Bonding

2. Bonding • Types of Bonds • Electronegativity • Polarity & Dipole Moment • Ions • Ionic Compounds • Lattice Energy • Covalent Bonds • Bond Energy • Lewis Structures • Resonance • Formal Charge • VESPR

3. Who bonds? Who doesn’t? • Atoms shift valence electrons to complete the outer energy level. • Gaining, losing or sharing to reach noble gas configuration • Monatomic elements do not easily form bonds

4. Types of Chemical Bonds • Intramolecular or Intermolecular bonds • Bonds form in order to lower the energy of being separated atoms

5. Chemical bonds form… • Electrical attraction between the nuclei of one atom and the valence electrons of another atom • Atoms arrange to lower the potential energy

6. Valence Electrons and Valence Shells • The outer energy level is called the VALENCE SHELL • Electrons in the outer energy level are called VALENCEELECTRONS

7. Valence Electrons • The number of electrons in the outer VALENCE shell can be determined by the location on the periodic chart. • Group 1A: 1e- Group 2A: 2e- Group 3A: 3e- etc Group 3B: 3e- Group 4B: 4e-

8. The Octet Rule Na has 1 valence electron. It loses that electron. Na+ can bond with a negative ly charge atom • To become like a noble gas & require an octet or eight electrons in the outer level in order to be stable • Hydrogen, & Helium need only two electrons to complete their outer energy level. (Li, Be, and B try to become like He)

9. Octet Rule • Chemical bonds tend to form so that by gaining, or losing, or sharing electrons, each atom has an octet (8) electrons in the outer energy level. • s2 p6 • Coordinate bonds are an exception to the octet rule. • Expanded valence bonding with electrons in the d sublevel can occur with elements bonding to F, O, and Cl.

10. Oxidation Numbers • Atoms in a pure elemental statehave an oxidation number of zero. O2 P4 S8 • Compounds have an oxidation number of zero. CO2 H2O • Atoms tend to gain or lose electrons to become stable or inert like the noble gases that they are closest to. The noble gases are inert or unreactive. • The oxidation number tells the number of electrons that each atom must gain or lose to become stable.

11. Oxidation Numbers • Hydrogen usually has an oxidation number of +1 because it loses one 1 electron. Electrons have a negative charge so Hydrogen will form a positive ion. • All Group 1 have oxidation number +1 • Group 2 have oxidation number +2 and lose 2 electrons • Oxygen (and other group 6A elements) gains 2 negative electrons (becomes O-2 ) with and oxidation number of -2.

12. Positive Oxidation Numbers • +1: H, K, Na, Ag, Hg (Mercurrous), Cu (Cuprous), Au, NH4 • +2: Ba, Ca, Co, Mg, Pb, Zn, Hg (Mercurric), Cu (Cuppric), Fe (Ferrous), Mn (Manganous), Sn (Stannous) • +3: Al, Au (auric), As (Arsenous), Cr, Fe (Ferric), P (Phosphorous), Sb (Antimonous), Bi (Bismuthous) • +4: C, Si, Mn (Manganic), Sn (Stannic), Pt, S • +5: As (Arsenic), P (Phosphoric), Sb (Antimonic), Bi (Bismuthic)

13. Negative Oxidation Numbers • -1: F Fluoride, Cl Chloride, Br Bromide, I Iodide • -2: O Oxygide, S Sulfide, • -3: N Nitrogide, P Phosphide • -4: C Carbide

14. Cations and Anions • Positive ions are CA IONS • Negative ions are ANIONS • Cations are positive ions that move toward the cathode (negative terminal) and accept electrons. • Anions are negative ions that move toward the anode (positive terminal) and give up electrons.

15. Electron Affinity • Electron affinity is the energy released when atoms attract electrons. • Atoms generally release energy when they acquire an electron. A + e- => A- + energy • Atoms that must be forced to accept an electron release it immediately. A + e- + energy => A- • Electron affinity decreases down the group and increases across the period generally. • Electron affinity, like other energy, is measured in kilojoules

16. Ionization Energy • Ionization energy is the energy required to remove an outer electron from a neutral atom. • Ionization energy increases across a period. • Ionization energy decrease down a group. • The energy required to move successive electrons (2nd and 3rd ionization energy ) increases.

17. Fluorine Electronegativity e- e - • The most electronegative (most able to attract electrons)are in the upper right corner (F) and the least electronegative are in the lower left corner. • The electronegativity is a scale of 0-4 with Fluorine being 4.0 the strongest and Francium being 0.7 the weakest.

18. Linus Pauling received the Nobel prize for developing the table that showed that pattern for electronegativity increasing across the period and decreasing down the period 2.2 4.0 2.0 2.2 0.7

20. Bond Types Intermolecular Intramolecular Dipole-Dipole Van der Waals London Dispersion • Ionic • Covalent • Polar covalent • Coodinating Covalent

21. Bond Energy DH Mg(s) + Cl2(g) → MgCl2(s)  • Mg(s) → Mg(g)   S = 148 kJ/mol • Mg(g) → Mg2+(g)   IP1 + IP2 = 2187 kJ/mol • Cl2(g) → 2 Cl(g)   D = 244 kJ/mol • 2 Cl(g) → 2 Cl–(g)   –2EA = –2(355) = –710 kJ/mol • Mg2+(g) + 2 Cl–(g) → MgCl2(s)   –Elattice • S + IP1 + IP2 + D – 2 EA –Elatttice – ΔHfo = 0 • Elatttice = S + IP1 + IP2 + D – 2 EA – ΔHfo = 148 + 2187 + 244 – 710 – (–641.3) = 2510 kJ/mol • ΔHfo = –641.3 kJ/mol

