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Chapter 10
Chemical Bonding
10-00-02_1un
Title
 Electron configuration of O
Caption
 The valence electrons are always in the last shell; here the n = 2 shell is the last one, and six valence electrons reside there.
Notes
 All Group 6 elements will have six valence electrons.
Keywords
 energy, electron, quantum number, shell, subshell, valence electron, electron configuration, oxygen
10-00-02_2un
Title
 Using dots to represent valence electrons
Caption
 The Lewis structure uses dots to represent the valence electrons. This system facilitates visualization of the number of electrons an atom needs to satisfy its octet.
Notes
 The Lewis structure is somewhat deceptive, in that it suggests that electrons are particles (dots). The wave-particle duality is ignored.
Keywords
 Lewis structure, electron, valence electron
10-00-03un
Title
 Convention for drawing Lewis dot structures
Caption
 The electrons (as dots) should be drawn as shown, where X is any element.
Notes
 The Lewis structure uses dots to represent the valence electrons. This system facilitates visualization of the number of electrons an atom needs to satisfy its octet. The left-right-top-bottom locations are meant to represent orbitals.
Keywords
 Lewis structure, electron, valence electron
10-00-05un
Title
 Lewis dot structure for phosphorus
Caption
 The Lewis structure uses dots to represent the valence electrons. This system facilitates visualization of the number of electrons an atom needs to satisfy its octet. In this case, phosphorus has five valence electrons
Notes
 The Lewis structure is somewhat deceptive, in that it suggests that electrons are particles (dots). The wave-particle duality is ignored.
Keywords
 Lewis structure, electron, valence electron
10-00-06un
Title
 Lewis dot structures for K and Cl
Caption
 These two elements will form ions: The K atom will give up its valence electron to Cl. In doing so, K's electron configuration becomes like that of a noble gas (low energy), and so does Cl's. This is an example of ionic bonding.
Notes
 K+ has an electron configuration 1s22s22p63s23p6 Cl- has an electron configuration 1s22s22p63s23p6 As ions, both elements have full octets, and are therefore at low energy. Getting to low energy is one reason elements undergo chemical reactions.
Keywords
 Lewis structure, electron, valence electron, potassium, chlorine
10-00-07un
Title
 K and Cl react with one another by forming ions
Caption
 These two elements will form ions: The K atom will give up its valence electron to Cl. In doing so, K's electron configuration becomes like that of a noble gas (low energy), and so does Cl's. This is an example of ionic bonding.
Notes
 K+ has an electron configuration 1s22s22p63s23p6 Cl- has an electron configuration 1s22s22p63s23p6 As ions, both elements have full octets, and are therefore at low energy. Getting to low energy is one reason elements undergo chemical reactions.
Keywords
 Lewis structure, electron, valence electron, potassium, chlorine
10-00-08un
Title
 Lewis dot structures for Mg and O
Caption
 These two elements will form ions: The Mg atom will give up its two valence electrons to O. In doing so, Mg's electron configuration becomes like that of a noble gas (low energy), and so does O's. This is an example of ionic bonding.
Notes
 Mg2+ has an electron configuration 1s22s22p6 O2- has an electron configuration 1s22s22p6 As ions, both elements have full octets, and are therefore at low energy. Getting to low energy is one reason elements undergo chemical reactions.
Keywords
 Lewis structure, electron, valence electron, magnesium, oxygen
10-00-09un
Title
 Mg and O react with one another by forming ions
Caption
 These two elements will form ions: The Mg atom will give up its two valence electrons to O. In doing so, Mg's electron configuration becomes like that of a noble gas (low energy), and so does O's. This is an example of ionic bonding.
Notes
 Mg2+ has an electron configuration 1s22s22p6 O2- has an electron configuration 1s22s22p6 As ions, both elements have full octets, and are therefore at low energy. Getting to low energy is one reason elements undergo chemical reactions.
Keywords
 Lewis structure, electron, valence electron, magnesium, oxygen
10-00-09_01un
Title
 Na and Br react with one another by forming ions
Caption
 These two elements will form ions: The Na atom will give up its valence electron to Br. In doing so, Na's electron configuration becomes like that of a noble gas (low energy), and so does Br's. This is an example of ionic bonding.
Notes
 Na+ has an electron configuration 1s22s22p6 Br- has an electron configuration 1s22s22p63s23p64s23d104p6 As ions, both elements have full octets, and are therefore at low energy. Getting to low energy is one reason elements undergo chemical reactions.
Keywords
 Lewis structure, electron, valence electron, sodium, bromine
10-00-10un
Title
 Lewis dot structures for Na and S
Caption
 These two elements will form ions: The Na atom will give up its valence electron to S. In doing so, Na's electron configuration becomes like that of a noble gas (low energy), But, S requires two electrons, not one, so a second Na must give its valence electron to S in order for S to get to low energy. The resulting formula will be Na2S. This is an example of ionic bonding.
Notes
 Na+ has an electron configuration 1s22s22p6 S2- has an electron configuration 1s22s22p63s23p6 As ions, both elements have full octets, and are therefore at low energy. Getting to low energy is one reason elements undergo chemical reactions.
Keywords
 Lewis structure, electron, valence electron, sodium, sulfur
10-00-11un
Title
 Na and S react with one another by forming ions
Caption
 These two elements will form ions: The Na atom will give up its valence electron to S. In doing so, Na's electron configuration becomes like that of a noble gas (low energy), But, S requires two electrons, not one, so a second Na must give its valence electron to S in order for S to get to low energy. The resulting formula will be Na2S. This is an example of ionic bonding.
Notes
 Na+ has an electron configuration 1s22s22p6 S2- has an electron configuration 1s22s22p63s23p6 As ions, both elements have full octets, and are therefore at low energy. Getting to low energy is one reason elements undergo chemical reactions.
Keywords
 Lewis structure, electron, valence electron, sodium, sulfur
10-00-12un
Title
 Lewis dot structures for Ca and Cl
Caption
 These two elements will form ions: The Ca atom will give up its two valence electrons to Cl. But, in doing so, each Cl atom can only accept one valence electron. Therefore, two Cl atoms must be present to accept all of the electrons that Ca will give. Ca's electron configuration becomes like that of a noble gas (low energy), and both Cl's will assume a noble gas configuration. The resulting formula will be CaCl2. This is an example of ionic bonding.
Notes
 Ca2+ has an electron configuration 1s22s22p63s23p6 Cl- has an electron configuration 1s22s22p63s23p6 As ions, both elements have full octets, and are therefore at low energy. Getting to low energy is one reason elements undergo chemical reactions.
Keywords
 Lewis structure, electron, valence electron, calcium, chlorine
10-00-13un
Title
 Ca and Cl react with one another by forming ions
Caption
 These two elements will form ions: The Ca atom will give up its two valence electrons to Cl. But, in doing so, each Cl atom can only accept one valence electron. Therefore, two Cl atoms must be present to accept all of the electrons that Ca will give. Ca's electron configuration becomes like that of a noble gas (low energy), and both Cl's will assume a noble gas configuration. The resulting formula will be CaCl2. This is an example of ionic bonding.
