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Molecular Geometry

Molecular Geometry. It’s all about the Electrons. Electrons decide how many bonds an atom can have They also decide the overall shape of the molecule OPPOSITES ATTRACT!. Lewis Structures. A Lewis structure is basically a diagram of how a molecule looks using dots to represent the electrons.

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Molecular Geometry

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  1. Molecular Geometry

  2. It’s all about the Electrons • Electrons decide how many bonds an atom can have • They also decide the overall shape of the molecule • OPPOSITES ATTRACT!

  3. Lewis Structures • A Lewis structure is basically a diagram of how a molecule looks using dots to represent the electrons. • There are 4 rules for making these structures and that is where the electrons come into play.

  4. Rule number 1 • Count the Number of valence electrons! • This means of all the atoms present • With a polyatomic anion, add one for each negative charge • With a polyatomic cation, subtract one for each positive charge • Ex: CO2 • C: 4 O: 6 6 + 6 + 4 = 16

  5. Rule Number 2 • Draw a “skeleton structure” for the molecule using all single bonds • This will most often be one central atom with several surrounding ones • Typically the central atom is written first • Ex: CO2 O-C-O

  6. Rule Number 3 • Determine the number of valence electrons still available for distribution • To do this simply deduct two valence electrons for each single bond written in step two • Ex: CO2 two single bonds so far so we subtract a total of 4 • 16 – 4 = 12

  7. Fourth Rule • Determine the number of electrons required to fill an octet for each atom • If this equals the number of electrons left, then place them on the atoms as unshared pairs • If the number of electrons available is less than the number needed then you need to make double or triple bonds in place of the single bonds

  8. CO2 (again) • O – C – O • So far we have used four electrons so 16 – 4 = 12 • Carbon still needs 4 more electrons and each Oxygen needs 6 more. 6 + 6 + 4 = 16 • But we only have 12 left so lets make some double bonds! • O – C – O becomes O = C = O

  9. Now Carbon doesn’t need anymore electrons and the Oxygen’s only need 4 more each. Since we used 4 electrons to make those into double bonds we now have exactly 8 electrons left. • Now we simply distribute them to the Oxygen atoms as unshared paired electrons.

  10. Practice!

  11. Resonance! • Resonance is invoked whenever a single Lewis structure does not adequately reflect the properties of a substance • In other words, resonance comes into play when you can make two structures that are the same in their placement of atoms but different in the bonds • SO2

  12. Resonance structures are NOT forms where the electrons move eternally between them • Resonance structures are equally plausible or they are not a resonance structure • Resonance forms differ in their distribution of electrons, NOT in their arrangement of atoms! • So just because a formula for a compound is the same it does not mean that it is a resonance structure

  13. Electronegativity • Electronegativity is a measure of how much an element wants to pull electrons towards itself • This is represented as a unit-less number ranging from 0 – 4.0 • Here’s a handy reference sheet with all the values. Guard it with your LIFE!

  14. So what? • These numbers can be used mathematically to know if a bond is ionic or covalent • It can also tell you if a covalent bond is more polar or less polar (more on polarity in a minute) • So all we have to do is subtract one from the other.

  15. Example • Fluorine has an electronegativity of 4.0 • Sodium has an electronegativity of 0.9 • 4.0 – 0.9 = 3.1 • So what does that mean? • It means that it is an ionic bond! • This makes sense since we know that a bond involving one metal and one non-metal is ionic.

  16. Example two • Fluorine has an electronegativity of 4.0 • Carbon has an electronegativity of 2.15 • 4.0 – 2.5 = 1.5 • This makes this bond covalent!

  17. Sharing is caring, but some elements are greedy! • This “greediness” shown by some elements like fluorine leads us to the next piece of this puzzle • The more unequal the sharing of electrons is in a bond, the more polar it is. • The smaller that difference in electronegativity, the less polar.

  18. Polar vs Non-polar • Polar: • Number greater than 0.4 • Unequal sharing of electrons • Water is an example • Non-Polar: • Number less than 0.4 • Equal sharing of electrons • Methane (CH4) is an example

  19. So why is this important? • Polarity is a major component of organic chemistry • Polarity also explains why certain substances can dissolve other substances while others cannot • Think oil and water

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