Chapter 6 Chemical Bonding Covalent molecules

1. Chapter 6Chemical Bonding Covalent molecules

2. PART 1:Chemical Bonds

3. Chemical Bonds • Mutual electrical attraction between the nuclei and valence electrons of different atoms that binds the atoms together • Atoms are usually less stable by themselves than combined with other atoms

4. Types of Bonding • Ionic Bonding • Chemical bonding resulting from the attractions of cations and anions • In pure ionic bonding, atoms give up electrons completely to other atoms, which accept them completely to form ions which attract one another

5. Types of Bonding • Covalent Bonding • Chemical bond that shares the electrons between two atoms • A pure covalent bond results in the electrons being “owned” equally by two atoms • Both nuclei have these electrons around them an equal amount of time • Example of purely covalent bond : F2 • Both atoms attract each others’ electrons equally (both have an electronegativity of 4.0, so the pull on the electrons is exactly the same from both atoms)

6. Classifying Bonds:Ionic or Covalent • Most compounds do not have purely covalent or ionic bonds • The bonds fall somewhere on a continuum • Look at the bond’s percentage of ionic character • Bonds with 0 to 50% ionic character are classified as covalent • Bonds with > 50% ionic character are classified as ionic

7. Ionic vs. Polar Covalent vs. Nonpolar covalent Bonds • Ionic Bonds • Electrons get transferred between atoms forming ions that attract one another • Polar Covalent Bond • Electrons shared between atoms unequally • Nonpolar Covalent Bond • Electrons shared between atoms equally

8. Part 2:Classifying bonds using electronegativity

9. Can use Electronegativity Difference as a Guide • We can predict the bond type between two atoms by calculating the absolute value of the electronegativity difference between them.

10. Ionic Bonding • If the absolute value of the electronegativity difference between two atoms is > 1.7, then the bond they form is probably ionic • Example #1: NaCl • Na – electronegativity of 0.9 • Cl – electronegativity of 3.0 • Electronegativity difference = | 0.9 – 3.0 | = 2.1 • An electron moves from the Na atom to the Cl atom causing Na to have a +1 charge and Cl to have a -1 charge • The Na + ion is then attracted to the Cl- ion forming a very strong bond

11. Polar Covalent Bonds • If the absolute value of the electronegativity difference between two atoms is >0.3 - 1.7, then the bond they form is probably polar covalent • Covalent because the electrons are shared between the atoms • Polar because the electrons are shared unequally between the atoms (spend more time around one nucleus) • Example: BF3 • Boron – electronegativity of 2.5 • Fluorine – electronegativity of 4.0 • Electronegativity difference of | 4.0-2.5|= 1.5 • Polar covalent - electrons spend more time around the F nucleus due to greater electronegativity • Electron cloud is distorted towards the Fluorine atoms (more electron density) resulting in a partial negative charge on the fluorine atoms and a partial positive charge on the boron atom.

12. Nonpolar covalent bonds • If the absolute value of the electronegativity difference between two atoms is 0.0- 0.3, then the bond they form is probably nonpolar covalent • Electrons are shared equally between the nuclei of the bonded atoms, resulting in a balanced distribution of electrical charge • Ionic character of 0 to 5% ionic character • Example: F2 • Electronegativity difference of |4.0 – 4.0| = 0.0

13. Part 3:Covalent Compounds

14. Covalent Compounds • Molecules • Neutral group of atoms held together by covalent bonds • Can exist on its own as an individual unit • Molecular formula • Shows the relative numbers of atoms of each kind in a chemical compound by using element symbols and subscripts to show the number of atoms • Examples: • H2O - water • O2 - oxygen • C12H22O11 - sucrose

15. Bond Energy • Energy required to break a chemical bond and form neutral isolated atoms • Units kj/mole • Energy required to break one mole of bonds

16. Octet Rule • Chemical compounds tend to form so that each atom by gaining, losing or sharing electrons has an octet of electrons in its highest occupied energy level

17. Exceptions to the Octet Rule • Hydrogen – needs total of 2 electrons • Beryllium – needs total of 4 electrons • Boron – needs 6 electrons • Aluminum – needs 6 electrons • Atoms of elements in groups 15, 16, 17, and 18 can sometimes form expanded octets • Can have more than 8 electrons surrounding the nucleus • Usually not N, O, or F

18. Lewis Structures AKA Electron Dot Notation • Symbols represent the nucleus and inner shell electrons • Dot pairs between the symbol represent pairs of shared electrons • Dots adjacent to symbol represent unshared pairs of electrons

