron and aluminum, from Group III or 13 , display different bonding behavior than previously discussed.

These atoms each have three valence electrons, so we would predict that these atoms want to bond covalently in order to gain 5 electrons through sharing to fulfill the octet rule.

However, compounds in which boron or aluminum atoms form five bonds are never observed, so we must conclude that simple predictions based on the octet rule are not reliable for Group III. Consider boron trifluoride BF 3.

The bonding is relatively simple to model with a Lewis structure if we allow each valence level electron in the boron atom to be shared in a covalent bond with each fluorine atom. In this compound, the boron atom only has six valence shell electrons, but the octet rule is satisfied by the fluorine atoms.

Lewis structure of boron trifluoride : Each pair of dots represents a pair of electrons. When placed between two atoms, the electrons are in a bond. A bond can be drawn as a line between two atoms, which also indicates two electrons. We might conclude from this one example that boron atoms obey a sextet rule.

However, boron will form a stable ion with hydrogen, BH 4 — , in which the boron atom does have a complete octet. In addition, BF 3 will react with ammonia NH 3 , to form a stable compound, NH 3 BF 3 , for which a Lewis structure can be drawn that shows boron with a complete octet. Boron trifluoride-ammonia complex : This covalent compound NH 3 BF 3 shows that boron can have an octet of electrons in its valence level. Compounds of aluminum follow similar trends.

Aluminum trichloride AlCl 3 , aluminum hydride AlH 3 , and aluminum hydroxide Al OH 3 indicate a valence of three for aluminum, with six valence electrons in the bonded molecule. However, the stability of aluminum hydride ions AlH 4 — indicates that Al can also support an octet of valence shell electrons. Although the octet rule can still be of some utility in understanding the chemistry of boron and aluminum, the compounds of these elements are harder to predict than for other elements. Some elements, most notably nitrogen, can form compounds that do not obey the octet rule.

One class of such compounds are those that have an odd number of electrons. As the octet rule requires eight electrons around each atom, a molecule with an odd number of electrons must disobey the octet rule.

Recall that the Lewis structure of a molecule must depict the total number of valence electrons from all the atoms which are bonded together. Nitric oxide has the formula NO. Therefore, no matter how electrons are shared between the nitrogen and oxygen atoms, there is no way for nitrogen to have an octet.