Presentation on theme: " Ionic bonds › Strongest (melting point?) › Charged ions Metallic bonds › Strong, but weaker than ionic › “Sea of electrons” – all shared among all."— Presentation transcript:
Ionic bonds › Strongest (melting point?) › Charged ions Metallic bonds › Strong, but weaker than ionic › “Sea of electrons” – all shared among all of them Covalent bonds › Weakest (melting point?) › Probably the most important substances of all are molecules
When you need to draw a Lewis structure, always list the number of electrons you have first. This is especially important for when you have polyatomic ions (charged particles). SO 4 2- is a polyatomic ion – what would its Lewis structure be?
First, how many electrons? › S 6 › O 6 x 4 = 24 › Charge 2 extra e - › TOTAL = 32 Place S in the center, surrounded by the 4 O’s. Make a single bond first on each one. Then place electrons until you have a model that obeys the octet rule.
Let’s try one more for ammonium. › N 5 › H 1 x 4 = 4 › Charge -1 e - (a positive charge means we’re short electrons) › Total – 1 = 8
Remember we can use electronegativity to calculate the polarity of a bond? We find the electronegativity values for each atom in the bond on the periodic table (p. 174) Then we subtract the smaller number from the bigger one – the difference tells us how polar the bond is. nonpolar polar mostly ionic
Consider the compound HCl. Is this bond nonpolar covalent, polar covalent, or mostly ionic? › Hydrogen’s electronegativity is 2.2 › Chlorine’s electronegativity is 3.19 › What is the difference? What about CCl?
Has anyone seen the movie Casino Royale ? Know the name of the main girl? Vesper. Now I have to ask: what are the odds that a scriptwriter would just happen to put a guy named Bond with a girl named Vesper? Clearly that writer knew their chemistry.
Say that 5 times fast! What does that name tell you about the theory? It has to do with valence shell electrons, it seems they come in pairs, and they repel each other. Why would electrons repel?
VSEPR theory only applies to molecules (substances with only covalent bonds). Its difficult to see VSEPR shapes on paper, so I’ll be showing you some shapes with a modeling kit.
In a molecule, there are two types of electron pairs: bonding pairs (those used to form bonds), and lone pairs (those not used to form bonds) For counting purposes, double bonds and triple bonds count as a single electron pair. Each electron pair wants to be as far away from all the others as possible. As you’ll see, this results in some very common shapes.
Electronegativity can tell us more than just the polarity of a bond. It can tell us if a molecule will be polar overall. First, consider CH 4. What is the geometry? What is the polarity of each bond in the molecule? Is the molecule polar? Look at CH 3 Cl. Calculate the polarity of each bond in the molecule. Is the molecule polar? Last one: CCl 4. Are the bonds polar? Is the molecule polar?
No! If all the polar bonds in a molecule cancel out, the molecule will be nonpolar.
Determine the geometry of each of the following molecules, then determine if the molecules are polar or nonpolar overall. (Remember, a compound can have polar bonds and still be nonpolar!) › H 2 O XeF 2 › SO 2 NH 3 › BF 3 › (the ones in bold don’t obey the octet rule for the central atom)
P. 181 (4, 6, 8, 10, 14, 20) P. 225 (1, 5, 7, 10, 12, 14b,c,e, 19) This assignment will be due the day before the unit test – you should have at least a week to finish it.