7.3 Lewis Symbols and Structures

• The following types of bonds are found in Silicones: Si-O, Si-C, C-H, and C-C.
• The symbols d+ and d- are used to designate the positive and negative atoms.

• In this chapter, we have discussed the various types of bonds between atoms and/or ion.
• These bonds involve the sharing or transfer of shell electrons between atoms.
• The typical method for depicting Lewis symbols and Lewis structures is explored in this section.

• Lewis symbols are used to describe electron configurations.

• Lewis symbols for elements of the third period of the periodic table are shown in Figure 7.9.

• Lewis symbols show the number of electrons for each element in the periodic table.

• Lewis symbols are used to show the transfer of electrons during the formation of ionic compounds.

• When atoms lose electrons, Lewis dots represent them, whereas anions are formed by atoms gaining electrons.
• There is no change in the total number of electrons.

• Each atom has a noble gas electron configuration.

• The number of bonds that an atom can form can be predicted from the number of electrons needed to reach an octet; this is especially true of the nonmetals of the second period of the periodic table.
• Each atom of a group 14 element has four electrons in its shell and therefore needs four more to reach an octet.

• There is a chapter called Chemical Bonding andMolecular Geometry illustrated for carbon in CCl4 and SiH4 The octet rule states that hydrogen only needs two electrons to fill its valence shell.

• Nitrogen has one lone pair and three unpaired electrons in the atomic Lewis symbol.
• The atoms form three bonds in NH3 to get an octet.

• A single bond is when a pair of atoms share one pair of electrons.
• A pair of atoms may need to share more than one pair of electrons.

• The Lewis structures can be written if we pair up the unpaired electrons on the atoms.

• Determine the number of electrons in the shell.
• For cations, subtract one electron from each This OpenStax book.
• Add one electron for each negative charge.

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

• Distribute the remaining electrons as lone pairs on the terminal atoms.

• The electrons should be placed on the central atom.

• Rearrange the electrons of the outer atoms in order to make bonds with the central atom.

• Determine the number of electrons in the molecule or ion.

• The OF2 O has 6 valence electrons/atom x 1 atom and 7 valence electrons/atom x 2 atoms.
• 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 electron pair.

• We need to use experimental evidence to choose the correct arrangement of atoms.
• The less positive elements are more likely to be central atoms.

• The less negative carbon atom occupies the central position with the oxygen and hydrogen atoms surrounding it.

• The exception is that hydrogen is not a central atom.
• fluorine can't be a central atom because it's the most negative element.

• Distribute the remaining electrons as lone pairs on the terminal atoms to complete their valence shells.

• The electrons should be placed on the central atom.

We already placed all of the electrons determined in Step 1

• Rearrange the electrons of the outer atoms in order to make bonds with the central atom.

• Nothing needs to be done because Si already has an octet.

• The carbon atom lacks an octet and we have distributed the valence electrons as lone pairs on the oxygen atoms.

• This still doesn't produce an octet, so we have to move another pair to form a triple bond.

• The cloud of Hcn was detected by the NASA's Cassini-Huygens mission.
• ethane, acetylene, and ammonia are found in Titan.

• Attach the atoms with single bonds by drawing a skeleton.

• Carbon dioxide, CO2, and carbon monoxide are products of the burning of fossil fuels.
• CO is toxic and CO2 has been implicated in global climate change.

• Since prehistoric times, carbon soot has been known to man, but it wasn't until recently that the structure of the main component of soot was discovered.
• The C60 buckminsterfullerene molecule, a new form of carbon, was discovered by Richard Smalley, Robert Curl, and Harold Kroto.
• There are a variety of applications for this type of molecule.
• Because of their size and shape, fullerenes have shown potential in various applications from hydrogen storage to targeted drug delivery systems.
• They have unique electronic and optical properties that can be used in solar powered devices.

• One of the leading advocates for fullerene chemistry was Richard Smalley, a professor of physics, chemistry, and astronomy at Rice University.

• The Lewis structures of many covalent molecules do not have eight electrons.
• The odd-electron molecule has an unpaired electron and is in one of the three categories.

• A noble gas configuration requires a central atom that has fewer electrons than needed.

• A noble gas configuration requires a central atom that has more electrons than is needed.

• When oxygen and nitrogen react at high temperatures, Nitric oxide is produced in internal combustion engines.

• The sum of the electrons is 11.
• The odd number tells us that we have a free radical, so we know that not every atom has eight electrons.

• The electrons are distributed around both atoms because there is no central atom.

• This step does not apply since there are no remaining electrons.

• We know that an odd-electron molecule can't have an octet for every atom, but we want to get each atom as close to an octet as possible.
• Nitrogen has five electrons around it.
• To form a NO double bond, we take one of the lone pairs from oxygen and use it.

• There are a few molecules that do not have a filled valence shell.
• These aremolecules with central atoms from groups 2 and 12 outer atoms that are hydrogen, or other atoms that do not form multiple bonds.
• The Lewis structures of beryllium dihydride, BeH2, and Boron trifluoride each have only four and six electrons.
• It is possible to draw a structure with a double bond between a boron atom and a fluorine atom, satisfying the octet rule, but experimental evidence indicates the bond lengths are closer to that expected for B-F single bonds.
• The best Lewis structure has three B-F single bonds and an electron deficient boron.
• The reactivity of the compound is similar to an electron deficient boron.
• The B-F bonds are slightly shorter than expected, indicating that there is at least one double bond in the molecule.

• An atom that doesn't have eight electrons is very reactive.
• It is easy to combine with a molecule with an atom.

• When we write the Lewis structures, we find that we have left over electrons after filling the outer atoms with eight electrons.
• The central atom has additional electrons assigned to it.

• There are a number of stable compounds in xenon.
• We looked at XeF4 earlier.

• The Lewis structure of any molecule can be drawn by following the six steps.
• Since not all of them apply, we can condense the last few steps.