Chapter 1 Structure Determines Physical and Chemical Properties

# Chapter 1 Structure Determines Physical and Chemical Properties

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## Chapter 1 Structure Determines Physical and Chemical Properties

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1. 1.1Atoms, Electrons, and Orbitals

2. + – • Atoms are composed of • Protons • positively charged • mass = 1.6726 X 10-27 kg • Neutrons • neutral • mass = 1.6750 X 10-27 kg • Electrons • negatively charged • mass = 9.1096 X 10-31 kg

3. A X Z Atomic Number and Mass Number • Atomic number (Z) = number of protons in nucleus • (this must also equal the number of electrons in neutral atom) • Mass number (A) = sum of number of protons + neutrons in nucleus

4. Schrödinger Equation • Schrödinger combined the idea that an electron has wave properties with classical equations of wave motion to give a wave equation for the energy of an electron in an atom. • Wave equation (Schrödinger equation) gives aseries of solutions called wave functions ( ).

5. Wave Functions • Only certain values of  are allowed. • Each  corresponds to a certain energy. • The probability of finding an electron at a particular point with respect to the nucleus isgiven by 2. • Each energy state corresponds to an orbital.

6. Figure 1.1 Probability distribution (2) for an electron in a 1s orbital.

7. A boundary surface encloses the regionwhere the probability of finding an electronis high—on the order of 90-95% 1s 2s Figure 1.2 Boundary surfaces of a 1s orbitaland a 2s orbital.

8. Quantum Numbers • Each orbital is characterized by a unique set of quantum numbers. • The principal quantum number nis a wholenumber (integer) that specifies the shell and isrelated to the energy of the orbital. • The angular momentum quantum numberis usually designated by a letter (s, p, d, f, etc) and describes the shape of the orbital.

9. s Orbitals • s Orbitals are spherically symmetric. • The energy of an s orbital increases with thenumber of nodal surfaces it has. • A nodal surface is a region where the probabilityof finding an electron is zero. • A 1s orbital has no nodes; a 2s orbital has one;a 3s orbital has two, etc.

10. The Pauli Exclusion Principle • No two electrons in the same atom can havethe same set of four quantum numbers. • Two electrons can occupy the same orbitalonly when they have opposite spins. • There is a maximum of two electrons per orbital.

11. 1s 2p 2s First Period • Principal quantum number (n) = 1 • Hydrogen Helium • Z = 1 Z = 2 • 1s 1 1s 2 H He

12. p Orbitals • p Orbitals are shaped like dumbells. • Are not possible for n = 1. • Are possible for n = 2 and higher.

13. p Orbitals • p Orbitals are shaped like dumbells. • Are not possible for n = 1. • Are possible for n = 2 and higher. • There are three p orbitals for each value of n (when n is greater than 1).

14. p Orbitals • p Orbitals are shaped like dumbells. • Are not possible for n = 1. • Are possible for n = 2 and higher. • There are three p orbitals for each value of n (when n is greater than 1).

15. p Orbitals • p Orbitals are shaped like dumbells. • Are not possible for n = 1. • Are possible for n = 2 and higher. • There are three p orbitals for each value of n (when n is greater than 1).

16. 1s 2p 2s Z Li 3 Be 4 B 5 C 6 Second Period • Principal quantum number (n) = 2

17. 1s 2p 2s Z N 7 O 8 F 9 Ne 10 Second Period

18. 1.2Ionic Bonds

19. Na+ (cation) Cl– (anion) Ionic Bonding • An ionic bond is the force of electrostaticattraction between oppositely charged ions

20. Ionic Bonding • Ionic bonds are common in inorganic chemistrybut rare in carbon atoms. • Carbon shows less of a tendency to form cationsthan metals do, and less of a tendency to formanions than nonmetals.

21. 1.3Covalent Bonds, Lewis Structures,and the Octet Rule

22. The Lewis Model of Chemical Bonding • In 1916 G. N. Lewis proposed that atomscombine in order to achieve a more stableelectron configuration. • Maximum stability results when an atomis isoelectronic with a noble gas. • An electron pair that is shared between two atoms constitutes a covalent bond.

23. . . H H : H H Covalent Bonding in H2 Two hydrogen atoms, each with 1 electron, • Sharing the electron pair gives each hydrogen an electron configuration analogous to helium. can share those electrons in a covalent bond.

24. .. .. . . : : F F .. .. .. .. : : : F F .. .. Covalent Bonding in F2 Two fluorine atoms, each with 7 valence electrons, • Sharing the electron pair gives each fluorine an electron configuration analogous to neon. can share those electrons in a covalent bond.

