Bonding, Electronegativity , & Bond Shape

1. Bonding, Electronegativity, & Bond Shape

2. Forming Chemical Bonds • According to the Lewis model • an atom may lose or gain enough electrons to acquire a filled valence shell and become an ion. An ionic bond is the result of the force of attraction between a cation and an anion. • an atom may share electrons with one or more other atoms to acquire a filled valence shell. A covalentbond is the result of the force of attraction between two atoms that share one or more pairs of electrons.

3. Types of Chemical Bonding Typically: 1. Metal with nonmetal: electron transfer leads to ionic bonding 2. Nonmetal with nonmetal: electron sharing leads to covalent bonding 3. Metal with metal: electron pooling leads to metallic bonding

4. Figure 9.2 The three models of chemical bonding

6. When drawing Lewis Dot Structures, start with one dot on the top and work your way around. Kr

7. Lewis Dot Diagrams Because valence electrons are so important to the behavior of an atom, it is useful to represent them with symbols. A Lewis Dot Diagram shows dots for each electron in the atom. Each dot represents one valence electron. Li Be B C N O F Ne

8. Ionic Compounds Na  Na+ + e- Cl + e-  [ Cl ]- Na+ + [ Cl ]-  Na+[ Cl ]- Now chlorine and sodium have 8 valence electrons, fulfilling the octet rule. Ionic compounds form crystals: latticed structures of atoms lined up in predictable ways. Remember, negative and positive charges attract, so Na+ and Cl- stick together, making an ionic compound.

9. Ionic bonding: Al + Cl Al [Al]3+ [ Cl ]3– Cl Cl Cl Al + 3Cl [Al]3+[Cl]3–

10. Covalent Bonds in NH3 Bonding pairs H   H: N : H   Lone pair of electrons (unshared pair)

11. Covalent bonding Cl Cl C Cl Cl H H H H H H H Cl N H O [Mg]2+ [Na] + [ F ]2– [ Cl ]– H O Q7 CCl4 - Covalent HCl - Covalent MgF2 - Ionic NH3 - Covalent NaCl - Ionic H2O - Covalent H2 - Covalent OH– - Covalent

12. Covalent Bonds Two nonmetal atoms form a covalent bond because they have less energy after they bonded H+H H: H = HH = H2 hydrogen molecule

13. Double Covalent Bond 2 pairs of electrons are shared between 2 atoms Example O2       O  + O  O::O      double bond

14. Triple Covalent Bond 3 pairs of electrons are shared between 2 atoms Example N2       N  + N  N:::N  triple bond

15. Drawing Lewis Structures 1. Determine the number of valence electrons in the molecule 2. Decide on the arrangement of atoms in the molecule 3. Connect the atoms by single bonds 4. Show bonding electrons as a single line; show nonbonding electrons as a pair of Lewis dots 5. In a single bond, atoms share one pair of electrons; in a double bond, they share two pairs, and in a triple bond they share three pairs.

16. Lewis Structures

17. Electronegativity Electronegativity is a measure of the attraction of an atom for the electrons in a chemical bond. The higher the electronegativity of an atom, the greater its attraction for bonding electrons. Electronegativity is related to ionization energy.

18. Electrons with low ionization energies have low electronegativities because their nuclei do not exert a strong attractive force on electrons. Elements with high ionization energies have high electronegativities due to the strong pull exerted on electrons by the nucleus.

19. In a group, the electronegativity decreases as atomic number increases, as a result of increased distance between the valence electron and nucleus (greater atomic radius). An example of an electropositive (i.e., low electronegativity) element is cesium; an example of a highly electronegative element is fluorine.

21. When two identical atoms form a covalent bond, as in H2 or Cl2, each have equal share of the electron pair. the electron density is the same on both ends of the bond, because the electrons are equally attracted to both nuclei.

22. But when different atoms bond, as in HCl, one nucleus attracts the electrons from a bond more strongly than the other. The electrons are shared, but unequally. There are partial charges on both sides of the bond, indicated by a lowercase Greek letter delta, . In HCl, the chlorine carries a partially negative charge (-) while the hydrogen carries a partially positive charge (+). This is because the chlorine atom attracts the electrons stronger than the hydrogen.

24. Covalent or Ionic ??? • To decide whether a bond is covalent or ionic find the difference in electronegativities < 2.0 covalent > 2.0 ionic Try KF, MgS, Cl2

25. 3.0 2.0 DEN 0.0 Figure 9.18 Boundary ranges for classifying ionic character of chemical bonds.

26. Figure 9.19 Percent ionic character of electronegativity difference (DEN).

27. Polarity A bond that carries a partial charge is called a polar covalent bond. Because there are two poles of charge involve, the bond is a dipole. If the entire molecule has a partial charge, it is a polar molecule. This relative attraction of electrons in a bond is called electronegativity. In HCl, the Cl is more electronegative than the H and attracts the electrons. The electrons spend more time around the Cl than the H. Looking it up on a periodic table or a similar table, one can see the H has an electronegativity of 2.1 while Cl has one of 2.9. Therefore, the electronegativity difference is 0.8. Electronegativity decreases as you go down a group and increases as you go left to right in a period.