19. Steps for Drawing Lewis Dot structures • Count up the number of valence electrons. • E.g. in CCl4 there is 1 C which has 4 valence electrons and 4 Cl atoms which each have 7 valence electrons • Therefore for C 1 x 4 = 4 and for Cl 4 x 7 = 28 • Adding 4 and 28 together gives you 32 electrons • Identify the central atom and put the other atoms around the central atom. • If C is present, put C in the center. • Otherwise, the least electronegative atom goes in the center. • H never goes in the center. • Halogens don’t usually go in the center unless they are bonded to other halogens

20. Steps for Drawing Lewis Dot Structures (cont’d) • Put a pair of electrons (dots) between the central atom and the outside atoms. • Start distributing the remaining electrons • Subtract those electrons placed in step 3 from the total in step 1 to find electrons remaining • Put the remaining electrons around the outside atoms first until all atoms either have an octet or are satisfied (Exceptions to the Octet Rule for the exceptions). • If there are electrons left over after satisfying the octet rule, put them as unshared pairs on the central atom.

21. Steps for Drawing Lewis Dot Structures (cont’d) • If there are too few electrons for all atoms (including the central atom) to have an octet (or be satisfied if it is an exception), then double and or triple bonds should be considered. • C, N, and O are most atoms likely to form double or triple bonds. 

22. Drawing Lewis Dot Structures for Polyatomic ions • To draw the Lewis dot structure for a polyatomic ion (a covalently bonded molecule that has a charge), you follow the same steps except the following: • When counting up the valence electrons in step 1, add the absolute value of the charge if the charge is negative to the total of the valence electrons. If the charge is positive, subtract the absolute value of the charge from the total valence electrons. • After completing the Lewis dot structure, put the entire structure in brackets with the charge on the right upper corner on the outside of the right bracket.

23. Multiple covalent bonds • Some elements, especially C, N, and O can share more than one electron pair. • Double bond – 2 pairs of electrons are shared • Triple bond – 3 pairs of electrons are shared

24. Multiple bonds energies and bond lengths • Double bonds have a greater bond energy and shorter bond lengths than single bonds • Triple bonds have a greater bond energy and shorter bond lengths than single bonds • Triple bonds stronger than double • Double bonds stronger than single

25. Resonance Structures • Bonding in molecules or ions that cannot be correctly represented by a single Lewis structure • Draw possible Lewis structures with double-headed arrow in between Bond between oxygens is a hybrid of a single and double bond

26. Part 4: Chemical Formulas, Naming, and Shapes

27. Chemical Formula • Chemical formula indicates the relative number of atoms of each kind in a chemical compound. • Molecular formula indicates the number of atoms of each element contained in a single molecule • E.g. C8H18 • Ionic compounds have formula units, not molecules (Remember, these ions are located in a lattice of positively and negatively charged ions – are not found as individual units) • One formula unit is used to designate the simplest ratio of the cations to anions • E.g. in a crystal of NaCl, the formula NaCl is used to represent one unit

28. Naming Binary Molecular Compounds (Chapter 7 p. 227-229) • Prefix system of naming – must learn prefixes

29. Naming Binary Molecular Compounds • Element with the smaller group number is given first in the name • If both elements are in same group, element whose period number is greater is given first. • If first element subscript is one, no prefix is given. • If subscript is greater than 1, then the prefix corresponding to the subscript is added to the beginning of the element name. • The second element is always given a prefix indicating the number of atoms and the element root + ide is used. • The –o or –a at the end of a prefix is usually dropped when the word following the prefix begins with another vowel i.e. monoxide instead of monooxide or pentoxide instead of pentaoxide

30. Binary Molecule Name Examples

31. VSEPR Theory • Theory used to determine the shape of molecules • Based on the premise that repulsion between the sets of valence electrons surrounding an atom causes the electron pairs to be oriented as far apart as possible. • Looks at the bonded pairs and lone pairs on the central atom only • Bonded pairs are set as far apart from one another as possible

32. Molecular Geometries • Some representative molecular geometries – need to learn molecular geometries located on VSEPR Summary sheet in Chapter 6 folder

33. How to read VSEPR Summary Sheet • Must determine molecule type from Lewis Dot structure – must draw this first • Let A = central atom (in example below, N) • Let B = outside atoms regardless of whether they are the same element or not (3 outside atoms ) • Let E = represent the # of lone pairs on the central atom (only) (one lone pair on N) • Let subscripts for B and E represent the number of outside atoms and number of lone pairs on the central atom • Example: NF3 is AB3E • N is central atom or A, • the 3 Fs are represented as B3 • and there is one lone pair on the central atom or E • Molecular geometry - from VSEPR Summary sheet is trigonal pyramidal

