Chem 150 Unit 1 - Molecules

1. Chem 150Unit 1 - Molecules • Organic and Biological Chemistry substances are made of molecules. Being able to predict the structures and interactions of molecules is important to our understanding of Organic and Biological Chemistry

2. Atomic Structure • Atoms are made up of three types of subatomic particles: • protons • neutrons • electrons • Two important properties of these particles are mass and charge.

3. Atomic Structure • The protons and neutrons packed in the center of the atom in a region called the nucleus. • The electrons occupy a diffuse region surrounding the nucleus. • The diameter of the volume occupied by the electrons is≈ 10,000 times that of the nucleus.

4. Atomic Structure • The electrons have a specific arrangement in the the diffuse cloud that surround the nucleus. • The Bohr model of the atom • Electrons circle the nucleus in specific orbits, with each orbit corresponding to a different energy level.

5. Atomic Structure • The electrons have a specific arrangement in the the diffuse cloud that surround the nucleus. • The Bohr model of the atom only works if the atom has a single electron. • Quantum mechanics was able to produce a better model.

6. Atomic Structure • According to quantum mechanics the electrons occupy orbitals in pairs. • The volumes of these orbitals have specific shapes.

7. Atomic Structure • The electrons in each energy level are divided into the orbitals. • Each orbital can hold a maximum of 2 electrons. • The number of orbitals in each energy level, or shell, increases with the number of the energy level.

8. Atomic Structure • For the ground states of the elements, the electrons fill the orbitals in a predictable way. • In the ground state the electrons occupy the available orbitals that are nearest to the center of the atom

9. Atomic Structure • The ground state distribution of the electrons is reflected in the organization of the periodic table

10. Valence Electrons • The electrons that are most important for determining the chemical and physical properties of an atom are those located on the surface in the highest occupied energy level. • This energy level is called the valence shell. • The electrons in the valence shell are called the valence electrons.

11. Valence Electrons • For the representative elements, the number of valence electrons can be determined from the group number for that element.

12. Electron Dot Structures • The Lewis electron dot structures are a convenient way of showing the number of valence electrons that an atom has.

13. The Octet Rule • Except for helium (He), the inert gases, (He, Ne, Ar, Kr, Xe & Rn) each have 8 electrons in their valence shell

14. The Octet Rule • The Octet Rule: • Atoms gain, lose, or share valence electrons in order to end up with eight valence electrons.

15. The Octet Rule • Chemistry is all about the different strategies that the elements use to attain 8 electrons in their valence shell • When an element has the same number of electrons as one of the inert gases, it is said to be isoelectron with that inert gas.

16. The Octet Rule • For more discussion on the octet rule, see the Chem 150Elaboration - The Octet Rule

17. Monoatomic Ions • A strategy that the representative elements use is to simply gain or lose electrons to become isoelectronic with one of the inert gases.

18. Monoatomic Ions • Predicting the charge on the transition metals is less straight forward than for the representative element.

19. Monoatomic Ions • Predicting the charge on the transition metals is less straight forward than for the representative element.

20. Compounds • Metals and non-metals combine to form binary ionic compounds. • Each monoatomic ion in a binary ionic compound has 8 valence electrons. 2 Na Cl2 2 Na+ + 2 Cl-

21. Covalent Bonds • Review the characteristics of Metals and Nonmetals, with respect to their approaches to adhering to the octet rule.

22. Covalent Bonds • Review the characteristics of Metals and Nonmetals, with respect to their approaches to adhering to the octet rule. • When metals combine with nonmetals, the metal give up their valence electrons to the nonmetal and each forms and ion: • An ionic compound is formed, which is held together by an ionic bond.

23. Covalent Bonds • When nonmetals combine with nonmetals, a different strategy is need. • When nonmetals combine with nonmetals, they share valance electrons. • The electrons are shared in pairs to form a covalent bond.

24. Compounds con’d • Compounds that are formed from nonmetals are held together by covalent bonds. • A group of atoms held together by covalent bonds is called a molecule. • Compounds that are composed of molecules are called molecular compounds. • Unlike ionic compounds, the group of atoms in a molecular compound stay together as a group when the compound changes it state (solid, liquid, gas or solution) • The formula that describes the composition of a molecule is called a molecular formula.

25. Compounds con’d • The structures of molecules can be represented by either electron dot structural formulas • or • line-bond structural formulas.

26. Compounds con’d • Compounds that are formed from nonmetals are held together by covalent bonds. • In some molecules, more than one pair of electrons is shared to form double bonds (2 pairs of electrons) or triple bonds (3 pairs of electrons).

