Chapter 1: Molecular Structure Review

1. Chapter 1: Molecular Structure Review • Organic chemistry: The study of the compounds of carbon. • Over 20 million organic compounds have been identified. • C is a unique atom among the elements. • It is small. • It forms strong covalent bonds with itself. • It can form up to 4 bonds with other atoms. • It forms single, double and triple bonds. • It is moderately high in electronegativity (2.5). • It forms strong bonds with C, H, O, N, S, the halogens, and some metals.

2. In organic chemistry, we make things! • Flavorings • Pharmaceuticals • Polymers • Fuels • Solvents and cleaners • Paints and adhesives • DNA and proteins • Cosmetics • etc…………………………!!! • Most organic chemicals have two sources. • Crude oil • Biomaterials

3. Very often to make what we want, we need to do a reaction to change one organic molecule into another. • For example, ethanol can be changed into acetaldehyde (used to make other molecules!) or acetic acid (vinegar).

4. Overview of an Atom A small dense nucleus contains positively charged protons, neutrons, and most of the mass of the atom. Extranuclear space, diameter 10-10 m, which contains negatively charged electrons. Electrons are waves and particles. We often want to think about electrons as clouds (orbitals). Electrons and Atomic Structure

5. An orbital (an electron cloud) can be represented in several different ways. • s orbital • p orbital

6. Electrons in an atom are organized into shells, subshells and orbitals. • Shells are defined by their principal quantum number, n. n = 1, 2, 3, 4, … • The smaller the n, the more negative the energy. • Negative energy is energy of attraction. (More negative energy, more attraction.)

7. Each shell has “n” number of subshells. • Subshells are labeled with letters: s, p, d and f. • The first shell has one subshell: 1s • The second shell has two subshells: 2s and 2p • The third shell has three subshells: 3s, 3p and 3d • The fourth shell has four subshells: 4s, 4p, 4d and 4f • Each subshell type contains a specific number of orbitals. • s (1 orbital), p (3 orbitals), d (5 orbitals), f (7 orbitals)

8. Each orbital contains a maximum of two electrons. • If two electrons are in an orbital, they must have opposite electron spins. • Pauli Exclusion Principle states that paired electrons in an orbital (or a covalent bond) must have opposite spins.

9. Atomic Electronic Configurations • Three principles: • Aufbau: orbitals fill from lowest energy to highest energy • Pauli exclusion: only two electrons per orbital, spins must be paired • Hund’s rule: e- tend to occupy a set of orbitals with the same energy (degenerate orbitals) before pairing up

10. All the electrons in an atoms can be described with an electron configuration. • The periodic table is organized based on where an atom’s valence (outer) electrons are.

12. Placement of electrons in an atom can be represented visually with an orbital energy level diagram. 3p 3s 2p 2s 1s carbon neon oxygen sulfur

13. Chemical Bonding Review • Electronegativity: the ability of an element to attract electrons when part of a compound. • Electronegativity is a periodic trend. • Fluorine is the most electronegative element.

14. Electronegativity is used to understand the bonding between atoms in a compound. • Two atoms with low electronegativities bond with metallic bonding. • Two atoms with high electronegativities bond with covalent bonding. • One atom with low electronegativity and one atom with high electronegativity bond with ionic bonding. • The difference of electronegativity is sometimes used to describe bonding. • Ionic bond: a chemical bond between an anion & a cation (ΔEN >1.9) • NaF 4.0 – 0.9 = 3.1  ionic bond • Covalent bond: • Nonpolar: (ΔEN < 0.5) • Polar: ΔEN, [0.5, 1.9] • HCl 3.0 – 2.1 = 0.9  polar covalent

15. Lewis Structure Review • Lewis structures are a visual representation of the valence electrons in a compound (especially molecular compounds). • The number of valence electrons for an atom is found by the column of the periodic table the element is in. • The Lewis structure for a single atom has its valence electrons arranged on the sides of a square (according to Hund’sPrinciple).

16. Drawing Lewis Structures • Procedure for drawing Lewis structures • Sum total number of valence electrons for all atoms in the structure. • If structure is an ion, add or subtract the appropriate number of electrons. • Arrange atoms in a skeleton. • Carbon will be somewhere in the middle. • Hydrogen and the halogens will always be on the outside. • Oxygen, nitrogen, sulfur, etc… will often but not always be on the outside. • Draw bonds between atoms • Carbon should have 4 bonds. • Oxygen and sulfur should have 2 bonds. • Nitrogen and phosphorus should have 3 bonds. • Hydrogen and the halogens will have 1 bond. • Complete octets of the electronegative atoms • If some atoms are not satisfying the octet rule, rearrange electrons to give all atoms the appropriate number of bonds.