Notes
 Ca2+ has an electron configuration 1s22s22p63s23p6 Cl- has an electron configuration 1s22s22p63s23p6 As ions, both elements have full octets, and are therefore at low energy. Getting to low energy is one reason elements undergo chemical reactions.
Keywords
 Lewis structure, electron, valence electron, calcium, chlorine
10-00-15un
Title
 Lewis dot structures for H and O
Caption
 H and O are both nonmetals; instead of forming ions, they will share valence electrons to get to low energy. Oxygen gets to low energy by satisfying its octet; hydrogen gets to low energy by satisfying its duet. Because O needs two electrons to satisfy its octet, two hydrogens will bond with the oxygen. The resulting formula will be H2O.
Notes
 The Lewis dot formula for H2O shows that each hydrogen shares a pair of electrons with the oxygen atom. Each pair consists of one electron provided by oxygen and one by hydrogen. Of course, oxygen will be the central atom in the H2O molecule. This is an example of covalent bonding.
Keywords
 Lewis structure, electron, valence electron, covalent bonding, hydrogen, oxygen
10-00-16un
Title
 Lewis dot structures for H2O
Caption
 H and O are both nonmetals; instead of forming ions, they will share valence electrons to get to low energy. Oxygen gets to low energy by satisfying its octet; hydrogen gets to low energy by satisfying its duet. Because O needs two electrons to satisfy its octet, two hydrogens will bond with the oxygen. The resulting formula will be H2O.
Notes
 The Lewis dot formula for H2O shows that each hydrogen shares a pair of electrons with the oxygen atom. Each pair consists of one electron provided by oxygen and one by hydrogen. Of course, oxygen will be the central atom in the H2O molecule. This is an example of covalent bonding.
Keywords
 Lewis structure, electron, valence electron, covalent bonding, hydrogen, oxygen
10-00-17un
Title
 Lewis dot structures for H2O show that octets and duets are satisfied
Caption
 H and O are both nonmetals; instead of forming ions, they will share valence electrons to get to low energy. Oxygen gets to low energy by satisfying its octet; hydrogen gets to low energy by satisfying its duet. Because O needs two electrons to satisfy its octet, two hydrogens will bond with the oxygen. The resulting formula will be H2O.
Notes
 The Lewis dot formula for H2O shows that each hydrogen shares a pair of electrons with the oxygen atom. Each pair consists of one electron provided by oxygen and one by hydrogen. Of course, oxygen will be the central atom in the H2O molecule. This is an example of covalent bonding. Hydrogen need only satisfy a duet to get to low energy, because H has no p subshell to fill in its valence shell. Oxygen, on the other hand, must fill both an s subshell and a p subshell in its valence shell.
Keywords
 Lewis structure, electron, valence electron, covalent bonding, hydrogen, oxygen, octet, duet
10-00-18un
Title
 Lone pairs and bonding pairs satisfy oxygen's octet in H2O
Caption
 H and O are both nonmetals; instead of forming ions, they will share valence electrons to get to low energy. Oxygen gets to low energy by satisfying its octet; hydrogen gets to low energy by satisfying its duet. Because O needs two electrons to satisfy its octet, two hydrogens will bond with the oxygen. The resulting formula will be H2O.
Notes
 The Lewis dot formula for H2O shows that each hydrogen shares a pair of electrons with the oxygen atom. Each pair consists of one electron provided by oxygen and one by hydrogen. Of course, oxygen will be the central atom in the H2O molecule. This is an example of covalent bonding. Oxygen shares two pairs of electrons with hydrogenÑthe bonding pairsÑand controls two pairs by itselfÑthe lone pairs.
Keywords
 Lewis structure, electron, valence electron, covalent bonding, hydrogen, oxygen, lone pair, bonding pair
10-00-19un
Title
 Lewis dot structures for H2O
Caption
 H and O are both nonmetals; instead of forming ions, they will share valence electrons to get to low energy. Oxygen gets to low energy by satisfying its octet; hydrogen gets to low energy by satisfying its duet. Because O needs two electrons to satisfy its octet, two hydrogens will bond with the oxygen. The resulting formula will be H2O.
Notes
 The Lewis dot formula for H2O shows that each hydrogen shares a pair of electrons with the oxygen atom. Each pair consists of one electron provided by oxygen and one by hydrogen. Of course, oxygen will be the central atom in the H2O molecule. This is an example of covalent bonding. Oxygen shares two pairs of electrons with hydrogenÑthe bonding pairsÑthat are often represented by dashes in the Lewis dot structure.
Keywords
 Lewis structure, electron, valence electron, covalent bonding, hydrogen, oxygen
10-00-20un
Title
 Lewis dot structure for a chlorine atom
Caption
 The Lewis structure uses dots to represent the valence electrons. This system facilitates visualization of the number of electrons an atom needs to satisfy its octet. In this case, chlorine has seven valence electrons.
Notes
 The Lewis structure is somewhat deceptive, in that it suggests that electrons are particles (dots). The wave-particle duality is ignored.
Keywords
 Lewis structure, electron, valence electron, chlorine
10-00-21un
Title
 Lewis dot structure for a chlorine molecule
Caption
 The Lewis structure uses dots to represent the valence electrons. This system facilitates visualization of the number of electrons an atom needs to satisfy its octet. In this case, each chlorine has seven valence electrons. By sharing an electron from each chlorine, both atoms can accumulate the eight electrons they need to have a full octet.
Notes
 This is an example of covalent bonding. As we have seen, a dash can represent a shared pair of electronsÑa covalent bond.
Keywords
 Lewis structure, electron, valence electron, chlorine
10-00-22un
Title
 Lewis dot structure for hydrogen
Caption
 The Lewis structure uses dots to represent the valence electrons. This system facilitates visualization of the number of electrons an atom needs to satisfy its octet, or in hydrogen, the duet. In this case, hydrogen has one valence electron.
Notes
 The Lewis structure is somewhat deceptive, in that it suggests that electrons are particles (dots). The wave-particle duality is ignored.
Keywords
 Lewis structure, electron, valence electron, hydrogen
10-00-23un
Title
 Lewis dot structure for a hydrogen molecule
Caption
 The Lewis structure uses dots to represent the valence electrons. This system facilitates visualization of the number of electrons an atom needs to satisfy its octet or, for hydrogen, its duet. In this case, each hydrogen has one valence electron. By sharing an electron from each hydrogen, both atoms can accumulate the two electrons they need to have a full duet.
Notes
 This is an example of covalent bonding. As we have seen, a dash can represent a shared pair of electronsÑa covalent bond.
Keywords
 Lewis structure, electron, valence electron, hydrogen
10-00-24un
Title
 Lewis dot structure for oxygen
Caption
 The Lewis structure uses dots to represent the valence electrons. This system facilitates visualization of the number of electrons an atom needs to satisfy its octet. In this case, oxygen has six valence electrons.
