Bonding and Forces of Attraction

1. Bonding and Forces of Attraction (mostly Chapter 9)

2. Metallic Bonding • In metallic bonding instead of sharing electrons between two atoms (covalent bonding), the valence electrons are shared between all the atoms in a positive metal crystal. These atoms are attracted to the negatively charged 'cloud’ or ‘sea' of electrons. • The movement or delocalization of the free electrons means that metallic bonded materials have good thermal and electrical conduction. http://www.ider.herts.ac.uk/school/courseware/materials/bonding.html

3. Ionic Bonds Ionic Bonds Occur when the nonmetal “steals” one or more electrons away from a metal. The nonmetal becomes a negative ion while the “victim” metal becomes a positive ion. The atoms are held together by their opposite charges.

4. Ionic Bonds Strength of crystal lattice depends on two factors, size and charge transferred. Smaller atoms have stronger ionic bonds. Ex: NaF is stronger than NaCl Atoms transferring more electrons are stronger. Ex: MgCl2 is stronger than NaCl.

5. Covalent Bonding The bonds of nature.

6. Covalent Bond • Shared valence electrons • Complete outer energy levels • Molecule has 2 or more nonmetal atoms covalently bond • Carbohydrates, proteins, fats, DNA, stupendous seven (H2, N2, O2, F2, Cl2, Br2, I2)

7. How do they form? Distance is to great E - E – repulsive P – E - attractive P – P – repulsive E – E - repulsive Repul = Attract Attractive forces balance the repulsive forces

8. Electronegativity • Electronegativity is the ability of atoms in a molecule to attract electrons to themselves. • On the periodic chart, electronegativity increases as you go… • …from left to right across a row. • …from the bottom to the top of a column.

9. What types of bonds are they? MgO, water, Calcium Carbide, Potassium Oxide, Nitrogen trihydride

10. Electronegativity and Bond Type • Ability to attract electrons in a chemical bond—subtract for bond type. • Ionic Bonds 1.7 or greater • Polar Covalent Bonds < 1.7 and >0.2 • Pure or Nonpolar Covalent bonds <0.2

11. Dash Line Represents the Cut Off

12. Polar Covalent Bonds • Electrons are shared, but not equally • Non polar bonds – equal sharing, H2, O2, Cl2

13. 1 valence e 1 s1 1 valence e 1 s1 1 s2 Stable Helium configuration

14. Single Covalent bond – shared electrons (also known as a sigma bond σ) 1s1 1s1 1s2 2s2 2p4 H –1s2 - He O - 1s2 2s2 2p6 - Ne Lone pairs – unshared elecrons The molecules are more stable because they have complete outer energy levels – Octet Rule

15. Bonding Capacity

16. Drawing Lewis Dot Structures • Count the valence electrons. • Predict the location of the atoms • Hydrogen is a terminal atom • The central atom has the smallest electronegativity. • Draw a pair of electrons between the central atom and the surrounding atoms. • Use the remaining electrons to complete the octets of each atom. If there are electrons left over, place them on the central atom. • If the central atom does not have a complete octet then try double or triple bonds. a. If the atom has 1, 2, or 3 valence electrons, it doesn’t require an octet.

17. STEP 1: count the total # of valence e- for all atoms involved in the bonding Carbon: 1 carbon with 4 valence electrons (1x4) = 4 CCl4 4+28 =32 Chlorine: 4 chlorine with 7 valence electrons (4x7) = 28

18. STEP 2–place the single atom in the center and other atoms around it evenly spaced Cl CCl4 4+28 =32 Cl Cl C Cl

19. STEP 3: place the electrons in pairs between the central atom and each non-central atom CCl4 4+28 =32 -8 =24 Cl Cl C Cl Cl

20. STEP 4: place the remaining electrons around the non- central atom until each has 8 electrons (H atoms have) only 2e- CCl4 4+28 =32 -8 =24 -24 =0 Cl C Cl Cl Cl

21. Step 5: If you run out of electrons before the central atom has an octet…form multiple bonds until it does. Ex—HCN Hydrogen—1 electron Carbon—4 electrons Nitrogen—5 electrons TOTAL—1+4+5=10 e-s H:C:N .. H:C:N: H:C:::N: ..

22. Drawing Lewis Dot Structures Draw Lewis Dot Structures for: PH3 H2S HCl CCl4 SiH4 CH2Cl2

23. Draw Lewis Dot Structures Cl2 NF3 CS2 BH3 CH4 SCl2 C2H6 BF3

24. Strength of Bonds • Based on proximity – bond length • Influenced by atom size and number of shared electrons • F2 is greater than Cl2 is greater than Br2 F2 1.43 x 10-10 m - single O2 1.21 x 10-10 m - double N2 1.10 x 10-10 m – triple

25. Sigma Bonds - σ • Overlapping electron orbitals create a bonding orbital – likely hood of finding an electron • S overlaps with S • S overlaps with P • P overlaps with P

26. Pi Bonds--π Pi bonds happen after sigma bonds. While sigma bonds are head to head bonds, pi bonds are side to side bonds of the same two atoms. Every bond has one sigma bond, but subsequent bonds between the same two atoms must have a different way of connecting so they use pi bonds!

