Covalent Bonds

1. Covalent Bonds Chapter 9

2. A Model of Bonding • Atoms either transfer electrons and then form ionic compounds or they share electrons to form covalent compounds. • In both cases, the bond forms because of an increase in stability.

3. A Model of Bonding • Dividing compounds into two bonding types, ionic and covalent, is convenient. • If you know the type of bonds in a compound, you can predict many of its physical properties.

4. A Model of Bonding • You can also reason in the opposite direction. If you know the physical properties of an unknown compound, you can predict its bond type. • But predictions may not always be correct because there is no clear-cut division between ionic and covalent compounds.

5. A Model of Bonding • A compound may be partly covalent and partly ionic. • A more realistic view of bonding is to consider that all chemical bonds involve a sharing of electrons.

6. A Model of Bonding • Electrons may be shared equally, but they also may be shared only slightly—almost not at all. • The properties of any compound, particularly its physical properties, are related to how equally the electrons are shared.

7. A Model of Bonding • The bonding between atoms in compounds can be viewed as a range of electron sharing measured by electronegativity difference, ∆EN. • This range contains three main classes of bonds—ionic, polar covalent, and covalent.

8. A Model of Bonding

9. A Model of Bonding • A purely ionic bond results when the sharing is so unequal that it is best described as a complete transfer of electrons from one bonding atom to another. • A purely covalent bond results when electrons are shared equally. • Most compounds fall somewhere in between these two extremes; they have some ionic characteristics and some covalent characteristics.

10. Electronegativity: An Attraction for Electrons • Electronegativity is the measure of the ability of an atom in a bond to attract electrons.

11. Electronegativity: An Attraction for Electrons • How each atom fares in a tug-of-war for shared electrons is determined by comparing the electronegativities of the two bonded atoms.

12. Electronegativity: An Attraction for Electrons • Atoms are assigned electronegativity values. • Atoms with large electronegativity values, such as fluorine, attract shared valence electrons more strongly than atoms such as sodium that have small electronegativities.

13. Electronegativity: An Attraction for Electrons • Electronegativity is a periodic property.

14. Electronegativity: An Attraction for Electrons • With only a few exceptions, electronegativity values increase as you move from left to right in any period of the periodic table.

15. Electronegativity: An Attraction for Electrons • Within any group, electronegativity values decrease as you go down the group.

16. Electronegativity: An Attraction for Electrons • That means that the most electronegative elements are in the upper-right corner of the table.

17. Electronegativity: An Attraction for Electrons • Fluorine has the highest value of 4.0. • It also follows that the elements having the lowest electronegativities are in the lower-left corner. • The electronegativity of cesium is 0.7. • The noble gases are considered to have electronegativity values of zero and do not follow the periodic trends.

18. Electronegativity: An Attraction for Electrons • The decrease in electronegativity as you move down a column occurs because the number of energy levels increases, and the valence electrons are farther from the positively charged nucleus. The nucleus has less attraction for its own valence electrons so they are held less tightly. Also, electrons in the inner energy levels tend to block the attraction of the nucleus for the valence electrons. This is known as the shielding effect.

19. Electronegativity: An Attraction for Electrons • The shielding effect increases as you move down a column because the number of inner electrons increases. • For example, magnesium has the electron configuration 1s22s22p63s2 and an electronegativity of 1.2; calcium has the configuration 1s22s22p63s23p64s2 and an electronegativity of 1.0.

20. Electronegativity: An Attraction for Electrons • Although both have two valence electrons, calcium has eight more inner electrons than magnesium. • These electrons shield the outer electrons from the attraction of the nucleus.

21. Because electronegativity varies in a periodic way, you can make predictions about differences in electronegativity by looking at the distance between bonding atoms on the table. Electronegativity: An Attraction for Electrons

22. In general, the farther the bonding atoms are from each other on the periodic table, the greater their electro-negativity difference. Electronegativity: An Attraction for Electrons

23. The Ionic Extreme • The greater the difference between the electronegativities of the bonding atoms, the more unequally the electrons are shared. • The electronegativity difference between two bonding atoms is often represented by the symbol ∆EN, where EN is an abbreviation for ElectroNegativity and ∆ is the Greek letter delta meaning “difference.”

24. The Ionic Extreme • ∆EN is calculated by subtracting the smaller electronegativity from the larger, so ∆EN is always positive. • For example, ∆EN for cesium and fluorine is 4.0 − 0.7 = 3.3.

25. When the electronegativity difference in a bond is 2.0 or greater, the sharing of electrons is so unequal that you can assume that the electron on the less electronegative atom is transferred to the more electronegative atom. Highly Unequal Sharing

26. Highly Unequal Sharing • This electron transfer results in the formation of one positive ion and one negative ion. • The bond formed by the two oppositely charged ions is classified as a mostly ionic bond.