Notes
 The Lewis structure is somewhat deceptive, in that it suggests that electrons are particles (dots). The wave-particle duality is ignored.
Keywords
 Lewis structure, electron, valence electron, oxygen
10-00-25un
Title
 Attempt to form an oxygen molecule
Caption
 With six valence electrons each, two oxygen atoms can use twelve electrons to satisfy both octets. By sharing one electron from each oxygen, we can't have more than one oxygen's octet satisfied. We'll need to try a different sharing pattern; a double bond will work.
Notes
 By sharing four electrons between the two atoms, both oxygens' octets would be satisfied.
Keywords
 Lewis structure, electron, valence electron, oxygen, double bond
10-00-26un
Title
 An oxygen molecule possesses a double bond
Caption
 With six valence electrons each, two oxygen atoms can use twelve electrons to satisfy both octets. By sharing one electron from each oxygen, we can't have more than one oxygen's octet satisfied. We'll need to try a different sharing pattern; a double bond will work.
Notes
 By sharing four electrons between the two atoms, both oxygens' octets would be satisfied.
Keywords
 Lewis structure, electron, valence electron, oxygen, double bond
10-00-27un
Title
 An oxygen molecule satisfies both octets with a double bond
Caption
 With six valence electrons each, two oxygen atoms can use twelve electrons to satisfy both octets. By sharing one electron from each oxygen, we can't have more than one oxygen's octet satisfied. We'll need to try a different sharing pattern; a double bond will work.
Notes
 By sharing four electrons between the two atoms, both oxygens' octets would be satisfied. Both octets consist of the double bond and two lone pairs.
Keywords
 Lewis structure, electron, valence electron, oxygen, double bond, octet
10-00-28un
Title
 Attempt to form a nitrogen molecule
Caption
 With five valence electrons each, two nitrogen atoms can use ten electrons to satisfy both octets. By sharing one electron from each nitrogen, we can't have either nitrogen's octet satisfied. We'll need to try a different sharing pattern; a triple bond will work.
Notes
 By sharing six electrons between the two atoms, both nitrogens' octets would be satisfied.
Keywords
 Lewis structure, electron, valence electron, nitrogen, triple bond
10-00-29un
Title
 Attempt to form a nitrogen molecule
Caption
 With five valence electrons each, two nitrogen atoms can use ten electrons to satisfy both octets. By sharing one electron from each nitrogen, we can't have either nitrogen's octet satisfied. We'll need to try a different sharing pattern; a triple bond will work.
Notes
 By sharing six electrons between the two atoms, both nitrogens' octets would be satisfied.
Keywords
 Lewis structure, electron, valence electron, nitrogen, triple bond
10-00-31un
Title
 Attempt to form a carbon dioxide molecule
Caption
 With six valence electrons each, two oxygen atoms can use twelve electrons to satisfy octets. Carbon brings four valence electrons. By sharing two electrons with each oxygen, carbon is able to satisfy the octets of both oxygens, but the carbon's octet remains unsatisfied. We'll need to try a different sharing pattern; double bonds will work.
Notes
 By sharing four electrons between each oxygen and the carbon, all octets would be satisfied.
Keywords
 Lewis structure, electron, valence electron, oxygen, carbon, carbon dioxide
10-00-32un
Title
 Lewis dot structure for a carbon dioxide molecule
Caption
 With six valence electrons each, two oxygen atoms can use twelve electrons to satisfy octets. Carbon brings four valence electrons. By sharing two electrons with each oxygen, carbon is able to satisfy the octets of both oxygens, but the carbon's octet remains unsatisfied. We'll need to try a different sharing pattern; double bonds will work.
Notes
 By sharing four electrons between each oxygen and the carbon, all octets would be satisfied. Each set of four electrons would form a double bond between carbon and an oxygen atom.
Keywords
 Lewis structure, electron, valence electron, oxygen, carbon, carbon dioxide
10-00-33un
Title
 Lewis dot structure for a carbon monoxide molecule
Caption
 With six valence electrons, an oxygen atom can use six electrons to satisfy octets. Carbon brings four valence electrons. By sharing six electrons, oxygen and carbon are able to satisfy their octets. This molecule provides an example of a triple bond.
Notes
 By sharing six electrons between the oxygen and the carbon, all octets would be satisfied. The set of six electrons would form a triple bond between carbon and oxygen.
Keywords
 Lewis structure, electron, valence electron, oxygen, carbon, carbon monoxide
10-00-35un
Title
 Lewis dot structure for carbon tetrachloride
Caption
 Carbon shares one of its valence electrons with each chlorine. Each chlorine likewise shares a valence electron with carbon. The result is a central carbon that satisfies its octet through formation of four single bonds.
Notes
 This is an example of a tetrahedral molecule.
Keywords
 Lewis structure, electron, valence electron, carbon, chlorine, carbon tetrachloride, covalent bonding
10-00-36un
Title
 Lewis dot structure for formaldehyde
Caption
 Carbon shares one of its valence electrons with each hydrogen, and two electrons with oxygen. Each hydrogen likewise shares a valence electron with carbon, and the oxygen shares two with carbon. The result is a central carbon that satisfies its octet through formation of single bonds with each H, and a double bond with O.
Notes
 This is an example of a trigonal planar molecule.
Keywords
 Lewis structure, electron, valence electron, carbon, oxygen, hydrogen, formaldehyde, covalent bonding
10-00-37un
Title
 Calculating the number of valence electrons in cyanide ion's Lewis dot structure
Caption
 The total number of ten comes from the carbon atom (four), the nitrogen atom (five), and the -1 charge (1). When we know the number of electrons, we can then build the Lewis dot structure in the usual manner.
Notes
 Note that negative ions require us to ADD the number of electrons designated by the charge's magnitude; positive ions require us to SUBTRACT the number of electrons designated by the charge's magnitude.
Keywords
 Lewis structure, electron, valence electron, cyanide, covalent bonding, polyatomic ion
10-00-39un
Title
 Attempt to form a cyanide ion
Caption
 With ten valence electrons available, C and N cannot satisfy both octets with a single bond. We'll need to try a different sharing pattern; a triple bond will work.
Notes
 By sharing six electrons between the two atoms, both atoms' octets would be satisfied.
Keywords
 Lewis structure, electron, valence electron, cyanide, covalent bonding, polyatomic ion
10-00-40un
Title
 Lewis dot structure for a cyanide ion
Caption
 With ten valence electrons available, C and N cannot satisfy both octets with a single bond. We'll need to try a different sharing pattern; a triple bond will work.
Notes
 By sharing six electrons between the two atoms, both atoms' octets would be satisfied. The square brackets tell us that the atoms form a polyatomic ion.
