Bonding & Molecular Shape

1. Bonding & Molecular Shape

2. Alkali Metals Alkaline Earth Metals Transition Metals Boron Group Carbon Group Nitrogen Group Oxygen Group (Chalcogens) Halogens Noble Gases Actinide Series Lanthanide Series Groups on the Periodic Table

3. Why Are There Patterns? • Elements have physical and chemical properties based upon their valence electrons. • These valence electrons are the electrons in the outermost energy level (s & p orbitals) • The number of valence electrons may be determined by using the periodic table

4. Be like the Noble Gases • Bonding is when atoms tend to gain, lose, or share electrons in order to gain a full set of valence electrons. • Metals will usually lose electrons which is why they form positive ions. Nonmetals will lose electrons and usually form negative ions.

5. Take a look at the number of electrons in the outer shell of each of the atoms below. Notice Sodium (Na) has 1 in its outer shell. It is easier to remove 1 than gain 7, so notice that the electron is leaving Sodium and heading to the Chlorine atom.

7. Lewis Dot Structures • Dots representing valence electrons are written outside of the chemical symbol. • Lewis dot structures are a quick and easy way to see how many valence electrons an atom has. • The Group Number on the Periodic Chart is the same as the valence electrons.

8. To the left is the electron configuration of the second period of elements. Notice that there are never more than 2 electrons per side.

10. Chemical Bonding • Any chemical reaction is a rearrangement of valence electrons to get a filled outer energy level. Chemical bonds are formed. • HOWEVER, while there are different types of bonds, they are all very similar.

11. The Nature of Bonding • All bonds are comprised of electrons (always in pairs) between the two atoms that are being bonded. • The are classified differently based on how even or uneven the sharing of electrons is. • The measure of how much an atom pulls on shared electrons is called Electronegativity.

12. Electronegativity • This idea was developed by Linus Pauling – 2 time Nobel Prize winner. • It is a measure of how strongly an atom grabs onto shared electrons. • Every atom has its own value. • Fluorine is the highest at 4.0 • We can use these values to classify which type of bond we are working with. • The symbol for it is En

13. Electronegativity # 2 • Here is how the type of bond is determined. • No matter how big the atom, it can be broken down to individual pairs of atoms being bonded together. • Look up the En for each atom. • Subtract the values so it comes out to be a positive number or zero. • This value is the key to learning more about the bond in question.

14. Ionic Bonds • In an ionic bond electrons are ‘transferred’ from one element to another. • The cation (+ ion) is attracted to the anion (- ion); a bond is formed. • Ions in an ionic bonds may be positive or negative.

16. Ionic Bonds Continued • Keep in mind that the electron can’t be totally transferred. There must always be a connection between both atoms and the shared electrons. If not, then they are just floating ions, like they are when in a solution.

18. Covalent Bonds • Sometimes electrons are shared more evenly between molecules. • While the atom that is more electronegative is winning the “tug of war for the shared electrons”, the difference is small and the electrons show balance between the two atoms.

20. Bond Lengths • Not all bond lengths are the same. • H-H bond is the shortest at (0.074 nm) • As you move down a group the bond length becomes longer. • Multiple bonds are shorter than single bonds • C-C  0.154 nm • C=C  0.134 nm

21. Bond Length

24. Bonding & Electronegativity • The electronegativity difference (ΔEn) between two atoms determines type of bond. • Non-polar covalent is when the ΔEn is ≤ .4 • Polar covalent is when the .4 < ΔEn < 1.67 • Ionic is when the ΔEnis ≥ 1.67

26. Practice • What kind of bond is formed between • Fluorine and Cesium • Potassium and Iodine • Phosphorus and Sulfur

27. Formulas • Empirical Formula – The ratio of elements in a compound (CH2O) • Molecular Formula – gives the actual number of atoms in a compound C6H12O6

28. Formulas • Models – Show which atoms are bonded to which. • Molecular Shape – Gives the angles of the bond

29. Formulas • Ball and Stick – Ball represents the inner atom, the sticks represent bonds • Space Filling – shows a more life-like view of the molecule

30. How Do We Know This? • VSEPR Theory – Valence-shell electron pair repulsion theory • In a small molecule, pairs of valence electrons are arranged as far apart from each other as possible. • Lone pairs of electrons repel more than electrons in bonds • Double bonds count as one pair!!

31. Effective Pairs • Covalent bonds count as 1 effective pair. • Double bonds and triple bonds count as 1 pair !! • Lone pairs count as 1 pair.

32. Shapes of Small Molecules • Linear – atoms connected form a straight line. • Bond forms an angle of 180o • Carbon dioxide and acetylene are examples.

33. Shapes of Small Molecules • Trigonal Planar – 3 atoms connected to a central atom. • Bond forms an angle of 120o • Boron trichloride & carbonate ion

34. Shapes of Small Molecules • Tetrahedral – 4 atoms connected to a central atom. • Bond forms an angle of 109.5o • Methane & carbon tetrachloride

35. Shapes of Small Molecules • Pyramidal – 3 sided pyramid • Lone electron pair not involved in bonding. • Bond forms an angle of 107.3o • Ammonia, Phosphorous Trichloride

36. Shapes of Small Molecules • Bent – 3 atoms not in a straight line • There are lone electron pairs not involved in bonding. • Bond forms an angle of 104.5o • Water & Nitrite Ion

37. Beyond Tetrahedral • Bromine Pentafluoride (square pyramidal BrF5) • Bromine Trifluoride (T-shaped BrF3)

38. Beyond Tetrahedral • Phosphorus Pentachloride (trigonalbipyramidal PCl5) • Sulfur Tetrafluoride (see-saw SF4)

39. Beyond Tetrahedral • Sulfur Hexafluoride (octahedral SF6) • Tetrachloroiodate Ion (square planar ICl4-1)

40. Practice • What shape is formed from • Nitrogen trichloride • Formaldehyde • Oxygen Difluoride • There is a tutorial on Shapes of Molecules on the website.

41. Hybrid Orbitals • In bonding the electron orbitals “mix” together forming hybrids • The hybrid orbital had characteristics of both. • Many types of hybrid orbitals are possible. • The number of hybrid orbitals created must always equal the number of atomic orbitals that are started with.

42. sp Hybrid Orbitals • sp hybrid orbitals combine a 1s and 1p orbital. • Forms linear molecules like BeF2.

43. sp2 Hybrid Orbitals • sp2 hybrid orbitals combine 1s and 2p orbitals. • Forms trigonal planar molecules like BF3.

44. sp3 Hybrid Orbitals • sp3 hybrid orbitals combine 1s and 3p orbitals. • Forms tetrahedral molecules like CH4. • Other hybrid orbitals form from the d orbital.

45. Polarity • Molecules as well as bonds may be polar or nonpolar. • Polar molecules are also called dipoles. • Polarity is determined by the shape of the molecule as well as the polarity of each of the bonds.

46. Polarity • Bond polarity is represented by vectors. • Arrows point to the more electronegative atom. • The polarity of the molecule is the sum of all the vectors.

48. Practice • Is it a polar molecule? • Chloroform (CHCl3) • Formaldehyde • Methane

49. Hydrogen Bonds • The weak intermolecular forces between polar molecules creates Hydrogen bonds. • H bonds have 5% of the strength of a covalent bond. • This might not seem like much but there are so many of them that they require a lot of energy to break. • This gives water unique properties like its high boiling pt.

50. Polarity Determines Physical Properties • Polar molecules will align in an electrical field. • The more polar a molecule, then the higher the boiling point. • Higher polarity gives a high surface tension.