Structure Determines Properties

1. Structure Determines Properties Chapter 1

2. Overview • A review the relationship between structure and properties is what chemistry is all about. • Lewis Structures • Arrhenius, Bronsted-Lowry, and Lewis pictures of acids and bases • Concludes with the effects of structure on acidity and basicity.

3. Atoms, Electrons, and Orbitals • Each atom has • Protons, + • Electrons, - • Neutrons, no charge • A neutral atom • # protons = # electrons • Atomic number (Z) • Number of protons

4. Atoms, Electrons, and Orbitals • Wave functions are the mathematical descriptions giving the probability of locating the electron. • Wave functions are called orbitals. • Orbitals are descried by specifying their size, shape, and directional properties.

5. Atoms, Electrons, and Orbitals • Rules of electron configuration • Aufbau Principle • The rule that electrons occupy the orbitals of lowest energy first. • Pauli Exclusion Principle • An atomic orbital may describe at most two electrons, each with opposite spin direction. • Hund’s Rule • Electrons occupy orbitals of the same energy in a way that makes the number or electrons with the same spin direction as large as possible.

6. Atoms, Electrons, and Orbitals • The letter s is preceded by the principal quantum number which specifies the shell and is related to the energy of the orbital. • Hydrogen: 1s1 • Helium: 1s2 • Electrons possess the property of spin. • The spin quantum number of an electron can have a value of either +1/2 or -1/2.

7. Atoms, Electrons, and Orbitals • Valence electrons of an atom are the outermost electrons, the ones most likely to be involved in chemical bonding and reactions • Four main group elements, the number of valence electrons is equal to its group number in the periodic table. • Structure determines properties and the properties of atoms depend on atomic structure.

8. Atoms, Electrons, and Orbitals • Neutral atoms have as many electrons as the number of protons in the nucleus. • Those electrons occupy orbitals in order of increasing energy, with no more than two electrons in any one orbital. • most frequently encountered atomic orbital in this text are s orbitals and p orbital.

9. Ionic Bonds • The attractive forces between atoms in a compound is a chemical bond • Ionic bond is the fore of attraction between oppositely charged ions. • Ionic bonds are very common in inorganic compounds, but rare in organic ones.

10. Ionic Bonds • Positively charged ions are called cations • Na(g) → Na+(g) + e- • Negatively charged ions are called anions. • Cl(g) + e- → Cl-(g)

11. Ionic Bonds • An ionic bond is the force of electrostatic attraction between two oppositely charged ions. • Ionic bonds in which carbon is the cation or anion are rare.

12. Covalent Bonds, Lewis Structures and Octet Rule • Covalent or shared electron pair is the bond formed when electrons are shared. • Lewis structures are structural formulas in which electrons are represented as dots.

13. Covalent Bonds, Lewis Structures and Octet Rule • The valence electrons that are not involved in the bonding are called unshared pairs. • Octet rule • In forming compounds, they gain, lose, or share electrons to achieve a stable electron configuration characterized by eight valence electrons.

14. Covalent Bonds, Lewis Structures and Octet Rule • The most common kind of bonding involving carbon is covalent bonding. • A covalent bond is the sharing of a pair of electrons between two atoms. • Lewis structures are written on the basis of the octet rule, which limits second-row elements to no more eight electrons in their valence shells. • In most of its compounds, carbon has four bonds.

15. Double and Triple Bonds • By pairing the unshared electrons of one carbon with its counterpart of the other carbon, a double bond results and the octet rule is satisfied for both carbons.

16. Double and Triple Bonds • The ten valence electrons of acetylene (C2H2) can be arranged in a structural formula that satisfies the octet rule when six of them are shared in a triple bond between the carbons.

17. Double and Triple Bonds • Many organic compounds have double or triple bonds to carbon. Four electrons are involved in a double bond: six in a triple.

18. Polar Covalent Bonds, Electronegativity and Bond Dipoles • Electrons in covalent bonds are not necessarily shared equally by the two atoms that they connect.

19. Polar Covalent Bonds, Electronegativity and Bond Dipoles • If one atom has a greater tendency to attract electrons toward itself than the other, the electron distribution is polarized, and the bond is described as polar covalent. • The tendency of an atom to attract the electrons in a covalent bond toward itself defines its electronegativity.

20. Polar Covalent Bonds, Electronegativity and Bond Dipoles • A dipole exists whenever opposite charges are separated from each other, and a dipole moment m is the product of the amount of the charge e multiplied by the distance d between the centers of charge.