Keywords
 Lewis structure, electron, valence electron, cyanide, covalent bonding, polyatomic ion
10-00-41un
Title
 Calculating the number of valence electrons in ammonium ion's Lewis dot structure
Caption
 The total number of eight comes from the nitrogen atom (five), each hydrogen atom (four: one from each H), and the +1 charge (delete one). When we know the number of electrons, we can then build the Lewis dot structure in the usual manner.
Notes
 Note that negative ions require us to ADD the number of electrons designated by the charge's magnitude; positive ions require us to SUBTRACT the number of electrons designated by the charge's magnitude. Also, recall that H satisfies a duet, not an octet.
Keywords
 Lewis structure, electron, valence electron, ammonium, covalent bonding, polyatomic ion
10-00-43un
Title
 Lewis dot structure for an ammonium ion
Caption
 With eight valence electrons available, H and N can satisfy N's octet and the four H's duets, with single bonds.
Notes
 By sharing eight electrons between the five atoms, all atoms are at low energy; the ion is stable. The square brackets tell us that the atoms form a polyatomic ion.
Keywords
 Lewis structure, electron, valence electron, ammonium, covalent bonding, polyatomic ion
10-00-44un
Title
 Lewis dot structure for a hypochlorite ion
Caption
 With fourteen valence electrons available, Cl and O can satisfy their octets with single bonds.
Notes
 The fourteen electrons come from Cl (7), O (6), and the -1 charge (1). The square brackets tell us that the atoms form a polyatomic ion.
Keywords
 Lewis structure, electron, valence electron, hypochlorite, covalent bonding, polyatomic ion
10-00-45un
Title
 NO: an exception to the octet rule
Caption
 NO has an odd number of electrons: 5 from N and 6 from O, for a total of 11. There is no way that 11 electrons can be distributed between two atoms so that both atoms are surrounded by an even number of electrons.
Notes
 The Lewis dot structure is, according to the text, a "simple theory," and as such, cannot be expected to be correct every time. This is a very interesting take on scientific theories, and students might be challenged to identify other "simple" theories that are not correct all the time.
Keywords
 Lewis structure, electron, valence electron, nitrogen monoxide, covalent bonding
10-00-46un
Title
 Boron gives examples of exceptions to the octet rule
Caption
 To obey the octet rule, BF3 would have to have a double bond between B and one F. F is quite unwilling to share more than one electron, so the double bond won't happen. In BH3, there simply are not enough electrons to provide B with an octet.
Notes
 The Lewis dot structure is, according to the text, a "simple theory," and as such, cannot be expected to be correct every time. This is a very interesting take on scientific theories, and students might be challenged to identify other "simple" theories that are not correct all the time.
Keywords
 Lewis structure, electron, valence electron, boron trifluoride, boron trihydride, covalent bonding
10-00-47un
Title
 SF6 and PCl5: examples of exceptions to the octet rule
Caption
 SF6 and PCl5 can violate the octet rule through the use of empty d orbitals: Both S and P can utilize empty d orbitals to hold pairs of electrons that help bond halogen atoms.
Notes
 The Lewis dot structure is, according to the text, a "simple theory," and as such, cannot be expected to be correct every time. This is a very interesting take on scientific theories, and students might be challenged to identify other "simple" theories that are not correct all the time.
Keywords
 Lewis structure, electron, valence electron, sulfur hexafluoride, phosphorus pentachloride, covalent bonding
10-00-50un
Title
 Attempt to form a sulfur dioxide molecule
Caption
 With six valence electrons each, two oxygen atoms and one sulfur atom can use 18 electrons to satisfy octets. By sharing electrons with each oxygen, both oxygen atoms have satisfied octets, but sulfur does not. We'll need to try a different sharing pattern; a double bond will work.
Notes
 By sharing four electrons between the sulfur and one of the oxygens, all octets would be satisfied.
Keywords
 Lewis structure, electron, valence electron, oxygen, sulfur, sulfur dioxide, double bond
10-00-51un
Title
 Lewis dot structure for a sulfur dioxide molecule
Caption
 With six valence electrons each, two oxygen atoms and one sulfur atom can use 18 electrons to satisfy octets. A double bond will work to satisfy all octets.
Notes
 By sharing four electrons between the sulfur and one of the oxygens, all octets would be satisfied. The selection of the oxygen atom that will double bond is arbitrary: Either one will work, as the two atoms are chemically equivalent.
Keywords
 Lewis structure, electron, valence electron, oxygen, sulfur, sulfur dioxide, double bond
10-00-53un
Title
 Sulfur dioxide's resonance structures
Caption
 The two oxygens are chemically equivalent, so it makes no difference to the molecule which oxygen assumes the double bond. In nature, this equivalence works itself out by the two structures averaging themselves: The measured S-O bonds are intermediate between single and double bonds.
Notes
 By sharing four electrons between the sulfur and one of the oxygens, all octets would be satisfied. The selection of the oxygen atom that will double bond is arbitrary: Either one will work, as the two atoms are chemically equivalent.
Keywords
 Lewis structure, electron, valence electron, oxygen, sulfur, sulfur dioxide, double bond, resonance structure
10-00-54un
Title
 Calculating the number of valence electrons in nitrate ion's Lewis dot structure
Caption
 The total number of 24 comes from the nitrogen atom (five), each oxygen atom (18: six from each O), and the -1 charge (add one). When we know the number of electrons, we can then build the Lewis dot structure in the usual manner.
Notes
 Note that negative ions require us to ADD the number of electrons designated by the charge's magnitude; positive ions require us to SUBTRACT the number of electrons designated by the charge's magnitude.
Keywords
 Lewis structure, electron, valence electron, nitrate ion, covalent bonding, polyatomic ion
10-00-56un
Title
 Attempt to form a nitrate ion
Caption
 With six valence electrons each, three oxygen atoms can use 18 electrons to satisfy octets. The -1 charge adds one more electron, and the nitrogen brings five. By sharing electrons with each oxygen, the nitrogen helps the oxygen atoms satisfy octets. The nitrogen, however, does not have a satisfied octet. We'll need to try a different sharing pattern; a double bond will work.
Notes
 By sharing four electrons between the nitrogen and one of the oxygens, all octets would be satisfied.
Keywords
 Lewis structure, electron, valence electron, oxygen, nitrogen, nitrate ion, double bond, polyatomic ion
10-00-57un
Title
 Lewis dot structure for nitrate ion
Caption
 With six valence electrons each, three oxygen atoms can use 18 electrons to satisfy octets. The -1 charge adds one more electron, and the nitrogen brings five. By sharing electrons with each oxygen, the nitrogen helps the oxygen atoms satisfy octets. If we put a double bond between the N and one of the O's, the N will have a satisfied octet without breaking the oxygen's octets.
Notes
 By sharing four electrons between the N and one of the oxygens, all octets would be satisfied. The selection of the oxygen atom that will double bond is arbitrary: Any one will work, as the three atoms are chemically equivalent.