27. Compounds, con’d • For more discussion on the formation of compounds, see the Chem 150Elaboration - Compounds

28. Compounds, con’d • Naming binary compounds depends on type of compound you have. • Binary ionic compounds with a representative metal • Name of metal+ name of nonmetal with the “-ide” ending • NaCl Sodium chloride • MgF2 Magnesium fluoride • Binary ionic compound with a transition metal • Name the metal with a Roman numeral to indicate the charge+ name of nonmetal with the “-ide” ending • CuCl Copper(I) chloride • CuCl2 Copper(II) chloride • Binary molecular compound • Name the least electronegative element first with a prefix to indicate its number + name of the more electronegative element with a prefix to indicate its number and the “-ide” ending • SO3 Sulfur trioxide • N2O5 Dinitrogen pentoxide

29. Polyatomic Ions • There are some species that are held together by covalent bond, but which are also ionic.

30. Polyatomic Ions

31. Structural Formulas • Structural formulas are used to explicitly show what atom is connected to what atom in a molecule. • Some molecules share the same molecular formula • These are called isomers These all have the same molecular formula: C3H8O

32. Structural Formulas • The different types of structural formulas include: • Electron dot structural formula • Line-bond structural formula • Condensed structural formula • Skeletal structural formula

33. Compounds, con’d • For more discussion on structural formulas, see the Chem 150Elaboration - Structural Formulas

34. Noncovalent Interactions • Molecules are held together by covalent bonds. • Molecules interact with other molecules through noncovalent interactions. • Characteristics of noncovalent interactions • Are much weaker than covalent bonds. • They are the interactions that hold molecules together in the solid, liquid and solution states. • They are easily disrupted by increasing the temperature. • The are primarily electrical in nature

35. Noncovalent Interactions • Charge/Charge Interaction • Also called a salt bridge. • Arises from permanent charges that attract or repel one another. • Permanent charges are also called formal charges.

36. Noncovalent Interactions • Calculating the formal charge on an atom in a molecule or polyatomic ion. • When calculating the number of electrons around an atom in a compound, divide the bonding electrons equally between the two atoms participating in a bond.

37. Noncovalent Interactions • Calculating the formal charge on an atom in a molecule or polyatomic ion. Formal Charge on hydrogens = 1 - 1 = 0 Formal Charge on nitrogen = 5 - 4 = +1

38. Noncovalent Interactions • Calculating the formal charge on an atom in a molecule or polyatomic ion. Formal Charge on singly bonded oxygens = 6 - 7 = -1 Formal Charge on doubly bonded oxygen = 6 - 6 = 0 Formal Charge on nitrogen = 5 - 4 = +1

39. δ- δ- δ- δ+ δ+ δ+ Noncovalent Interactions • Dipole/Dipole and Ion Dipole Interactions • Some molecules, which are uncharged, have their valence electrons distributed unevenly, leading to a molecule having a positive and a negative end. • This situation leads to a permanent dipole. • These dipoles can interact with ions and other dipoles.

40. Noncovalent Interactions • Determining if a molecule is polar or not: • Determining the electron distribution in a molecule is a very complicated calculation, there are, however, some simple rules that can be applied to determine if a molecule is polar or not. • Two questions are asked: • Does the molecule contain any polar covalent bond? • If so, are these bonds arranged in a way that they will cancel each other out?

41. Noncovalent Interactions • To answer the first question, identify any polar covalent bonds that the molecule has: • This is done by comparing the electronegativity of the two atoms participating in each of the bonds. • If the one atom has a substantially higher electronegativity that the other, the valence electrons are not shared equally and the bond is polar.

42. Noncovalent Interactions • The electronegativity values for the elements (Table 4.1 in Raymond)

43. Noncovalent Interactions • The electronegativities can be used to predict whether a bond is ionic, polar covalent, or pure covalent: ΔE.N. > 2 0.5 > E.N.

44. Noncovalent Interactions • The electronegativities can be used to predict whether a bond is ionic, polar covalent, or pure covalent: • In general: • The bond is polar if N, O, F or Cl are bonded any of the other non-metals, including H. • The bond is nonpolar if C or H are bond to any of the other non-metals, except N, O, F or Cl; or if N, O, F or Cl are bonded to N, O, F or Cl

45. Noncovalent Interactions • Some examples of polar covalent bonds:

46. Noncovalent Interactions • If a molecule contains more than one polar covalent bond, you then need to determine the geometry and shape of the molecule around the atoms participating in the polar covalent bonds. • This is done by counting up the groups of electrons around each atom • Groups of electrons include: • A non-bonding pair of electrons • A single covalent bond • A double covalent bond (counts as a single group of electrons) • A triple covalent bond (counts as a single group of electrons)

49. Noncovalent Interactions • Use arrows to represent the polar bonds and look to see if the arrows cancel one another out

50. No Does the molecule contain any polar covalent bonds? Nonpolar Yes Determine the shape of the molecule. Yes Are the polar covalent bonds arranged in space in a way that causes them to cancel one another out? Nonpolar No Polar Noncovalent Interactions • Putting is all together: