"Chapter 9 Chemical Bonding"
Chapter 9 Chemical Bonding Bonding of Atoms Bonding Extremes • In chapter 4 you learned that there were two types of bonds. • Ionic Bonds – Electrons are transferred from one atom to another forming ions. • Covalent (molecular) Bonds – Electrons are shared between atoms • This is simple, but not all bonds are one type or the other. Instead, there is a continuum from one type to the other. 12/9/2011 NAHS Chemistry D. Shepherd 2 Examples of the Extremes •Na + Cl [Na]+ + [ Cl ]- •This is an example of the ionic extreme. • C + O + O O C O • This is an example of the covalent extreme. 12/9/2011 NAHS Chemistry D. Shepherd 3 A tug of war between atoms • Picture two atoms, each one is represented by a person below. Both pull on the other’s valence electrons. Neither is strong enough to pull the electrons totally away, but one can pull a bit harder and the electrons move toward the harder pull. d- d+ e- e- The electrons are shared, but not equally, the resulting molecule has a somewhat negative end and a somewhat positive end. 12/9/2011 NAHS Chemistry D. Shepherd 4 Electronegativity • The relative attraction of an atom for the electrons in a bond is called electronegativity • Electronegativity is a periodic property and each element has been assigned a quantitative number as shown in table 9.2 on page 304. • The difference in the electronegativity values between any two elements will determine the bonding type that will occur between them. 12/9/2011 NAHS Chemistry D. Shepherd 5 Some Examples • Take a look at Hydrogen • Look at Hydrogen and and Chlorine Fluorine • The value for H is 2.1 • H is still 2.1 • The value for Cl is 3.0 • F is 4.0 • The electronegativity difference (DEN) is .9 DEN = 4-2.1 or 1.9 • Using figure 9.1 on page It is also polar, but to a 303, the bond type is read greater degree than is as polar covalent since .9 the HCl. is between 0.5 and 2.0 12/9/2011 NAHS Chemistry D. Shepherd 6 For you to try • Lithium and Bromine DEN = 1.8 is classified as • Cesium and Fluorine ionic, but has less ionic character than K and Cl • Fluorine and Fluorine DEN = 3.3 Ionic • Sodium and Chlorine DEN = 0 Pure Covalent • Potassium Chlorine DEN = 2.1 Ionic • Oxygen and Oxygen DEN = 2.2 Ionic • Silicon and Carbon DEN = 0 Pure Covalent • Phosphorus and DEN = .7 Polar Bromine DEN = .7 Polar 12/9/2011 NAHS Chemistry D. Shepherd 7 Periodic Trends in Electronegativity Electronegativity Increases 12/9/2011 NAHS Chemistry D. Shepherd 8 Variation in % Ionic Character 100% 50% Percentage Ionic Character to the bond 0 0.5 1.0 1.5 2.0 2.5 3.0 Electronegativity difference between atoms 12/9/2011 NAHS Chemistry D. Shepherd 9 From the Graph • Once an electronegativity difference of 1.7 is reached, the bond has better than 50% ionic character. • The bond might be labeled correctly as polar covalent, but for the most part the rules you learned about metals and nonmetals forming ionic compounds and two nonmetals forming covalent compounds remains a good rule of thumb for deciding which bonding type is present. 12/9/2011 NAHS Chemistry D. Shepherd 10 Polar Bonding • We’ve looked at the extremes which are a bond with a DEN so large that it is completely ionic or with a DEN that is zero where the bond is completely covalent. • Now we will look closer at the unequal sharing that occurs in polar bonds and what it means. 12/9/2011 NAHS Chemistry D. Shepherd 11 Unequal Sharing • Hydrogen and Chlorine are 2 nonmetals and by our earlier rule they will form a covalent bond. • Their DEN is .9 which puts them in the polar covalent range. • Since Cl is more electronegative than H, it will have a stronger attraction for the electrons and the Cl end of the molecule will have a somewhat negative charge while the H end will have a somewhat positive charge. 