Keywords
 Lewis structure, electron, valence electron, oxygen, nitrogen, nitrate ion, double bond, polyatomic ion
10-00-59un
Title
 Nitrate ion's resonance structures
Caption
 The three oxygens are chemically equivalent, so it makes no difference to the ion which oxygen assumes the double bond. In nature, this equivalence works itself out by the three structures averaging themselves: The measured N-O bonds are intermediate between single and double bonds.
Notes
 By sharing four electrons between the N and one of the oxygens, all octets would be satisfied. The selection of the oxygen atom that will double bond is arbitrary: Any one will work, as the three atoms are chemically equivalent.
Keywords
 Lewis structure, electron, valence electron, oxygen, nitrogen, nitrate ion, double bond, polyatomic ion, resonance structure
10-00-60un
Title
 Nitrite ion's resonance structures
Caption
 The two oxygens are chemically equivalent, so it makes no difference to the ion which oxygen assumes the double bond. In nature, this equivalence works itself out by the two structures averaging themselves: The measured N-O bonds are intermediate between single and double bonds.
Notes
 By sharing four electrons between the N and one of the oxygens, all octets would be satisfied. The selection of the oxygen atom that will double bond is arbitrary: Either one will work, as the two atoms are chemically equivalent.
Keywords
 Lewis structure, electron, valence electron, oxygen, nitrogen, nitrite ion, double bond, polyatomic ion, resonance structure
10-00-61 un
Title
 Ozone's resonance structures
Caption
 The two terminal oxygens are chemically equivalent, so it makes no difference to the molecule which oxygen assumes the double bond. In nature, this equivalence works itself out by the two structures averaging themselves: The measured O-O bonds are intermediate between single and double bonds.
Notes
 By sharing four electrons between the central O and one of the terminal oxygens, all octets would be satisfied. The selection of the terminal oxygen atom that will double bond is arbitrary: Either one will work, as the two atoms are chemically equivalent.
Keywords
 Lewis structure, electron, valence electron, oxygen, ozone, double bond, resonance structure
10-00-62un
Title
 Lewis dot structure for an oxygen molecule
Caption
 With six valence electrons each, two oxygen atoms can use 12 electrons to satisfy octets. A double bond will work to satisfy both octets.
Notes
 By sharing four electrons between the oxygens, all octets would be satisfied.
Keywords
 Lewis structure, electron, valence electron, oxygen, double bond
10-00-63un
Title
 UV light causes ozone to decompose
Caption
 High-energy UV light breaks bonds in ozone, leaving an oxygen molecule and a single oxygen atom. The unbonded oxygen atom is at high energy, because its octet is not satisfied.
Notes
 Challenge students to speculate on what happens to the unbonded O atom after it is formed.
Keywords
 Lewis structure, electron, valence electron, oxygen, ozone, double bond
10-00-64un
Title
 Why are these Lewis structures incorrect for ozone?
Caption
 The leftmost structure has too many valence electrons: Ozone has 18; the structure shown has 20. The rightmost structure also has 20 valence electrons. Moreover, the octets of the terminal atoms are not satisfied.
Notes
 Students should be given the opportunity to correct these structures.
Keywords
 Lewis structure, electron, valence electron, oxygen, ozone, double bond
10-00-65b
Title
 Carbon dioxide has a linear structure
Caption
 This spacefilling model shows that carbon dioxide has its three atoms arranged in a straight line. If we take the carbon atom as being at the vertex of the angle formed by the three atoms, a 180o angle would be measured.
Notes
 This structure makes sense: The two double bonds are composed of electrons, all of negative charge. The bonds will therefore repel each other, and a 180o angle is the arrangement that gets the electrons as far apart from one another as possible.
Keywords
 Lewis structure, electron, valence electron, oxygen, carbon, carbon dioxide, double bond, linear, geometry
10-00-66c
Title
 Formaldehyde has a trigonal planar structure
Caption
 This spacefilling model shows that formaldehyde has its four atoms arranged in a triangle. If we take the carbon atom as being at the vertex of the angles formed by the three terminal atoms, a 120o angle would be measured.
Notes
 This structure makes sense: The C=O double bond and the two C-H single bonds are composed of electrons, all of negative charge. The bonds will therefore repel each other, and a 120o angle is the arrangement that gets the electrons as far apart from one another as possible.
Keywords
 Lewis structure, electron, valence electron, oxygen, carbon, hydrogen, formaldehyde, double bond, single bond, trigonal planar, geometry
10-00-67un
Title
 Lewis dot structure for a carbon dioxide molecule
Caption
 With six valence electrons from each of two oxygen atoms, and four valence electrons from one carbon atom, 16 electrons are available to satisfy octets. A double bond between each oxygen and the carbon will work to satisfy all octets.
Notes
 By sharing four electrons between the carbon and each of the oxygens, all octets would be satisfied.
Keywords
 Lewis structure, electron, valence electron, oxygen, carbon, carbon dioxide, double bond
10-00-68un
Title
 Lewis dot structure for a formaldehyde molecule
Caption
 With six valence electrons from the oxygen atom, four valence electrons from the carbon atom, and one valence electron from each of two hydrogen atoms, 12 electrons are available to satisfy octets and duets. A double bond between the oxygen and the carbon, and single bonds between C and each H, will work to satisfy all octets and duets.
Notes
 By sharing four electrons between the carbon and the oxygen, all octets would be satisfied.
Keywords
 Lewis structure, electron, valence electron, oxygen, carbon, hydrogen, formaldehyde, double bond, single bond
10-00-69d
Title
 Methane has a tetrahedral structure
Caption
 This ball-and-stick model shows that methane has its five atoms arranged in a tetrahedron. If we take the carbon atom as being at the vertex of the angles formed by the four terminal atoms, a 109.5o angle would be measured.
Notes
 This structure makes sense: The C-H bonds are composed of electrons, all of negative charge. The bonds will therefore repel each other, and a 109.5o angle is the arrangement that gets the electrons as far apart from one another as possible.
Keywords
 Ball-and-stick model, electron, valence electron, carbon, hydrogen, methane, single bond, tetrahedral, geometry
10-00-70e
Title
 Spacefilling model shows that methane has a tetrahedral structure
Caption
 This spacefilling model shows that methane has its five atoms arranged in a tetrahedron. If we take the carbon atom as being at the vertex of the angles formed by the four terminal atoms, a 109.5o angle would be measured.
Notes
 This structure makes sense: The C-H bonds are composed of electrons, all of negative charge. The bonds will therefore repel each other, and a 109.5o angle is the arrangement that gets the electrons as far apart from one another as possible.
Keywords
 spacefilling model, electron, valence electron, carbon, hydrogen, methane, single bond, tetrahedral, geometry
10-00-71un
Title
 Lewis dot structure for an ammonia molecule
Caption
 With five valence electrons from the nitrogen atom and one valence electron from each of three hydrogen atoms, eight electrons are available to satisfy octets and duets. Single bonds between N and each H will work to satisfy all octets and duets.
Notes
 By sharing two electrons between the nitrogen and each hydrogen, all octets and duets would be satisfied.
