Covalent Bonding & Molecular Geometry

1. Covalent Bonding & Molecular Geometry Mrs. Daniels .2 Chemistry November 2006

2. Rules for Lewis Dot Structures • 1. Count and add up total number of valence electrons • 2. Choose your central atom (usually the first element in the formula…but never hydrogen) and arrange remaining atoms around the central atom • 3. Draw two electrons between each atom (representing the electron pair that they are sharing) • 4. Place remaining electrons around the atoms satisfying their outer shells (start with outer atoms…in case you run out!)

3. Rules for Lewis Dot Structures • What if you run out of electrons? • Can you have a double bond or a triple bond? • Can atoms share more than one pair? • How many pairs are they sharing in a double bond? • 2 • How many pairs are they sharing in a triple bond? • 3

4. Lewis Dot with Multiple Bonds • Determine how many of each atom you will need to form a molecule. Then draw the Lewis Dot structures for each of the following: • H + Cl --> • P + Cl --> • H + S --> • Cl + Cl -->

5. Lewis Dot with Multiple Bonds • Determine how many of each atom you will need to form a molecule. Then draw the Lewis Dot structures for each of the following: • H + Cl --> HCl • P + Cl --> PCl3 • H + S --> H2S • Cl + Cl --> Cl2 • Which of the above have polar covalent bonds? What is unique about Cl2?

6. Diatomic Molecules • First of all, we’ve been using the term “molecule” • What is a molecule? • A covalently bonded compound • A diatomic molecule is one that is made up of two of the same atoms • Since they are the same atom and have equal electronegativities, they are 100% non polar (and are the only molecules who are)

7. Back to Lewis Dot Structures • For all of those math minded people, let’s look at an easy way to determine the number of bonds that will be formed: • Determine the #of electrons needed to satisfy each atom’s outer shell (most want 8) • Then add up the # of electrons available • Subtract the two and the answer will tell you how many electrons must be SHARED N - A = # Shared (S)

8. Ex. Carbon Dioxide CO2 • Needed: Each atom wants 8, so 8 x 3 = 24 • Available: Carbon has 4 & each oxygen has 6, so 4 + 6(2) = 16 • 24-16 = 8 electrons must be shared • Put Carbon in the center of your board • Draw an oxygen on either side and draw the C sharing half of the 8 with each O O::C::O Then, fill in the rest of your available dots

9. Now try SO3 • N = 32 • A = 24 • Shared = 8 electrons • What will be your central atom? • Now draw 3 oxygens around it and place your shared electrons • What type of bonds do you have?

10. Resonance • Does it matter which Sulfur to oxygen bond you drew as a double bond? • No, it could have been on any of them • These are called resonance structures

11. Polyatomic ion revisit • How well do you know your polyatomic ions? • Now, you’re going to draw Lewis dot structures of them • How does this differ from what you’ve been doing? • The only difference is in your number of available electrons (add in those that you’ve gained or subtract those you lost)

13. Covalent Bonding • According to electronegativity, what determines the type of bond that will form? • 2+ = ionic • 1.7 or less = covalent • 1.7 - 2 = two nonmetals = covalent metal and nonmetal = ionic

14. Covalent Bonding • What is the difference between ionic and covalent bonds? • Ionic bonds involve an actual transfer of electrons from one atom to another, thus creating charged particles or “ions” • Covalent bonds do NOT involve a transfer…instead atoms SHARE electrons

15. Covalent Bonding • Have you ever shared something with someone? • Was it perfectly equal sharing? • Sometimes atoms don’t share equally • One atom will have possession of the electron more of the time than the other • This situation is called bond polarity

16. Polar Covalent • If the difference in electronegativites is 0-0.5 = then the bond is nonpolar covalent • In other words, they share pretty equally • If the difference in electronegativities is 0.5-1.69 = then the bond is polar • The more electronegative atom has a stronger pull and will have the electron more often than the less electronegative atom

18. Covalent Bonds - Polar or Not? • Use the table of electronegativities to determine whether the following bonds are polar or non-polar covalent: • O with O • P with H • Se with F • N with O • C with H • S with O

20. Money Activity • Okay, you’ve done it once with ionic bonding…now let’s do it with covalent. • This time when you bond, you’ll have to “share” your valence electrons. In order to demonstrate sharing, you must both have your hands on the “electrons” at the same time.

22. Molecular Geometries • VSEPR - stands for Valence Shell Electron Pair Repulsion • What charge is an electron? • So what do negative charges do with each other? • This model just puts that into practice…the electron pairs are going to repel each other and be as far apart as possible

23. http://www.shef.ac.uk/chemistry/vsepr

24. Molecular Geometries • Recall that when you draw Lewis Dot structures, the first atom written is usually the central atom • We refer to this as “A” • The atoms attached to the central atom directly are referred to as “X” • Any other electrons that are unshared around the central atom are called “E”

25. Molecular Geometries • Add these to your geometric shape table • A2 and AX2 are linear • AX3 are trigonal planar • AX2E are bent • AX4 are tetrahedral • AX3E are trigonal pyramidal • AX2E2 are bent