17. Simple Lewis Structures

18. Lewis Structures with Double and Single Bonds

19. Formal Charge • Formal charge: the charge on an atom in a molecule or polyatomic ion. • Write a Lewis structure for the molecule or ion. • Assign each atom all its unshared (nonbonding) electrons and one-half its shared (bonding) electrons. • Compare this number with the number of valence electrons in the neutral, unbonded atom. • If the number is less than that assigned to the unbonded atom, the atom has a positive formal charge. • If the number is greater, the atom has a negative formal charge.

20. S C N Method to Determine Formal Charge Draw a circle (with a pencil or in your mind’s eye) around an atom in Lewis structure, ensuring that the circle cuts through any bonds that the atom has. Determine the normal # of valence e- for the atom. Subtract the number of e- inside the circle (keeping in mind that since the circle cuts through the e- atom’s bonds; therefore, only e- per bond is inside the circle). Let us examine the thiocyanate ion to show how the method is applied. For the Sulfur atom 1. Normal # of valence e- for the S atom = 6 2. e- inside the circle = 6 3. Formal charge of S atom: 6 – 6 = 0 For the Carbon atom 1. Normal # of valence e- for the C atom = 4 2. e- inside the circle = 4 3. Formal charge of S atom: 4 – 4 = 0 For the Nitrogen atom 1. Normal # of valence e- for the N atom = 5 2. e- inside the circle = 6 3. Formal charge of S atom: 5 – 6 = -1

21. Resonance • A bond, (the second and/or third bond in a double or triple bond), may involve more than two atoms. • Or there may be alternate ways to put lone pairs of electrons into a Lewis structure. • A conventional Lewis structure has no way of illustrating a bond that involves more than three atoms. • A “work-around” solution is to employ resonance structures. • Thus a more accurate structure for a particular compound is a blend (sometimes equal, other times not) of the different resonance structures. • A double-headed arrow is used to connect resonance structures.

22. All acceptable contributing structures must: 1. Have the same number of valence electrons. 2. Obey the rules of covalent bonding. • No more than 2 electrons in the valence shell of H. • No more than 8 electrons in the valence shell of a 2nd period element. • 3rd period elements may have up to 12 electrons in their valence shells. 3. Differ only in distribution of valence electrons. 4. Have the same number of paired and unpaired electrons.

23. An example of resonance: carbonate ion • The carbonate ion has one double bond, but it is distributed equally among all three C – O pairs. • The resonance structures show that a C – O pair doesn’t have a single or a double bond, but rather each pair has 1 1/3 bonds. • In the same vein, we can say that each oxygen atom doesn’t have zero charge or negative one charge; but rather, each oxygen atom has – 2/3 charge.

24. Other examples of resonance

25. Nonequivalent resonance structures • Nitrous oxide, N2O, laughing gas, can have multiple Lewis structures that satisfy the octet rule. • We surmise that the actual structure is a blend of the Lewis structures, but in this case, the blend is not an equal blend of all of the structures. Some structures contribute more than others.

26. Some resonance structures are more important than others. • Structures that satisfy the octet are more important than those that don’t. • More important structures have atoms that have small formal charges. • When a structure has formal charges on the atoms, the most important structure places negative formal charge on the most electronegative atom. • The more important structures are those involve a minimum of charge separation.

27. Molecular Geometry • The geometry (shape) of a molecule is described using the Valence-Shell Electron-Pair Repulsion (VSEPR) theory. • VSEPR is based on two concepts. • Atoms are surrounded by regions (domains) of electron density. • Regions of electron density repel each other. • Electron domain geometry is determined first from number of domains surrounding an atom • Domains can be single, double or triple bonds. • Domains can be a nonbonding pair of electrons or a single unpaired electron (rare). • Atoms follow where electrons domains go; thus, molecular geometry is determined after electron domain geometry.

28. Electron domain geometry is determined from number of domains surrounding an atom. • Bond angles are set by angles in an electron domain geometry.

29. Molecular Geometries with 2 Electron Domains • An atom with 2 electron domains must have a linear electron domain geometry. • The electron domains have a 180 angle of separation. • If 1 or 2 atoms are bonded to the electron domains, the molecular geometry is linear.

30. Molecular Geometries with 3 Electron Domains • An atom with 3 electron domains must have a trigonal planar electron domain geometry. • The electron domains have a 120 angle of separation. • If 2 atoms are bonded, the molecular geometry is bent. • If 3 atoms are bonded, the molecular geometry is trigonal planar.

31. Molecular Geometries with 4 Electron Domains • An atom with 4 electron domains must have a tetrahedral electron domain geometry. • The electron domains have a 109.47 angle of separation.

32. Molecular Geometries with 4 Electron Domains • If 1 atom is bonded to the electron domains, the molecular geometry is linear. • If 2 atoms are bonded, the molecular geometry is bent.

33. Molecular Geometries with 4 Electron Domains • If 3 atoms are bonded, the molecular geometry is trigonal pyramidal. • If 4 atoms are bonded, the molecular geometry is tetrahedral.

