1. Bonding Theories ABonding is the way atoms attach to make molecules. BBonding theories allow us to: predict the shapes of molecules and to predict the properties of substances based on the type of bonding within the molecules design and build molecules with particular sets of chemical and physical properties C.One way to represent the bonding in substances is to draw Lewis Diagrams: known as electron dot symbols uses symbols of the elements to represent the nucleus and inner electrons in atoms uses dots around the symbol to represent the valence electrons put one electron on each side first, then pair the electrons remember that elements in the same group have the same number of valence electrons; therefore, their Lewis dot symbols will look alike

2. Types of bonds Ionic—complete transfer of 1 or more electrons from one atom to anotherIonic—complete transfer of 1 or more electrons from one atom to another Covalent—some valence electrons shared between atomsCovalent—some valence electrons shared between atoms Most bonds are somewhere in between.Most bonds are somewhere in between.

3. Electronegativity

4. Electronegativity & Bond Polarity If the difference in electronegativity between bonded atoms is 0, the bond is pure covalent. –equal sharing If the difference in electronegativity between bonded atoms is 0.1 to 0.3, the bond is non-polar covalent. If the difference in electronegativity between bonded atoms is 0.4 to 1.9, the bond is polar covalent. If the difference in electronegativity between bonded atoms is larger than or equal to 2.0, the bond is ionic. A dipole is a material with positively and negatively charged ends. Polar bonds or molecules have one end slightly positive,  + ; and the other slightly negative,  - –They are not “full” charges and come from nonsymmetrical electron distribution Dipole Moment, , is a measure of the size of the polarity –measured in Debyes, D

5. Polarity of Molecules In order for a molecule to be polar it must 1)have polar bonds electronegativity difference bond dipole moments - measured 2)have an unsymmetrical shape Polarity effects the intermolecular forces of attraction polar bonds, and unsymmetrical shape causes molecule to be polar polar bonds, but non-polar molecule because pulls cancel

6. Several notes Polar bonds don’t necessarily make polar molecules. We need to add in the molecule shape. Metals have EN less than 2 Nonmetals have EN greater than 2 F has the largest EN, remember it has the smallest radius (except for noble gases) Cs has the smallest EN, it’s the largest radius

7. More Notes HF bond is very “polar” as 4.0-2.1 = 1.9 This is the “covalent” bond with the largest delta EN HF bond is more polar than a HO, than a HN, than an HC bond NaI = 2.5-.9 = 1.6 This delta EN is smaller than most ionic compounds. NaF is more ionic than NaI Keep the rule: Metal + nonmetal = ionic Nonmetal + nonmetal = covalent.

8. Stable Electron Configuration And Ion Charge Metals form cations by losing enough electrons to get the same electron configuration as the previous noble gas. Nonmetals form anions by gaining enough electrons to get the same electron configuration as the next noble gas.

9. Electron Configurations and Bonding When a nonmetal and a Group 1,2, or 3 metal react to form a binary ionic compound, the ions form in such a way that the valence-electron configuration of the nonmetal is completed to achieve the configuration of the next noble gas, and the valence orbitals of the metal are emptied to achieve the configuration of the previous noble gas. In this way both ions achieve noble gas electron configuration.

10. Electron Configurations and Bonding When two nonmetals react to form a covalent bond, they share electrons in a way that completes the valence-electron configurations of both atoms. That is, both nonmetals attain noble gas electron configurations by sharing electrons.

11. Ionic Bonds Na + Cl - Na + Cl - Na + Essentially all ionic compounds are solids These solids have a definite geometric structure called a crystal These strong electrostatic forces that hold these ions together in a fixed position are called ionic bonds

12. Salt versus Sugar Salt Dissolves in water Conducts electricity Melts at 801 ºC Sugar Dissolves in water Does not conduct electricity Chars at 160 ºC (smells like caramel) What’s the difference between these structures? Salt has bonds that involve the gaining or losing of electrons Sugar has bonds that involve the sharing of electrons

13. Lewis Formulas of Molecules Lewis diagrams show patterns of valence electron distribution within the molecule. They are useful for understanding the bonding in many compounds. They allow us to predict shapes of molecules. They allow us to predict properties of molecules and how they will interact with each other. Some of the common bonding patterns: –Carbon, C, has 4 bonds & 0 lone pairs 4 bonds: 4 single, or 2 double, or single + triple, or 2 single + double –N = 3 bonds & 1 lone pair, –O = 2 bonds & 2 lone pairs, –H and halogen = 1 bond; H cannot form an octet –Be = 2 bonds & 0 lone pairs; it cannot form an octet –B = 3 bonds & 0 lone pairs; it cannot form an octet. B C NOF

14. Lewis Symbols for Ions Cations have Lewis symbols without valence electrons. –Lost in the cation formation Anions have Lewis symbols with 8 valence electrons, the Octet Rule. –Electrons gained in the formation of the anion Li Li +1 Metal to nonmetal bonding. The metal loses electrons to form a cation The nonmetal gains electrons to form an anion. An ionic bond results when a positive ion, +, is attracted to a negative ion, -. –The larger the charge, the stronger the attraction. –The smaller the ion, the stronger the attraction. The Lewis Theory allows us to predict the correct formulas for ionic compounds... [:F:] -1... :F:..