26. Let’s Try a Few: • NH3 - • A: one central atom (Nitrogen) • X: three atoms attached to central atom • E: one lone pair of electrons (unshared) • So, its AXE designation would be AX3E

27. H2O - • Oxygen is the central atom (A) • Two hydrogens are attached to central atom (X2) • Two unshared (lone) pairs (E2) • So, AX2E2

28. Try CCl4 • There are no electrons unshared on the central atom • AX4 • Try NCl3 • AX3E

30. Polar and Nonpolar molecules • We already know how to determine whether or not a bond is polar or non-polar • Recall…what is polar? • Unequal sharing of electrons • If one atom has a stronger pull on the electrons, they will have a partial negative charge (the other atom will have a partial positive charge)

31. Polar and Nonpolar molecules • Draw the Lewis dot structure for BCl3 • Is the B - Cl bond polar or nonpolar? • The electronegativity of B is 2.01 • The electronegativity of Cl is 3.00 • The difference is 0.99 and is therefore polar • Is the molecule polar?

32. Polar and Nonpolar molecules • What is the molecular geometry of BCl3? • Draw the Lewis dot structure • Trigonal planar is the molecular geometry • Draw arrows showing the pull of electrons and the partial charges? • Do they cancel each other out? • Yes, so the molecule is nonpolar (even though it has polar bonds within it)

33. Polar and Nonpolar molecules • Let’s look at another example: • Ex. In NH3, the N - H bond is polar • The electronegativity of N is 3.07 • The electronegativity of H is 2.20 • The difference is 0.87, which is greater than 0.5 and is therefore polar • So, is the molecule polar or nonpolar?

34. Polar and Nonpolar molecules • Let’s think about the molecular geometry • Draw the Lewis dot structure for NH3 on your white board • There is a lone pair of electrons off the N • Draw an arrow showing the partial charges in each of the polar bonds • Is there equal pulling in opposite directions? • Do the charges cancel each other out?

35. Polar and Nonpolar molecules • No, there is an overall partial negative charge at the top and an overall partial positive charge at the bottom • Therefore, the molecule is polar

36. Polar and Nonpolar Molecules • So… • Can you have a nonpolar molecule that has polar bonds in it? • YES • But can you have a polar molecule that has only nonpolar bonds? • NO • There is NO pulling of electrons if they are non polar

38. Let’s go back and review atomic structure: • Where are the protons? • Where are the neutrons? • Where are the electrons? • Do we know at any given moment EXACTLY where the electrons are in an atom? • NO, we know where the highest probability is for them to be found (these are the orbitals s, p, d, and f) • What are the electrons doing in these orbitals?

39. Are the electrons in a molecule doing anything differently? • No, they may be SHARED between two atoms in a molecule in a covalent bond, but they are still buzzing about • Is it possible for a majority of the electrons in a molecule to be found on the same side of the molecule? • Yes, what would that do (temporarily) to that side of the molecule?

40. Intermolecular Forces • Part of the molecule can be partially negative for a short time while the electrons are on that side • The other side would then be partially positive • Even though they are weak charges, what will opposite charges do? • Can a partial positive of one molecule be attracted to a partial negative on another?

41. Intermolecular Forces • When this occurs, it is called induced dipole(a.k.a. London dispersion forces) • Dipole means that there are two poles (+ and -) • Where else have we seen dipoles? • In polar covalent bonds between two atoms • When the + pole of one molecule is attracted to the - pole of another molecule, the attraction or force is called a dipole-dipole

42. Intermolecular Forces • Hydrogen bonding is the third type of intermolecular force (really it is a very strong version of a dipole-dipole) • When an atom of hydrogen bonds to fluorine, oxygen, or nitrogen (FON) there is a large difference in electronegativity • The electrons spend more time around the electronegative atom and less time around the less electronegative atom

43. Hydrogen Bonding • The atom who has “possession” of the electron the majority of the time takes on a PARTIAL negative charge • (remember it does not own the electron, so it is not an ion) • The hydrogen then will take on a PARTIAL positive charge • The bond between the two is a POLAR COVALENT BOND…not a hydrogen bond

44. Hydrogen Bonding • If these molecules are in a solution and can come into contact with other molecules like themselves, the partial positive begin to attract the partial negatives and form a bond… • A hydrogen bond • This is what happens in water • Let’s draw it out

45. Hydrogen Bonding + + - - + + As you can see, a water molecule can be joined with up to 4 neighbors via hydrogen bonding + + - = hydrogen bond + - + + - +

46. Intermolecular Forces • Types: • Induced dipole • Dipole-dipole • Hydrogen bonds • So, why do we call these INTERmolecular forces? • These are forces or attractions BETWEEN molecules that pull them closer together • In order to overcome these attractions, more energy must be added

49. Binary Compounds • Many of the examples that we’ve just used were “binary compounds” • Made up of only 2 elements (no matter what the ratio) • Ternary compounds are made up of 3 different elements

50. Molecular Compounds • Not all binary compounds are ionic • Many are molecular (made up of 2 non-metallic elements) • Prefixes are often used to name molecular compounds because non-metals can bond in various ratios • For example, what ratio does Carbon bond to Oxygen? • CO or CO2