Keywords
 Lewis structure, electron, valence electron, nitrogen, hydrogen, ammonia, single bond
10-00-72f
Title
 Ammonia's electron geometry is tetrahedral
Caption
 This ball-and-stick model shows that ammonia has its four atoms and one lone pair arranged in a tetrahedron.
Notes
 This structure makes sense: The N-H bonds and the lone pair are composed of electrons, all of negative charge. The bonds and lone pair will therefore repel each other, setting up the tetrahedral geometry.
Keywords
 ball-and-stick model, electron, valence electron, nitrogen, hydrogen, single bond, tetrahedral, lone pair, geometry
10-00-73g
Title
 Ammonia's molecular geometry is trigonal pyramidal
Caption
 This ball-and-stick model shows that ammonia has its four atoms arranged in a trigonal pyramidal geometry.
Notes
 This structure makes sense: The N-H bonds and the lone pair are composed of electrons, all of negative charge. The bonds and lone pair will therefore repel each other, setting up the trigonal pyramidal molecular geometry. Even though the lone pair is not pictured, we can see its influence: The three H atoms (with their N-H bonds) are pushed down and away from the top of the molecule.
Keywords
 ball-and-stick model, electron, valence electron, nitrogen, hydrogen, single bond, trigonal pyramidal, lone pair, geometry
10-00-74h
Title
 Water's electron geometry is tetrahedral
Caption
 This ball-and-stick model shows that water has its three atoms and two lone pairs arranged in a tetrahedron.
Notes
 This structure makes sense: The O-H bonds and the lone pairs are composed of electrons, all of negative charge. The bonds and lone pairs will therefore repel each other, setting up the tetrahedral geometry.
Keywords
 ball-and-stick model, electron, valence electron, oxygen, hydrogen, single bond, tetrahedral, lone pair, geometry
10-00-75i
Title
 Water's molecular geometry is bent
Caption
 This ball-and-stick model shows that water has its three atoms arranged in a bent geometry.
Notes
 This structure makes sense: The O-H bonds and the lone pair are composed of electrons, all of negative charge. The bonds and lone pair will therefore repel each other, setting up the bent molecular geometry. Even though the lone pairs are not pictured, we can see their influence: The two H atoms (with their O-H bonds) are pushed out of a linear geometry.
Keywords
 ball-and-stick model, electron, valence electron, oxygen, hydrogen, single bond, bent, lone pair, geometry
10-00-76un
Title
 Lewis dot structure for a water molecule
Caption
 With six valence electrons from the oxygen atom and one valence electron from each of two hydrogen atoms, eight electrons are available to satisfy octets and duets. Single bonds between O and each H will work to satisfy all octets and duets.
Notes
 By sharing two electrons between the oxygen and each hydrogen, all octets and duets would be satisfied.
Keywords
 Lewis structure, electron, valence electron, oxygen, hydrogen, water, single bond
10-00-77un
Title
 Carbon dioxide is an example of a linear molecule
Caption
 Ball-and-stick models are more reliable than Lewis dot structures in expressing molecular geometry.
Notes
 Students might be asked to summarize how ball-and-stick models differ from Lewis dot structures: What purposes do each serve?
Keywords
 ball-and-stick model, Lewis dot structure, linear, molecular geometry, carbon dioxide
10-00-78un
Title
 Formaldehyde is an example of a trigonal planar molecule
Caption
 Ball-and-stick models are more reliable than Lewis dot structures in expressing molecular geometry.
Notes
 Students might be asked to summarize how ball-and-stick models differ from Lewis dot structures: What purposes do each serve?
Keywords
 ball-and-stick model, Lewis dot structure, trigonal planar, molecular geometry, formaldehyde
10-00-79un
Title
 Sulfur dioxide is an example of a bent molecule
Caption
 Ball-and-stick models are more reliable than Lewis dot structures in expressing molecular geometry.
Notes
 Students might be asked to summarize how ball-and-stick models differ from Lewis dot structures: What purposes do each serve?
Keywords
 ball-and-stick model, Lewis dot structure, bent, molecular geometry, sulfur dioxide
10-00-80un
Title
 Methane is an example of a tetrahedral molecule
Caption
 Ball-and-stick models are more reliable than Lewis dot structures in expressing molecular geometry.
Notes
 Students might be asked to summarize how ball-and-stick models differ from Lewis dot structures: What purposes do each serve?
Keywords
 ball-and-stick model, Lewis dot structure, tetrahedral, molecular geometry, methane
10-00-81un
Title
 Ammonia is an example of a trigonal pyramidal molecule
Caption
 Ball-and-stick models are more reliable than Lewis dot structures in expressing molecular geometry.
Notes
 Students might be asked to summarize how ball-and-stick models differ from Lewis dot structures: What purposes do each serve?
Keywords
 ball-and-stick model, Lewis dot structure, trigonal pyramidal, molecular geometry, ammonia
10-00-82un
Title
 Water is an example of a bent molecule
Caption
 Ball-and-stick models are more reliable than Lewis dot structures in expressing molecular geometry.
Notes
 Students might be asked to summarize how ball-and-stick models differ from Lewis dot structures: What purposes do each serve?
Keywords
 ball-and-stick model, Lewis dot structure, bent, molecular geometry, water
10-00-83un
Title
 Lewis dot structure for a PCl3 molecule
Caption
 With five valence electrons from the phosphorus atom and seven valence electrons from each of three chlorine atoms, 26 electrons are available to satisfy octets. Single bonds between P and each Cl will work to satisfy all octets.
Notes
 By sharing two electrons between the phosphorus and each chlorine, all octets would be satisfied.
Keywords
 Lewis structure, electron, valence electron, phosphorus, chlorine, phosphorus trichloride, single bond
10-00-84un
Title
 Phosphorus in PCl3 has a lone pair
Caption
 With five valence electrons from the phosphorus atom and seven valence electrons from each of three chlorine atoms, 26 electrons are available to satisfy octets. Single bonds between P and each Cl will work to satisfy octets, but the lone pairs provide the additional electrons needed to complete all octets.
Notes
 By sharing two electrons between the phosphorus and each chlorine, all octets would be satisfied. Not all electrons are shared between atoms; those that are not shared, but are still used to satisfy octets, are called lone pairs.
Keywords
 Lewis structure, electron, valence electron, phosphorus, chlorine, phosphorus trichloride, single bond, lone pair
10-00-85un
Title
 Lewis dot structure for nitrate ion
Caption
 With six valence electrons each, three oxygen atoms can use 18 electrons to satisfy octets. The -1 charge adds one more electron, and the nitrogen brings five. By sharing electrons with each oxygen, the nitrogen helps the oxygen atoms satisfy octets. If we put a double bond between the N and one of the O's, the N will have a satisfied octet without breaking the oxygen's octets.