27. Multiple Covalent Bonds – Double and Triple Bonds 6 valence electrons 6 valence electrons 12 valence electrons Octet satisfied More stable and stronger 1 sigma bond 1 pi bond

28. 5 valence electrons 5 valence electrons 10 valence electrons Octet satisfied More stable and stronger 1 sigma bond 2 pi bonds

29. Molecular Shapes • The shape of a molecule plays an important role in its reactivity. • By noting the number of bonding and nonbonding electron pairs we can easily predict the shape of the molecule.

30. What Determines the Shape of a Molecule? • Simply put, electron pairs, whether they be bonding or nonbonding, repel each other. • By assuming the electron pairs are placed as far as possible from each other, we can predict the shape of the molecule.

31. Valence Shell Electron Pair Repulsion Theory (VSEPR) “The best arrangement of a given number of electron domains is the one that minimizes the repulsions among them.”

32. Molecular Shape Chart

33. Polarity

34. Molecular Polarity • Symmetrical molecules have equal charge distribution so are nonpolar like tetrahedral CCl4 with 109.5o between chlorine atoms. • Unshared electron pairs in molecules push the shared pairs closer together, so NH3 (trigonal pyramid) has 107o between hydrogens, and water, H2O (bent), has 104.5o between hydrogen atoms.

35. Hybridization • The mixing of atomic orbitals in an atom to generate a new set of bonding orbitals (hybrids) • Different shapes than atomic orbitals • Requires energy but the energy is returned during bond formation 2sp3

36. Hybrid Orbitals • Consider beryllium: • In its ground electronic state, it would not be able to form bonds because it has no singly-occupied orbitals.

37. Hybrid Orbitals But if it absorbs the small amount of energy needed to promote an electron from the 2s to the 2p orbital, it can form two bonds. This new orbital set is hybridized with equal energy—2sp orbitals 1s 2sp orbitals

38. Group 2A elements make sp hybridizedorbitals Group 3A elements make sp2 hybridized orbitals

39. Elements in group 4A have 4 valence electrons and need four bonds to make an octet, therefore, they will have sp3 hybridization.

40. Questions • Why do some solids dissolve in water but others do not? • Why are some substances gases at room temperature, but others are liquid or solid? • What gives metals the ability to conduct electricity, what makes non-metals brittle? • The answers have to do with … Intermolecular forces

41. Intermolecular forces There are 2 types of attraction in molecules: intramolecular bonds & intermolecular forces • intramolecular bonds attraction between atoms in a molecule (covalent, ionic, and metallic bonds) • Intermolecular forces (IMF) the attraction between molecules (hydrogen bonds, dipole-dipole interactions, and London dispersion forces • These are listed from strongest to weakest in each category.

42. Intermolecular Forces Collectively, all intermolecular forces of attraction are known as van der Waals forces. They are not chemical bonds, but simply attractions between two covalent molecules. For different substances of similar molar mass, the weakest of these are London dispersion forces. Carbon dioxide, O=C=O, is a good example.

43. + – H Cl + – + – + – + – Dipole - Dipole attractions • We have seen that polar molecules have a separation of charge. • This happens in both ionic and polar bonds (the greater the EN, the greater the dipoles) • Molecules are attracted to each other in a compound by these +δ and -δ opposite charges

44. Polarity in Bonds vs. in Molecules • NH3 and H2O are both polar molecules with polar bonds. N-H: 3.04 - 2.20 = 0.84 O-H: 3.44 – 2.20 = 1.24 Electronegativity difference is between 0.2 and 1.7 so bonds in molecule are polar Unshared electron pairs in trigonal pyramid and bent molecules make both molecules polar • CH4 and BeH2 are both nonpolar molecules with polar bonds. C-H: 2.55 – 2.20 = 0.35 H-Be: 2.20 – 1.57 = 0.63 Again, electronegativity difference is between 0.2 and 1.7 so bonds in molecule are polar Symmetry of the tetrahedral and 3-atom linear molecules make both molecules nonpolar

45. London forces (temporary dipoles) Since non-polar molecules do not have dipoles, how can they attract each other to form solids or liquids? • London forces! (named after Fritz London) • London forces exist in non-polar molecules • Since electrons move in atoms, there are instants when the electrons shift to one side, causing that side to be slightly negative. This induces the surrounding molecules to do the same. • This results in the opposite side becoming slightly positive.

46. London forces Induced dipole: Instantaneous dipole: Eventually electrons are situated so that tiny dipoles form A dipole forms in one atom or molecule, inducing a dipole in the other

47. Polar molecules already have a separation of charge, due to either unshared electron pairs, or a change of one of the atoms around the central atom. Nonpolar molecules have even charge distribution because they are symmetrical.

48. Hydrogen bonding Hydrogen bonding is a special type of dipole - dipole attraction that is very strong • It occurs only when hydrogen is bonded to either N, O, or F • The high electronegativity difference of N-H, O-H, and H-F bonds cause these to be about 5x stronger than normal dipole-dipole forces • Compounds containing these bonds are important in biological systems (hold strands of DNA together; cause shape of enzymes) • An exceptionally important example is water!

49. In conclusion: • If two molecules have roughly the same molar mass, determine which has the strongest attractions by drawing the electron dot structure. • Those with the highest boiling point, are attracted to each other the strongest, so for boiling point: hydrogen bonded > dipole-dipole > nonpolar

50. To Help You Review! • Read through the questions that follow to review this unit. These questions will help you on your test. The answers follow!