15. Predict the formula of the compound that forms between calcium and chlorine. Draw the Lewis dot symbols of the elements 1) Transfer all of the valance electrons from the metal to the nonmetal. 2) Add more to each atom as you go until all of the electrons are lost from the metal atoms and all the nonmetal atoms have 8 electrons: the Octet Rule. Ca ∙ ∙ Cl ∙ ∙ ∙ ∙∙ ∙∙ Ca ∙ ∙ Cl ∙ ∙ ∙ ∙∙ ∙∙ ∙ ∙ ∙ ∙∙ ∙∙ Ca 2+ CaCl 2

16. Covalent Bonds Exist mostly between two nonmetals. They are typical of molecular species. The atoms are bonded together to form molecules –strong attraction, but weaker, in general, than ionic bonds. Within a bond each atom donates an electron to form sharing pairs of electrons that attain octets around each atom. Covalent molecules are generally weakly attracted to each other. F F H F F H O H H O FF

17. Double Covalent Bond Two atoms share two pairs of electrons –4 electrons, total The double bond is shorter and stronger than a single bond. O O O O OO

18. Triple Covalent Bond Two atoms sharing 3 pairs of electrons –6 electrons, total The triple bond is shorter and stronger than either the single or double bond. N N N N NN

19. Lone Pair Electrons Electrons that are shared by atoms are called bonding pairs. Electrons that are not shared by atoms but belong to a particular atom are called lone pairs –also known as nonbonding pairs O S O Lone Pairs Bonding Pairs

20. Polyatomic Ions The polyatomic ions are attracted to opposite ions by ionic bonds –They tend to form crystal lattices. Atoms within the polyatomic ion are held together by covalent bonds.

21. Writing Lewis Structures for Covalent Molecules 1)Attach the atoms together in a skeletal structure: –The most metallic element is generally central. –The halogens and hydrogen are generally terminal. –Many molecules tend to be symmetrical. –In oxyacids, the acid hydrogens are attached to an oxygen atom. 2)Calculate the total number of valence electrons that are available for bonding: –Use the group number on the periodic table. 3)Attach the atoms with pairs of electrons and subtract these electrons from the total number of valence electrons available: –bonding electrons

22. 5)If there are not enough electrons to complete an octet on the central atom, bring in pairs of electrons from an attached atom to share with the central atom until it has an octet. 4) Add the remaining electrons in pairs to complete the octets of all the atoms: - remember H only takes 2 electrons (duet rule) - don’t forget to keep subtracting from the total - complete octets on the terminal atoms first, then work toward the central atoms

23. Example HNO 3 1)Write the skeletal structure –since this is an oxyacid, H is on the outside attached to one of the O’s; N is central 2)Count Valence Electrons and Subtract Bonding Electrons from Total N = 5 H = 1 O 3 = 3x6 = 18 Total = 24 e - Electrons Start24 Used8 Left16 3) Complete the Octets from the outside-in. H is already complete with 2 electrons that form one bond.

24. 5) If the central atom does not have an octet, bring in electron pairs from the outside atoms to share in common bonding patterns, if possible. 4)Re-Count the total number of valence electrons. N = 5 H = 1 O 3 = 3x6 = 18 Total = 24 e- Electrons Start24 Used8 Left16

25. Writing Lewis Diagrams for Polyatomic Ions 3)Complete the Octets from the outside-in 1)Write the skeletal structure for NO 3 - –N is central because it is the most metallic 2)Count Valence Electrons and Subtract Bonding Electrons from Total N = 5 O 3 = 3x6 = 18 (-) = 1 Total = 24 e - Electrons Start24 Used6 Left18

26. 4)Re-Count Electrons N = 5 O 3 = 3∙6 = 18 (-) = 1 Total = 24 e- Electrons Start24 Used6 Left18 5)If the central atom does not have an octet, bring in electron pairs from the outside atoms to share with the central atom –follow the common bonding patterns if possible - -

27. Exceptions to the Octet Rule H & Li, lose one electron to form cations –Li now has an electron configuration like He –H can also share or gain one electron to have a configuration like He Be shares 2 electrons to form two single bonds B shares 3 electrons to form three single bonds There can be expanded octets for elements in Period 3 or below –use empty valence d orbitals some molecules have odd numbers of electrons –NO

28. Resonance Lewis structures often do not accurately represent the electron distribution in a molecule –Lewis structures imply that O 3 has a single (147 pm) and double (121 pm) bond, but the actual bond length is between121 and 147, (128 pm) The real molecule is a hybrid of all possible Lewis structures Resonance stabilizes the molecule –maximum stabilization comes when resonance forms contribute equally to the hybrid ↔

29. Molecular Geometry Molecules are 3-dimensional objects We often describe the shape of a molecule with terms that relate to geometric figures These geometric figures have characteristic “corners” that indicate the positions of the surrounding atoms with the central atom in the center of the figure The geometric figures also have characteristic angles that we call bond angles Linear –2 atoms on opposite sides of a central atom –180° bond angles Trigonal Planar –3 atoms form a triangle around the central atom –Planar –120° bond angles Tetrahedral –4 surrounding atoms form a tetrahedron around the central atom –109.5° bond angles 180° 120° 109.5°

30. Predicting Molecular Geometry VSEPR Theory –Valence Shell Electron Pair Repulsion The shape around the central atom(s) can be predicted by assuming that the areas of electrons on the central atom will try to get as far from each other as possible –areas of negative charge will repel each other Each Bond counts as 1 area of electrons –single, double or triple all count as 1 area Each Lone Pair counts as 1 area of electrons –Even though lone pairs are not attached to other atoms, they do “occupy space” around the central atom Lone pairs take up slightly more space than bonding pairs –Effects bond angles

31. Linear –2 areas of electrons around the central atom, both bonding Or a two atom molecule –180° Bond Angles Trigonal - 3 areas of electrons around the central atom - 120° bond angles - All Bonding = trigonal planar - 2 Bonding + 1 Lone Pair = bent Tetrahedral - 4 areas of electrons around the central atom - 109.5° bond angles - All Bonding = tetrahedral -3 Bonding + 1 Lone Pair = trigonal pyramid - 2 Bonding + 2 Lone Pair = bent

32. Adding Dipole Moments