Explain each of the following observations in terms of the electronic strucutre and/ or bonding of the compunds involved.
(a)At ordinary conditions, HF( normal boiling point=20C) is a liquid, whereas HCl(normal boiling point=-114C) is a gas. (b)Molecules of AsF3 are polar, whereas molecules of AsF5 are nonpolar.
(c)The N-O bonds in the NO2- ion are equal in length,whereas they are unequal in HNO2.
(d)For sulfur, the fluorides SF2,SF4, and SF6 are known to exist, whereas for oxygen only OF2 is known to exist.
What are you thinking on these? I don't want to do your homework for you. But I can help point you in the right direction if you can be specific about what you don't understand.
It is difficult to draw Lewis dot structure on these boards. Do you know how ro draw them. If you do, draw the one for HNO2 as well as the one for NO2^-. YOu will see that the HNO2 has a single bond from N to O and another from O to H whereas the other O is double bonded to N. So HNO2 has a single bond to one O and a double bond to the other. For the nitrite ion, you will see a double bond to one O and a single bond to the other; however, you may choose to draw the double bond to EITHER O and the single bond to EITHER O. These are resonanting structures and the N to O bond distance is a hybrid (an average) of the two structures and that makes them the same distance. Here is a site for Lewis structure of nitrite ion and another for the Lewis structure of HNO2. Remember that a line is the same as two dots. Also, to help you count the electrons, there are 18 for nitrite and 18 for HNO2.
http://en.wikipedia.org/wiki/Nitrite
http://en.wikipedia.org/wiki/Nitrous_acid
Let me know if you have any questions.
I need help with a
lewis structure
what the name of compound whicg are the boiling point are 114
do you have ms larriba?
HAHAHA I HAVE LARRIBA. that's funny, i'm glad i'm not the only one looking up the homework answers.
(a) The difference in boiling points of HF and HCl can be explained by the strength of intermolecular forces between molecules. In the case of HF, it forms hydrogen bonds due to the large electronegativity difference between hydrogen and fluorine. These hydrogen bonds are relatively strong, leading to higher intermolecular forces and making it harder for the molecules to escape and boil.
On the other hand, HCl does not form hydrogen bonds but rather exhibits weaker dipole-dipole interactions. Consequently, the intermolecular forces in HCl are weaker, allowing the molecules to escape and boil more easily. Therefore, HF exists as a liquid at ordinary conditions, whereas HCl is a gas.
(b) The polarity of a molecule is determined by the presence of a net dipole moment. In AsF3, the fluorine atoms are more electronegative than the arsenic atom, resulting in partial negative charges on the fluorine atoms and a partial positive charge on the arsenic atom. AsF3 has a trigonal pyramidal molecular geometry, which is asymmetric and contributes to the overall dipole moment, making AsF3 a polar molecule.
In contrast, AsF5 adopts a trigonal bipyramidal molecular geometry, which leads to symmetry in the molecule. This symmetry cancels out the dipole moments of individual As-F bonds, resulting in a net dipole moment of zero. Consequently, AsF5 is a nonpolar molecule.
(c) In the NO2- ion, the central nitrogen atom has a single bond with one oxygen atom and a double bond with the other oxygen atom. Due to resonance, the double bond can shift between the two oxygen atoms. As a result, the two N-O bonds in NO2- are equivalent in length, resulting in a symmetrical structure.
In HNO2, the nitrogen atom is bonded to one oxygen atom through a single bond and has a single bond with a hydrogen atom. The electron distribution is not delocalized through resonance in HNO2, leading to an unequal distribution of electron density. This unequal distribution results in a shorter N-O bond with the oxygen atom and a longer N-O bond with the hydrogen atom, making the N-O bonds in HNO2 unequal in length.
(d) The ability of elements to form stable compounds depends on their electron configurations and the availability of valence electrons for bond formation. In the case of sulfur, it has six valence electrons, allowing it to expand its octet by forming multiple bonds with fluorine atoms.
Sulfur can form SF2 by having a single bond with two fluorine atoms, SF4 by having a double bond with one fluorine atom and single bonds with the remaining three, and SF6 by having a triple bond with one fluorine atom and single bonds with the remaining five. The ability of sulfur to accommodate various numbers of fluorine atoms allows the existence of SF2, SF4, and SF6.
On the other hand, oxygen has only two valence electrons, limiting its ability to form multiple bonds. Therefore, the only known fluoride of oxygen is OF2, where oxygen forms a double bond with a fluorine atom. The electron configuration of oxygen does not allow the formation of additional fluorides like SF4 or SF6.