Thursday, December 14, 2006
1214
Honors: Lewis structures and more Lewis structures today. We learned about resonance structures which we draw and separate with double arrows. Resonance structures represent the sharing of certain valence electrons among several bonding nuclei throughout the molecule. Whenever you need to double bond or triple bond the central atom with one of several possible terminal atoms, you draw a resonance structure for each possible double or triple bonding pair of atoms.
We also drew structures for some polyatomic ions. Just keep in mind that Lewis structures for ANY ions, whether monatomic or polyatomic ions, ALWAYS get put in brackets with the magnitude of the charge FIRST and the sign of the charge NEXT outside the brackets. We also started to predict molecular and electronic geometry of a molecule or ion based on its Lewis structure. That will be the main consideration when we discuss molecular polarity.
By the way, some of you may not know the meaning of the word "polar". In covalent bonds, polar means that a partial negatively charged "pole" develops near the more electronegative atom and a partial positively charged "pole" develops near the less electronegative atom of the bonding atoms. So, just as the Earth (and our school) has a North "pole" and a South "pole", atoms that unequally share negative electrons develop opposite electrically charged "poles" at opposite ends of the bonding region.
Regents: we did some in-class review and learned about metallic bonding and how it accounts for:
1. the good electrical conductivity of metals due to the low Zeff on the valence electrons
2. the high melting point of metals due to the relatively strong metallic bonds due to the simultaneous attraction for all neighboring valence electrons among nearby metal atoms' nuclei. This multiple attraction creates strong bonding.
3. the malleability of metals; when you hammer the metal atom lattice, the mobile electrons just readjust their positions so that they are still between and attracted to multiple metal atoms' nuclei; the metal lattice doesn't break off.
We then contrasted a metal lattice with an ionic lattice of a salt. Salts are not malleable because when the lattice is hit with a hammer or mallet, the row or rows of ions will line up with like charges next to each other, which causes electronstatic repulsion which will make the row or rows of ions "break off' so the salt will crumble/be brittle. There will be a brief quiz tomorrow on ionic and molecular compounds, formulas, ionic charges, and Lewis structures. Our next full period exam will be given next Tuesday.
AP: we reviewed electronic and molecular geometry and related that to molecular polarity. We predicted intermolecular forces of attraction types based on molecular polarity. We reviewed delocalized bonding and counted the number of pi and sigma bonds in a given molecule or ion.
We started the seminal Molecular Orbital Theory. We STARTED it. Don't freak out if you don't get it yet. There isn't much to get! We haven't even gotten into the application of the theory! The practical questions that you will see are as easy if not EASIER than following the regular Aufbau principle in orbital diagrams of atoms. You'll see tomorrow as we go through the lesson using the new supercharged iBook!
We also drew structures for some polyatomic ions. Just keep in mind that Lewis structures for ANY ions, whether monatomic or polyatomic ions, ALWAYS get put in brackets with the magnitude of the charge FIRST and the sign of the charge NEXT outside the brackets. We also started to predict molecular and electronic geometry of a molecule or ion based on its Lewis structure. That will be the main consideration when we discuss molecular polarity.
By the way, some of you may not know the meaning of the word "polar". In covalent bonds, polar means that a partial negatively charged "pole" develops near the more electronegative atom and a partial positively charged "pole" develops near the less electronegative atom of the bonding atoms. So, just as the Earth (and our school) has a North "pole" and a South "pole", atoms that unequally share negative electrons develop opposite electrically charged "poles" at opposite ends of the bonding region.
Regents: we did some in-class review and learned about metallic bonding and how it accounts for:
1. the good electrical conductivity of metals due to the low Zeff on the valence electrons
2. the high melting point of metals due to the relatively strong metallic bonds due to the simultaneous attraction for all neighboring valence electrons among nearby metal atoms' nuclei. This multiple attraction creates strong bonding.
3. the malleability of metals; when you hammer the metal atom lattice, the mobile electrons just readjust their positions so that they are still between and attracted to multiple metal atoms' nuclei; the metal lattice doesn't break off.
We then contrasted a metal lattice with an ionic lattice of a salt. Salts are not malleable because when the lattice is hit with a hammer or mallet, the row or rows of ions will line up with like charges next to each other, which causes electronstatic repulsion which will make the row or rows of ions "break off' so the salt will crumble/be brittle. There will be a brief quiz tomorrow on ionic and molecular compounds, formulas, ionic charges, and Lewis structures. Our next full period exam will be given next Tuesday.
AP: we reviewed electronic and molecular geometry and related that to molecular polarity. We predicted intermolecular forces of attraction types based on molecular polarity. We reviewed delocalized bonding and counted the number of pi and sigma bonds in a given molecule or ion.
We started the seminal Molecular Orbital Theory. We STARTED it. Don't freak out if you don't get it yet. There isn't much to get! We haven't even gotten into the application of the theory! The practical questions that you will see are as easy if not EASIER than following the regular Aufbau principle in orbital diagrams of atoms. You'll see tomorrow as we go through the lesson using the new supercharged iBook!
# posted by C @ 5:52 PM 