34. Polar and Nonpolar Molecules • A molecule is polar if: • It is asymmetrical. (It is unbalanced.) • It has polar covalent bonds (polar bonds). • Difference in electroneg. between atoms is greater than 0.5. • An important aside: The carbon-hydrogen bond is not polar.

35. A water molecule is a polar molecule. • It is asymmetrical. (We can identify two different halves.) • It has polar bonds. • On the ESP (electrostatic potential) map to the right, red indicates negative charge and blue indicates positive charge. • This difference should make sense since oxygen is more electronegative than hydrogen.

36. A carbon dioxide molecule is a nonpolar molecule. • It is symmetrical. (We cannot identify two different halves.) • It is nonpolar despite that it has polar bonds. • The polar bonds cancel each other out because the molecular geometry.

37. An ammonia molecule has three polar covalent bonds, and because of its geometry, is a polar molecule.

38. Chloromethane and formaldehyde are polar molecules. • Acetylene is a nonpolar molecule. A more accurate ESP map of acetylene

39. Intermolecular Forces • The attraction that molecules have for each other has significant consequences for a substance’s chemical and physical properties. • The basic principles needed to understand intermolecular forces are • Opposite charges attract. • The greater charges, the greater the attraction.

40. We’ll consider three types of charge distributions when examining intermolecular forces. • Ions • Dipoles • Induced Dipoles • A molecule with a dipole (or dipole moment) is a polar molecule and vice versa. • An induced dipole is created when a molecule’s electron cloud is distorted by another charge (ion, dipole, induced dipole). • The strength of an induced dipole depends on a molecule’s polarizability which in turn depends on the number of electrons a molecule has.

41. Intermolecular forces are classified based on the types of charges that are being attracted to each other. • Ion – dipole forces • Dipole – dipole forces • Hydrogen bonding • Induction forces (dipole – induced dipole) • (London) dispersion forces (induced dipole – induced dipole)

42. London Dispersion Forces • All molecules have dispersion forces [all molecules have electron clouds that have induced (temporary) dipoles.] • The positive side of the induced dipole of one molecule is attracted the negative side of another molecule’s induced dipole.

43. Dipole forces • When the positive end of a polar molecule attracts the negative end of another polar molecule (and vice versa); the attraction is called a dipole force. (also known as the dipole-dipole force)

44. Hydrogen Bonding • A special case of dipole force (stronger) • Molecules that have hydrogen bonding must have a hydrogen atom covalently bonded to a highly electronegative atom (N, O or F) • Hydrogen of one molecule (+ charge) is attracted to lone pair of electrons of another molecule (- charge)

45. Hydrogen Bonding • O-H is more polar than N-H, so they create stronger hydrogen bonding

46. Boiling Points and Melting Points • Intermolecular forces between molecules affect the melting points and boiling points. • Stronger intermolecular forces make higher melting and boiling points. • Knowing boiling points and melting points of molecules provides insight into the structure of molecules, since intermolecular forces depend on structure.

47. Hydrogen bonding is stronger than the dipole force. • Molecules with more hydrogen bonding have higher melting and boiling points. • Hydrogen bonding with the O – H bond is stronger than the N – H bond.

48. Solubility • Intermolecular forces have a strong effect on the formation of solutions. • “Like dissolves like” • Polar solutes dissolve in polar solvents. (dipole forces or hydrogen bonding) • Ethanol dissolves in water • Nonpolar solutes dissolve in nonpolar solvents. (dispersion forces) • Oil dissolves in gasoline • Ionic substances are dissolved because of ion – dipole force between solute (ions) and solvent (dipole)

49. Valence Bond Theory of Covalent Bonding • A covalent bond forms when a portion of an atomic orbital of one atom overlaps a portion of an atomic orbital of another atom. • A bond that forms from overlap directly between two nuclei is a sigma () bond. • A bond that forms from p orbital that overlap off of the line connecting two nuclei is a pi () bond.

50. A Serious Problem! • Overlap of 2s atomic orbitals of one atom and 2p atomic orbitals of another atom would give bond angles of approximately 90°. • Instead we observe bond angles of approximately 109.5°, 120°, and 180°. • An Elegant Solution!! • Hybridization of atomic orbitals. • 2nd row elements use sp3, sp2, and sp hybrid orbitals for bonding. • We study three types of hybrid atomic orbitals: • sp3(one s orbital + three p orbitals give four sp3 hybrid orbitals). • sp2(one s orbital + two p orbitals give three sp2 hybrid orbitals). • sp (one s orbital + one p orbital give two sp hybrid orbitals). • Overlap of hybrid orbitals can form two types of bonds, depending on the geometry of the overlap: • σ bonds are formed by “direct” overlap. • π bonds are formed by “parallel” overlap.