34. Bond Angles • Angles between the bonded pairs of atoms around the central atom • Bond angles between bonded pairs tend to be larger for those atoms without lone pairs on the central atom • Lone pairs cause increased repulsion and force bonded pairs closer together

35. Polar molecules • Electrons are unevenly distributed throughout the molecule • Because of this there are areas of partial positive charge and areas of partial negative charge • Forms dipoles • Equal but opposite charges separated by short distances More positive more negative Partial - Partial +

36. Factors Determining Polarity of Molecules • Polar Bonds present in molecule • Symmetry of molecule • Do dipoles when taken together, cancel out or form a dipole in one direction? • If cancel out - then nonpolar • If overall dipole is present – polar

37. Guidelines for Determining Polarity • If molecule type is AB2,AB3,AB4, AB5, AB6, AB2E3, or AB4E2, and the outside atoms are the same element, then the molecule is nonpolar • If molecule type is AB and A and B are the same element, then the molecule is nonpolar • All others are polar • E.g. CH4 is AB4 all of the B’s are the same (H), therefore the molecule is nonpolar • CHCl3 is AB4 but not all of the B’s are H, (1 H and 3 Cls), therefore the molecule is polar

38. Hybridization • When bonding occurs, orbitals from the bonding atoms become mixed and form new orbitals that are equivalent energies (i.e. hybrid orbitals) • Hybrid orbitals • Orbitals of equal energy produced by the combination of two or more orbitals on the same atom • # of orbitals produced = # of orbitals that have combined

39. Hybrid orbitals and Multiple Bonds • All single bonds are sigma bonds – σ • Atomic orbitals overlap and form hybrid orbitals • Double and Triple bonds have one sigma bond • In addition, double bonds also have one π bond • Pi bonds (π) are overlapping atomic p orbitals on adjacent atoms • Triple bonds also have two π bonds

40. Hybrid orbitals • Count double and triple bonds as one bonding pair since there is only one sigma bond

41. Part 5: Intermolecular Forces

42. Intermolecular Forces vs. Intramolecular Forces • Intermolecular forces • Between molecules • 3 types • Dipole-dipole interaction • Hydrogen Bonding • London Dispersion Forces • Intramolecular Forces • Within molecules • True bonding • Covalent bonding Weaker Stronger

43. 3 Types of Intermolecular Forces • Dipole-dipole interactions • Attraction between molecules that occurs between oppositely charged ends of partially charged molecules • Occurs between polar molecules

44. Types of Intermolecular Forces • Hydrogen bonding – special case of dipole-dipole interaction • NOT Bonding!!!!! – intermolecular force • Occurs when highly polarized H is attracted to the lone pair of a highly electronegative N, O, and F • Occurs with molecules having N, O, or F directly bonded to a H • Strongest intermolecular force

45. Types of Intermolecular Force • London Dispersion Forces • Weakest intermolecular force • Occurs due to random motion of electrons in molecules forming temporary dipoles (areas of partial charge) which attract each other (oppositely charged ends attract) • Due to temporary nature of the dipoles, attractions are fairly weak but when added together can be significant • Occurs in all molecules (polar and nonpolar)

46. Strength of Intermolecular Forces • Hydrogen Bonding > Dipole-Dipole Interactions> London Dispersion Forces • Strength of intermolecular forces determines whether molecules have • High surface tension • High boiling point • Capillary action

47. Determining the Intermolecular Forces associated with a Molecule Is the molecule polar or nonpolar If polar, is there An N,O, or F directly Bonded to an H in the structure? If nonpolar, London Dispersion Forces only If yes, then the molecule has hydrogen bonding and London Dispersion Forces If no, then the molecule has a dipole-dipole interaction and London Dispersion Forces

48. Properties of Substances associated with Intermolecular Forces • Surface Tension • Force that pulls adjacent parts of a liquid’s surface together, thereby decreasing surface area to the smallest possible size • Molecules at the surface do not have other like molecules on all sides of them and consequently they cohere more strongly to those directly associated with them on the surface forming a surface "film” or “skin”.

49. Surface Tension in Molecules • The greater the strength of the intermolecular forces between molecules in a substance the greater the surface tension. • Hexane and Water • Which will have the greatest surface tension? • Hexane is nonpolar and only has London dispersion forces whereas water is very polar and has hydrogen bonding as well as London Dispersion Forces • Water has higher surface tension

50. Surface tension and Temperature • The greater the temperature of the substance, the lower the surface tension. • The greater the temperature, the greater the average kinetic energy of the molecules, the further apart they are from one another and the lower the force of attraction between them.