21. Polar Covalent Bonds, Electronegativity and Bond Dipoles • An important difference between a C-H bond and a C-O bond, and that is the direction of the dipole moment. In a C-H bond the electrons are drawn away from H, toward C. In a C-O bond, electrons are drawn from C toward O.

22. Polar Covalent Bonds, Electronegativity and Bond Dipoles • When two atoms that differ in electronegativity are covalently bonded, the electrons in the bond are drawn toward the more electronegative element.

23. Formal Charge • Lewis structures frequently contain atoms that bear a positive or negative charge. • Electrons in covalent bonds are counted as if they are shared equally by the atoms they connect, but unshared electrons belong to a single atom.

24. Formal Charge • Formal charge = Group number in periodic table – Electron count • Electron count = ½(Number of shared electrons) + Number of unshared electrons

25. Formal Charges • It will always be true that a covalently bonded hydrogen has no formal charge (formal charge = 0). • It will always be true that a nitrogen with four covalent bonds has a formal charge of +1. (A nitrogen with four covalent bonds cannot have unshared pairs, because of the octet rule.) • It will always be true that an oxygen with two covalent bonds and two unshared pairs has no formal charge. • It will always be true that an oxygen with one covalent bond and three unshared pairs has a formal charge of -1.

26. Formal Charges • Counting electrons and assessing charge distribution in molecules is essential to understanding how structure affects properties. A particular atom in a Lewis structure may be neutral, positive, or negatively charged. The formal charge of an atom in the Lewis structure of a molecule can be calculated by comparing its electron count with that of the neutral atom itself.

27. Structural Formulas of Organic Molecules • Different compounds that have the same molecular formula are called isomers. If they are different because their atoms are connected in a different order, they are called constitutional isomers.

28. Structural Formulas of Organic Molecules • Isomers can be either constitutional isomers (differ in connectivity) or stereoisomers (differ in arrangement of atoms in space).

29. Resonance • Sometimes more than one Lewis formula can be written for a molecule, especially if the molecule contains a double or triple bond.

30. Resonance • Lewis formulas show electrons as being localized; they either are shared between two atoms in a covalent bond or are unshared electrons belonging to a single atom. In reality, electrons distribute themselves in the way that leads to their most stable arrangement. This sometimes means that a pair of electrons is delocalized, or shared by several nuclei.

31. The Shapes of Simple Molecules • The shapes of molecules can often be predicted on the basis of valence shell electron-pair repulsion.

32. Molecular Dipole Moments • Knowing the shape of a molecule and the polarity of its various bonds allows the presence or absence of a molecular dipole moment and its direction to be predicted.

33. Curved Arrows and Chemical Reactions • Curved arrows increase the amount of information provided by a chemical equation by showing the flow of electrons associated with bond making and bond breaking.

34. Acids and Bases: Arrenhius View • According to the Arrhenius definitions, an acid ionizes in water to produce protons (H+) and a base produces hydroxide ions (OH-). The strength of an acid is given by its equilibrium constant Ka for ionization in aqueous solution. • Or more conveniently by its pKa: pKa = - log10Ka

35. Acids and Bases: Bronsted-Lowry • According to the Bronsted-Lowry definitions, an acid is a proton donor and a base is a proton acceptor. B: + H-A → B-H + :A- • The Bronsted-Lowry approach to acids and bases is more generally useful than the Arrhenius approach.

36. How Structure Affect Acid Strength • The strength of an acid depends on the atom to which the proton is bonded. • The two main factors are the strength of the H-X bond and the electronegativity of X. • Bond strength is more important for atoms in the same group of the periodic table; electronegativity is more important for atoms in the same row. • Electron delocalization in the conjugate base, usually expressed via resonance between Lewis structures, increase acidity by stabilizing the conjugate base.

37. Acid-Base Equilibria • The position of equilibrium in an acid-base reaction lies to the side of the weaker acid. Stronger acid + stronger base Weaker acid + weaker base • The equilibrium will be to the side of the acid that holds the proton more tightly. Keq =

38. Lewis Acids and Lewis Bases • A Lewis acid is an electron-pair acceptor. A Lewis-base is an electron-pair donor. • The Lewis approach incorporates the Bronsted-Lowry approach as a subcategory in which the atom that accepts the electron pair in the Lewis acid is a hydrogen.