12/9/2011 NAHS Chemistry D. Shepherd 12 Hydrogen Chloride d+ d- The electron from H is pulled toward Cl the Cl. H 12/9/2011 NAHS Chemistry D. Shepherd 13 Water is a polar molecule • Water is H2O, but let’s look at each H-O bond separately. • The DEN for H-O is 1.4, which makes it covalent. • The Oxygen is most electronegative and the electrons will be pulled in that direction • Let’s look at a model like the one for HCl 12/9/2011 NAHS Chemistry D. Shepherd 14 The Hydrogen – Oxygen bond d- The electrons are O H d+ pulled harder by the oxygen than the hydrogen atom. 12/9/2011 NAHS Chemistry D. Shepherd 15 A look at the Lewis structure of Water Oxygen has 6 valence O electrons and Hydrogen has 1. d- O H d+ H H H H d+ O H 12/9/2011 NAHS Chemistry D. Shepherd 16 Water’s Polarity and it’s shape • Water is bent because of it’s shared and unshared pairs of electrons. These two pair of electrons are These two pair H shared. of electrons are O H unshared. The unshared pair repel (like charges repel) and the molecule is “bent”. The angle is 105 o. 12/9/2011 NAHS Chemistry D. Shepherd 17 Water Polarity and it’s shape 2 • We’ll talk more about water in a later chapter, but if the 4 pair of electrons surrounding water in the H2O molecule were all equal the bond angle would be that of a tetrahedron (109.5o) • The nonbonding electrons require more room and “squeeze” the bonding pair closer and thus reduce the bond angle to 105 o. 12/9/2011 NAHS Chemistry D. Shepherd 18 Coffee Filter Chromatography • The mini lab on page 312 will help to illustrate some of water’s unique properties. • It will also give you some insight into polar and non-polar substances and how they might be expected to behave. • Follow the directions closely and answer the analysis questions. 12/9/2011 NAHS Chemistry D. Shepherd 19 Lewis Structures and Molecular Shape • The following steps may be used to determine the correct Lewis Structure. • 1.) Decide which atoms are bonded • 2.) Count ALL valence electrons • 3.) Place 2 electrons in each bond. • 4.) Complete the octets of the atoms attached to the central atom by adding e- in pairs. • 5.) Place any remaining electrons on the central atom in pairs. • 6.) If the central atom does not have an octet, form double bonds. If necessary, form triple bonds. 12/9/2011 NAHS Chemistry D. Shepherd 20 Examples of Lewis Structures and Molecular Shape • SO3 - and ClO4 – Draw Lewis Structures • First SO3 – S 1 atom x 6 electrons = 6 O 3 atoms x 6 electrons = 18 Total = 24 e- There are no electrons left and S does not have 8 O S O electrons so one pair from one of the oxygen atoms needs to be pulled in. O 12/9/2011 NAHS Chemistry D. Shepherd 21 Sulfur trioxide’s shape The molecule will have one double bond between one Oxygen and the Sulfur and single bonds joining the other two Oxygen atoms to the sulfur. The shape of the O S O molecule will be somewhat triangular. O 12/9/2011 NAHS Chemistry D. Shepherd 22 Websites with 3D Models • Virtual Chemistry site • If you use this site at home, you’ll need to get a Chime Plugin The MDL Chime Site • You will be building gumdrop models and will need to consider the spacing and direction some of the parts of the molecules are directed. 12/9/2011 NAHS Chemistry D. Shepherd 23 Formal Charge and Lewis Structures • [Formal Charge] = [Number of e- in valence shell of the isolated atom] – [(Number of bonds to the atom) + (Number of unshared e-)] • Let’s look at the structures that are possible for Sulfuric acid. H2SO4. • There is a structure that obeys the octet rule, but it is not the structure that is found by experiment. 12/9/2011 NAHS Chemistry D. Shepherd 24