27. Highly Unequal Sharing • Many bonds are classified as ionic, but they have varying degrees of ionic character. • The greater the difference in the electronegativities of the two atoms, the more ionic the bond. • The greater the distance between the bonding atoms on the periodic table, the more ionic the bond between the atoms.

28. Highly Unequal Sharing • The electronegativity differences in lithium fluoride, sodium chloride, and potassium bromide show that they are best represented as ionic compounds.

29. Equal Sharing, Covalent Bonds • Two of the same atoms form a covalent bond, for example, when two fluorine atoms form the fluorine molecule, F2. • Here is the electron dot structure for F2.

30. Equal Sharing, Covalent Bonds • Because two atoms of the same element are forming the bond, the difference in electronegativities is zero. • In the fluorine molecule, a pair of valence electrons are shared equally.

31. Equal Sharing, Covalent Bonds • This type of bond is a pure covalent bond. • All other diatomic elements (Cl2, Br2, I2, O2, N2, and H2) have pure covalent bonds. • In all these molecules, the electrons are shared equally. • Most of the elemental diatomic molecules are gases at room temperature—Cl2, F2, O2, N2, and H2.

32. Unequal Sharing in Polar Covalent Bonds • When the electronegativity difference between bonding atoms is between 0.5 and 2.0, the electron sharing is not so unequal that a complete transfer of electrons takes place. • Instead, there is a partial transfer of the shared electrons to the more electronegative atom. • The less electro-negative atom still retains some attraction for the shared electrons.

33. Unequal Sharing in Polar Covalent Bonds • The bond that forms when electrons are shared unequally is called apolar covalentbond. • A polar covalent bond has a significant degree of ionic character.

34. Polar covalent bonds are called polar because the unequal electron sharing creates two poles across the bond. Unequal Sharing in Polar Covalent Bonds • Just as a car battery or a flashlight battery has separate positive and negative poles, so polar covalent bonds have poles.

35. Unequal Sharing in Polar Covalent Bonds • The negative pole is centered on the more electronegative atom in the bond. This atom has a share in an extra electron.

36. Unequal Sharing in Polar Covalent Bonds • The positive pole is centered on the less electronegative atom. This atom has lost a share in one of its electrons.

37. Unequal Sharing in Polar Covalent Bonds • Because there was not a complete transfer of an electron, the charges on the poles are not 1+ and 1−, but δ+ and δ−.

38. Unequal Sharing in Polar Covalent Bonds • These symbols, delta plus and delta minus, represent a partial positive charge and a partial negative charge.

39. This separation of charge, resulting in positively and negatively charged ends of the bond, gives the polar covalent bond a degree of ionic character. Unequal Sharing in Polar Covalent Bonds

40. Unequal Sharing in Polar Covalent Bonds • The ∆EN for the O — H bond is 1.4, so water molecules have polar bonds. • When atoms of hydrogen and oxygen bond by sharing electrons, the shared pair of electrons is attracted toward the more electronegative oxygen.

41. The Shapes of Molecules • Models can help you visualize the three-dimensional structures of molecules. • Consider the simplest molecule that exists—hydrogen, H2. • Two hydrogen atoms share a pair of electrons in a nonpolar covalent bond, as shown in the electron dot structure.

42. The Shapes of Molecules • A hydrogen molecule is linear. • You could model other diatomic molecules, such as oxygen, nitrogen, chlorine, iodine, fluorine, and even hydrogen chloride, HCl.

43. The Shapes of Molecules • The model always predicts the same geometry: linear.

44. Modeling Water • Remember that the dot diagram models the arrangement of valence electrons in the molecule in two dimensions. • The electron dot diagram for a water molecule shows that each hydrogen shares a pair of electrons with the oxygen.

45. The eight valence electrons are distributed in such a way that the oxygen atom has an octet of electrons and the stable electron configuration of the noble gas neon. Modeling Water

46. Modeling Water • Each hydrogen has two valence electrons and the stable configuration of helium. • Two pairs of valence electrons are involved in the bonding. These electrons are called bonding pairs. • The other two pairs of valence electrons are not involved in the bonding. These are called nonbonding pairs, or lone pairs.

47. Modeling Water • A space-filling model shows the electron clouds of each atom as spheres. • The clouds overlap when two atoms form a bond. • This model is a good representation of the water molecule. 

48. Modeling Carbon Dioxide • Model carbon dioxide, CO2, as you did water. • Begin by drawing the electron dot diagrams for the two atoms. • Carbon has four valence electrons, and each oxygen atom has six.

49. Modeling Carbon Dioxide • To obtain a stable octet of electrons, the carbon atom needs four more electrons, and each oxygen atom needs two more electrons. • Therefore, each oxygen atom must share two pairs of electrons with the carbon atom.

50. Modeling Carbon Dioxide • A bond formed by sharing two pairs of electrons between two atoms is called a double bond, as illustrated by the electron dot structure for carbon dioxide.