Notes
 By sharing four electrons between the N and one of the oxygens, all octets would be satisfied. The selection of the oxygen atom that will double bond is arbitrary: Any one will work, as the three atoms are chemically equivalent. The square brackets alert us to the fact that nitrate is a polyatomic ion.
Keywords
 Lewis structure, electron, valence electron, oxygen, nitrogen, nitrate ion, double bond, polyatomic ion
10-00-86un
Title
 Lewis dot structure for nitrate ion shows that N has no lone pairs
Caption
 With six valence electrons each, three oxygen atoms can use 18 electrons to satisfy octets. The -1 charge adds one more electron, and the nitrogen brings five. By sharing electrons with each oxygen, the nitrogen helps the oxygen atoms satisfy octets. If we put a double bond between the N and one of the O's, the N will have a satisfied octet without breaking the oxygen's octets. N's octet is satisfied without any lone pairs.
Notes
 By sharing four electrons between the N and one of the oxygens, all octets would be satisfied. The selection of the oxygen atom that will double bond is arbitrary: Any one will work, as the three atoms are chemically equivalent. The square brackets alert us to the fact that nitrate is a polyatomic ion.
Keywords
 Lewis structure, electron, valence electron, oxygen, nitrogen, nitrate ion, double bond, polyatomic ion, lone pair
10-00-87un
Title
 Communicating three-dimensional structures in two dimensions: bond orientation
Caption
 Chemists are aware of the importance of molecular geometry. To communicate the three-dimensional aspects of geometry on paper, chemists have agreed on the symbols shown to indicate whether a bond lies in the paper, or whether it extends behind or above the paper.
Notes
 Students might be given some actual three-dimensional models, and be asked to write the equivalent two-dimensional expression.
Keywords
 bond, molecular geometry, Lewis structure
10-00-88un
Title
 Major molecular geometries expressed in two dimensions
Caption
 These are the five molecular geometries most frequently used in this book.
Notes
 Note that only tetrahedral and trigonal pyramidal geometries are truly three-dimensional.
Keywords
 bond, molecular geometry, Lewis structure, linear, trigonal planar, bent, tetrahedral, trigonal pyramidal
10-01-02j
Title
 Probability map of the O-H bond
Caption
 Because O is more electronegative than H, the electrons in the O-H bond are pulled toward the O end of the bond. The uneven distribution sets up regions of negative and positive charge, producing a measurable dipole moment.
Notes
 The greater the difference in electronegativity of the two bonding atoms, the greater the dipole moment, and the more polar the bond.
Keywords
 dipole moment, oxygen, hydrogen, polar, polar covalent, electronegativity
10-02
Title
 The periodic table shows electronegativity trends
Caption
 Electronegativity of the elements. Linus Pauling introduced the scale shown here. He arbitrarily set the electronegativity of fluorine to 4.0 and computed all other values relative to fluorine.
Notes
 Electronegativity trends follow the trends for metallic character: the lower the electronegativity, the higher the metallic character.
Keywords
 periodic table, electronegativity, Pauling
10-03
Title
 Spacefilling model of a chlorine molecule
Caption
 In Cl2, the two Cl atoms share the electrons evenly. This is a pure covalent bond.
Notes
 The electron sharing is even because the chlorine atoms have the same electronegativity
Keywords
 electronegativity, covalent bond, chlorine
10-04
Title
 Spacefilling model of a sodium chloride molecule
Caption
 In NaCl, Na completely transfers an electron to Cl. This is an ionic bond.
Notes
 Na has low electronegativity; Cl has high electronegativity. The Cl will completely remove from Na its only valence electron. The result is that both atoms will have full octets. The positive and negative ions that form attract one another to form the compound.
Keywords
 electronegativity, ionic bond, chlorine, sodium, sodium chloride
10-05
Title
 Spacefilling model of a hydrogen fluoride molecule
Caption
 In HF, the electrons are shared, but the shared electrons are more likely to be found on the F than on the H. The bond is polar covalent.
Notes
 In HF, the two atoms share electrons, forming a covalent bond. However, H has lower electronegativity than F, so the electrons will be pulled toward the F atom. The result is that the molecule has a dipole moment.
Keywords
 electronegativity, polar covalent, fluorine, hydrogen, hydrogen fluoride, dipole moment, polar
10-06
Title
 A spectrum of chemical bonding
Caption
 The type of bondÑpure covalent, polar covalent, or ionicÑis related to the electronegativity difference between the bonded atoms.
Notes
 The greater the difference in electronegativity of the bonding atoms, the more ionic the bonding. Students might be asked to speculate on why there is a "pure covalent" bond, but not a "pure ionic" bond. Students should focus their responses on electronegativity.
Keywords
 electronegativity, polar covalent, covalent, ionic, dipole moment, polar
10-06-02k
Title
 Carbon dioxide is a nonpolar molecule
Caption
 Despite having two polar covalent bonds, carbon dioxide is nonpolar. The reason for this is the molecule's linear geometry: The electrons are pulled in equal but opposite directions, causing them to cancel one another.
Notes
 For a molecule to be polar, bond polarity must work together with geometry.
Keywords
 electronegativity, polar covalent, covalent, dipole moment, polar
10-06-03l
Title
 Water is a polar molecule
Caption
 Water has two polar covalent bonds, but the polar covalent bonds pull electrons to the oxygen's region of the molecule, to make the molecule polar. The reason for this is the molecule's bent geometry: The electrons are not pulled in equal and opposite directions, despite the bonds being identical.
Notes
 For a molecule to be polar, bond polarity must work together with geometry.
Keywords
 electronegativity, polar covalent, covalent, dipole moment, polar
10-06-04m
Title
 Adding dipole moments to determine if a molecule is polar
Caption
 Molecules will be nonpolar if the molecular geometries cause the dipole moments to cancel out.
Notes
 The scheme shown applies in cases where the bonds are identical. If any of the bonds in the molecule differ, the dipole moments will not cancel, and the molecule will be polar.
Keywords
 bond, molecular geometry, linear, trigonal planar, bent, tetrahedral, trigonal pyramidal, dipole moment, polar
10-06-06n
Title
 Ammonia is a polar molecule
Caption
 Ammonia has three polar covalent bonds, but the polar covalent bonds pull electrons to the nitrogen's region of the molecule, to make the molecule polar. The reason for this is the molecule's trigonal pyramidal geometry: The electrons are not pulled in equal and opposite directions, despite the bonds being identical.
Notes
 For a molecule to be polar, bond polarity must work together with geometry.
Keywords
 electronegativity, polar covalent, covalent, dipole moment, polar, ammonia, trigonal pyramidal
10-07
Title
 Polar molecules attract one another
Caption
 Just as the north pole of one magnet is attracted to the south pole of another, so the positive end of one molecule with a dipole is attracted to the negative end of another molecule with a dipole.
Notes
 Water is the example here, but this is a general result. Challenge students to speculate on the effect of the attraction on properties such as solubility and boiling point.
Keywords
 electronegativity, polar covalent, covalent, dipole moment, polar
10-09-02p
Title
 Which end of this detergent molecule is polar?