28. Because the bonding pair in the carbo-fluorine bond is pulled towards the fluorine end of the bond, that end is left rather more negative than it would otherwise be. The carbon end is left rather short of electrons and so becomes slightly positive. The symbols + and  - mean "slightly positive" and "slightly negative". You read  + as "delta plus" or "delta positive". We describe a bond having one end slightly positive and the other end slightly negative as being polar.

31. The polar nature of the elements sets up an overall charge on the molecule. This overall charge is represented by the symbol and called a dipole moment. Each dipole moment for each bond faces the - molecule. Then all the dipole moments are “added” and face the general direction of all the other dipole moments.

32. These 2 molecules have individual dipole moments for each bond, but they face opposite directions, so they cancel each other out for a dipole moment of 0.

34. 3-D Characteristics of Molecules • Atoms and molecules have 3 dimensions • Shapes of molecules lead to additional properties of covalent compounds • Polar covalent Bonding • When electrons are not shared equally between two atoms • Bond that is certain % ionic • Nonpolar covalent Bonding • Electrons are shared equally • Diatomic atoms

35. Shapes of Molecules Number of electron pairs 2 (= negative charge clouds) Number of bonded atoms 2 Angle 180° Name of shape linear

36. Here are some basic compounds that you see everyday. Do you know what they are? As you can see, even though they are made of the same elements (carbon and hydrogen), they have very different shapes. This is due to the way the atoms bond to each other.

37. In this chart you can see the 6 basic shapes of molecules. Most molecules in the world take one or more of these shapes. This chart is nice because it shows the number of valence electrons involved in bonding.

38. Here are a few of the odd molecule shapes you may encounter. The figure shows the non-bonding pairs of electrons. Remember on a Lewis Dot Structure, you may have only a single dot or a double dot. The double dots do not bind to anything. Those double dots are non-bonding pairs. Non-bonding electrons Na O Bonding electrons

39. As seen in the figure to the left, each shape has a particular bond angle. This angle is determined by how much stress in placed on the bonds. Electrons do not like to be near each other and will push apart as much as possible. Bond angles do not change. Look at the trigonal planar molecule. The electrons can only push so far before the electrons on the other side begin to push back. These forces keep the bonds in the same place.

40. VSEPR Theory

41. Using the VSEPR Model 1. Draw the electron-dot structure • Identify the central atom • Count the total number of electron pairs around central atom • Predict the electron shape • Predict the shape of the molecule using the bonding atoms

42. Valence Shell Electron-Pair Repulsion Theory or VSEPR • molecular shape is determined by the repulsions of electron pairs • Electron pairs around the central atom stay as far apart as possible. • electron pair geometry - based on number of regions of electron density • Consider non-bonding (lone pairs) as well as bonding electrons. • Electron pairs in single, double and triple bonds are treated as single electron clouds. • molecular geometry - based on the electron pair geometry, this is the shape of the molecule

43. Figure 10.4 A balloon analogy for the mutual repulsion of electron groups.

44. VSEPR - electrons (e-) in bonds/orbitals repel each other  4 groups/bonds repulsion yields a tetrahedral shape ~109.5o  3 bonds repulsion yields a trigonal shape ~120o  2 bonds repulsion yields a linear shape 180o eg: CH4 or NH3, HCO2H, CO2

45. The steps in determining a molecular shape. Molecular formula Step 1 Lewis structure Count all e- groups around central atom (A) Step 2 Electron-group arrangement Note lone pairs and double bonds Step 3 Count bonding and nonbonding e- groups separately. Bond angles Step 4 Molecular shape (AXmEn)

46. PROBLEM: Draw the molecular shape and predict the bond angles (relative to the ideal bond angles) of (a) PF3 and (b)COCl2. SOLUTION: (a) For PF3 - there are 26 valence electrons, 1 nonbonding pair SAMPLE PROBLEM 10.7 Predicting Molecular Shapes with Two, Three, or Four Electron Groups The shape is based upon the tetrahedral arrangement. The F-P-F bond angles should be <109.50 due to the repulsion of the nonbonding electron pair. The final shape is trigonal pyramidal. <109.50 The type of shape is AX3E

47. 124.50 1110 SAMPLE PROBLEM 10.7 Predicting Molecular Shapes with Two, Three, or Four Electron Groups continued (b) For COCl2, C has the lowest EN and will be the center atom. There are 24 valence e-, 3 atoms attached to the center atom. C does not have an octet; a pair of nonbonding electrons will move in from the O to make a double bond. Type AX3 The shape for an atom with three atom attachments and no nonbonding pairs on the central atom is trigonal planar. The Cl-C-Cl bond angle will be less than 1200 due to the electron density of the C=O.

48. Figure 9.12 The attractive and repulsive forces in covalent bonding.

49. Bond Energy • Bond energy or Bond Strength - the energy required to overcome the attraction of covalently bonded atoms. • It is defined as energy required to break bonds in 1 mole of gaseous atoms. • Bond energy depends on the specific elements involved. • It can vary from molecule to molecule so table values are averages.