25. .. .. : : : F F .. .. The Octet Rule In forming compounds, atoms gain, lose, or share electrons to give a stable electron configuration characterized by 8 valence electrons. • The octet rule is the most useful in cases involving covalent bonds to C, N, O, and F.

26. . .. . . . : C F .. . .. : : F .. .. .. : : : F C : F .. .. .. : : F .. Example Combine carbon (4 valence electrons) andfour fluorines (7 valence electrons each) to write a Lewis structure for CF4. The octet rule is satisfied for carbon and each fluorine.

27. .. : : F .. .. .. : : : F C : F .. .. .. : : F .. : : F .. .. : : C F F .. .. : : F .. Example It is common practice to represent a covalentbond by a line. We can rewrite .. as

28. 1.4Double Bonds and Triple Bonds

29. .. .. .. .. : : : : : : : : C C O O O O : : : : : : H H C N C N Inorganic Examples Carbon dioxide Hydrogen cyanide

30. H H H H .. .. : : : : H H C C C C H H : : : : : H H H H C C C C Organic Examples Ethylene Acetylene

31. 1.5Polar Covalent Bonds and Electronegativity

32. Electronegativity Electronegativity is a measure of the abilityof an element to attract electrons toward itself when bonded to another element. • An electronegative element attracts electrons. • An electropositive element releases electrons.

33. Pauling Electronegativity Scale • Electronegativity increases from LEFT to RIGHT and BOTTOM to TOP

34. .. .. : : : : N N F F .. .. Generalization • The greater the difference in electronegativitybetween two bonded atoms; the more polar the bond. H—H nonpolar bonds connect atoms ofthe same electronegativity

35.  .. ..     : H H O H F .. .. Generalization • The greater the difference in electronegativitybetween two bonded atoms; the more polar the bond.    : : C O O .. .. polar bonds connect atoms ofdifferent electronegativity

36. ..   : H F .. Electrostatic Potential Maps • Electrostatic potential maps show the chargedistribution within a molecule. Solidsurface Red is negative charge;blue is positive.

37. ..   : H F .. Electrostatic Potential Maps • Electrostatic potential maps show the chargedistribution within a molecule. Transparentsurface Red is negative charge;blue is positive.

38. ..   : H F .. Electrostatic Potential Maps • Electrostatic potential maps show the chargedistribution within a molecule. + - H Li Red is negative charge;blue is positive.

39. 1.6Structural Formulas of Organic Molecules

40. Constitution The order in which the atoms of a molecule are connected is called its constitution or connectivity. The constitution of a molecule must be determined in order to write a Lewis structure.

41. H H H C H C C H : H : H O (CH3)2CHOH or CH3CHCH3 H OH Condensed structural formulas Lewis structures in which many (or all) covalent bonds and electron pairs are omitted. can be condensed to:

42. CH3CH2CH2CH3 is shown as CH3CH2CH2CH2OHis shown as OH Bond-line formulas Omit atom symbols. Represent structure by showing bonds between carbons and atoms other than hydrogen. Atoms other than carbon and hydrogen are called heteroatoms.

43. H Cl Cl C H2C CH2 H2C CH2 C H H Bond-line formulas is shown as Omit atom symbols. Represent structure by showing bonds between carbons and atoms other than hydrogen. Atoms other than carbon and hydrogen are called heteroatoms.

44. 1.7Formal Charge • Formal charge is the charge calculated for an atom in a Lewis structure on the basis of an equal sharing of bonded electron pairs.

45. .. : O .. N H O .. : : O Nitric acid Formal charge of H • We will calculate the formal charge for each atom in this Lewis structure. ..

46. .. : O .. N H O .. : : O Nitric acid Formal charge of H • Hydrogen shares 2 electrons with oxygen. • Assign 1 electron to H and 1 to O. • A neutral hydrogen atom has 1 electron. • Therefore, the formal charge of H in nitric acid is 0. ..

47. .. : O .. N H O .. : : O Nitric acid Formal charge of O • Oxygen has 4 electrons in covalent bonds. • Assign 2 of these 4 electrons to O. • Oxygen has 2 unshared pairs. Assign all 4 of these electrons to O. • Therefore, the total number of electrons assigned to O is 2 + 4 = 6. ..

48. .. : O .. N H O .. : : O Nitric acid Formal charge of O • Electron count of O is 6. • A neutral oxygen has 6 electrons. • Therefore, the formal charge of O is 0. ..

49. .. : O .. N H O .. : : O Nitric acid Formal charge of O • Electron count of O is 6 (4 electrons from unshared pairs + half of 4 bonded electrons). • A neutral oxygen has 6 electrons. • Therefore, the formal charge of O is 0. ..