Caption
 C-H bonds have a very low dipole moment, but C-O and O-H bonds have a higher dipole moment. The higher dipole moments are on the right side of the molecule; the right side will be polar.
Notes
 Students might be challenged to explain how detergents simultaneously interact with water and grease.
Keywords
 electronegativity, polar covalent, covalent, dipole moment, polar
10-09-03un
Title
 Lewis dot structure for sulfur
Caption
 The Lewis structure uses dots to represent the valence electrons. This system facilitates visualization of the number of electrons an atom needs to satisfy its octet. In this case, sulfur has six valence electrons.
Notes
 The Lewis structure is somewhat deceptive, in that it suggests that electrons are particles (dots). The wave-particle duality is ignored.
Keywords
 Lewis structure, electron, valence electron, sulfur
10-09-04un
Title
 Li and Br react with one another by forming ions
Caption
 These two elements will form ions: The Li atom will give up its valence electron to Br. In doing so, Li's electron configuration becomes like that of a noble gas (low energy), and so does Br's. This is an example of ionic bonding.
Notes
 Li+ has an electron configuration 1s2 Br- has an electron configuration 1s22s22p63s23p64s23d104p6 The Li+ is at low energy because it has the same electron configuration as helium. The Br- is at low energy because its electron configuration is the same as krypton's. Getting to low energy is one reason elements undergo chemical reactions.
Keywords
 Lewis structure, electron, valence electron, lithium, bromine, lithium bromide
10-09-05un
Title
 Lewis dot structures for K and S
Caption
 These two elements will form ions: The K atom will give up its valence electron to S. In doing so, K's electron configuration becomes like that of a noble gas (low energy), But, S requires two electrons, not one, so a second K must give its valence electron to S in order for S to get to low energy. The resulting formula will be K2S. This is an example of ionic bonding.
Notes
 K+ has an electron configuration 1s22s22p63s23p6 S2- has an electron configuration 1s22s22p63s23p6 As ions, both elements have full octets, and are therefore at low energy. Getting to low energy is one reason elements undergo chemical reactions.
Keywords
 Lewis structure, electron, valence electron, potassium, sulfur
10-09-06un
Title
 K and S react with one another by forming ions
Caption
 These two elements will form ions: The K atom will give up its valence electron to S. In doing so, K's electron configuration becomes like that of a noble gas (low energy), But, S requires two electrons, not one, so a second K must give its valence electron to S in order for S to get to low energy. The resulting formula will be K2S. This is an example of ionic bonding.
Notes
 K+ has an electron configuration 1s22s22p63s23p6 S2- has an electron configuration 1s22s22p63s23p6 As ions, both elements have full octets, and are therefore at low energy. Getting to low energy is one reason elements undergo chemical reactions.
Keywords
 Lewis structure, electron, valence electron, potassium, sulfur
10-09-09un
Title
 Lewis dot structure for a carbon disulfide molecule
Caption
 With six valence electrons from each of two sulfur atoms, and four valence electrons from one carbon atom, 16 electrons are available to satisfy octets. A double bond between each sulfur and the carbon will work to satisfy all octets.
Notes
 By sharing four electrons between the carbon and each of the oxygens, all octets would be satisfied.
Keywords
 Lewis structure, electron, valence electron, sulfur, carbon, carbon disulfide, double bond
10-09-10un
Title
 Selenium dioxide's resonance structures
Caption
 The two oxygens are chemically equivalent, so it makes no difference to the molecule which oxygen assumes the double bond. In nature, this equivalence works itself out by the two structures averaging themselves: The measured Se-O bonds are intermediate between single and double bonds.
Notes
 By sharing four electrons between the selenium and one of the oxygens, all octets would be satisfied. The selection of the oxygen atom that will double bond is arbitrary: Either one will work, as the two atoms are chemically equivalent.
Keywords
 Lewis structure, electron, valence electron, oxygen, selenium, selenium dioxide, double bond, resonance structure
10-09-12un
Title
 SeO2 is a polar molecule
Caption
 SeO2 has two polar covalent bonds, but the polar covalent bonds pull electrons to the oxygen's region of the molecule, to make the molecule polar. The reason for this is the molecule's bent geometry: The electrons are not pulled in equal and opposite directions, despite the bonds being identical.
Notes
 For a molecule to be polar, bond polarity must work together with geometry.
Keywords
 electronegativity, polar covalent, covalent, dipole moment, polar, selenium, selenium dioxide
10-09-13un
Title
 Lewis structures in Chapter 10, problem 41
Caption
 Find the errors in these Lewis structures.
Notes
 errors present
Keywords
 Lewis structure, electron, valence electron, ion, ionic bonding
10-09-14un
Title
 Lewis structures in Chapter 10, problem 42
Caption
 Find the errors in these Lewis structures.
Notes
 errors present
Keywords
 Lewis structure, electron, valence electron, ion, ionic bonding
10-09-15un
Title
 Lewis structures in Chapter 10, problem 49
Caption
 Find the errors in these Lewis structures.
Notes
 errors present
Keywords
 Lewis structure, electron, valence electron, ion, ionic bonding
10-09-16un
Title
 Lewis structures in Chapter 10, problem 50
Caption
 Find the errors in these Lewis structures.
Notes
 errors present
Keywords
 Lewis structure, electron, valence electron, ion, ionic bonding
10-09-17un
Title
 Reactions in Chapter 10, problem 96
Caption
 Write the Lewis structures for the reactants and products.
Notes
 Some of the reactants and products do not obey the octet rule; these are called free radicals.
Keywords
 Lewis structure, electron, valence electron, free radicals
10-09-18q
Title
 DNA in Chapter 10, problem 95
Caption
 Free radicals, molecules containing unpaired electrons, may attack biological molecules.
Notes
 The problem asks the student to write Lewis structures for some free radicals.
Keywords
 Lewis structure, electron, valence electron, free radicals, DNA
10-09-21t
Title
 Chapter 10 Problem 97
Caption
 Is the pictured spacefilling model correct for H2Se?
Notes
 No; H2Se is bent.
Keywords
 spacefilling model, electron, valence electron, molecular geometry
10-09-22u
Title
 Chapter 10 Problem 97
Caption
 Is the pictured spacefilling model correct for CSe2?
Notes
 Yes; CSe2 is linear.
Keywords
 spacefilling model, electron, valence electron, molecular geometry
10-09-23v
Title
 Chapter 10 Problem 97
Caption
 Is the pictured spacefilling model correct for PCl3?
Notes
 No; PCl3 is trigonal pyramidal.
Keywords
 spacefilling model, electron, valence electron, molecular geometry
10-09-24w
Title
 Chapter 10 Problem 97
Caption
 Is the pictured spacefilling model correct for CF2Cl2?
Notes
 Yes; CF2Cl2 is tetrahedral.
Keywords
 spacefilling model, electron, valence electron, molecular geometry
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