22. Ionic Bonds (non-metals & metals) When an atom with a high electro -negativity is near an atom with a low electro- negativity, the atom with the high electro negativity will “steal” electrons to complete its outer energy level. (electro negativity must differ by 1.7 or more) • Electrostatic forces of opposite charges hold ions together. • Ex: Na (0.9) + Cl(3.0) => Na+Cl-

23. Ionic Solids • Ex. 1 mole of ionic crystal compound is formed from gaseous ions calcium fluoride • Ca2+ + F1- + F1- CaF2 • Lattice • High melting & Melting Boiling point. • Form ELECROLYTES (aqueous solutions that carry electric current) when dissolved • Formula unit : simplest whole number ratio of positive and negative ions that will combine so that the charges balance.

24. Ionic Compounds • Ionic compoundsdo not form a single molecule. They form a crystal lattice of charged particles. • They tend to have high melting pointsand do not conduct until they begin to become molten. The formula unit tells the ration of the atoms

25. Lattice Energy • Lattice energy is the energy released when 1 mole of ionic crystal compound is formed from gaseous ions.

26. Covalent bonds • Covalent bonds electrons are shared between nuclei lowering overall potential

27. Covalent Bonds (non-metals) • When the atoms have the same electronegativity or the difference in the electronegativity is very slight (less than 0.5) the atoms tend to share electrons. • These covalent bonds form Non-polar covalent compounds and have low melting points, are generally gases & liquids & do not conduct. F F

28. Molecules • A molecule is a group of atoms held together by a covalent bond. • Diatomic molecules are molecules that contain 2 atoms • The molecular formula gives the relative types and number of atoms in a single molecule of the substance

29. Covalent Bond Character • As two atoms approach the nucleus of the atoms will attract the electrons of the opposite atom. • The electrons will repel each other and the nuclei will repel each other. • This creates a great deal of positive and negative potential energy (attraction is negative) • The attractive force is greater when the atoms are far apart. • As the atoms move together the repelling force increases and the attractive force decreases. • When the atoms are close enough together that the forces balance a bond forms holding them at the distance apart.

30. Coordinate Covalent Bonds • In some covalent bonds, like NH+4, one member supplies both electrons in the sharing. • This is due to energy changes that occur as the atoms approach each other. • Atoms will move to the state that is optimal and has lowest energy. This may require one atom giving both electrons to create a bond.

31. Bond Length & Energy • The distance between bonded atoms when they are at the minimum energy is the bond length. • The energy required to break them apart to form free neutral atoms is called the bond energy. The higher the bond energy the stronger the bond • (Exothermic reactions form more stable compounds with low potential energy)

32. Polar Covalent Bonds • When the difference in the electronegativity is between 0.5 and 1.7 the bond is not yet ionic but it is not an equal sharing of electron. • One atom is stronger in its affinity for electrons. It will hold the electrons a little more than the less electronegative atom creating a DIPOLE or polar molecule. H O H The Hs will be slightly positive and the O slightly negative

33. Examples of Bonds • Bonding with Hydrogen Electronegativity 2.1 • Oxygen 3.0 3.5 – 2.1 = 1.4 polar covalent • Fluorine 4.0 4.0 – 2.1 = 1.9 ionic • Astatine 2.2 2.2 – 2.1 = 0.1 non-polar

34. Polyatomic Ions • Polyatomic Ions are charged groups of atoms that are covalently bonded together. • Polyatomic ions act as a single atom. • NH4+ ammonium OH- hydroxide • NO3- nitrate CO3-2 carbonate • PO43- phosphate C2H3O2- acetate • SO32- sulfate ClO3- chlorate

35. Lewis/Structural Formulas • Electron Dot or Lewis Structures can be used to represent molecules. • Unshared pairs or lone pairs are drawn as dots. • Shared pairs are indicated as dashes or lines between the atoms or dots closer to the atom that is more electro negative. O C O O C O

36. Lewis Structures • Count all valence electrons • Add or substract charge if polyatomic ion • Draw skeletal structure with two electron between each pair of bonded atoms (least electronegative is usually center) • Add remaining electrons to complete outer shells on surrounding atoms , then central atom • If the central atom has less than 8 remove an e- from an outer atom and form another bond with the central atom • If the central atom has more than 8 that should be okay

37. More Lewis

38. Double and Triple Bonds • Sometimes in order to achieve the noble gas configuration it is necessary for atoms to share two pairs of electrons forming a double bond or even 3 pairs or electrons forming a triple bond. • Multiple bonds usually are shorter and have higher bond energy.

39. Bond Summary & Extension Bond name Bond length & strength Longest , least Intermediate, intermediate Shortest, greatest • Single-sigma s • Double -1 sigma s, 1 pi p • Triple- 1 sigma s, 2 pi p

40. Resonance Structures • In Molecules like SO2 the molecule changes between different forms or resonates

41. Electrostatic Repulsion • The VSEPR (Valence-Shell Electron Repulsion) Model gives us a way to predict the shape of molecules. • It is based on the idea that sets of like charges ( valence electrons) surrounding an atom will repel as far as possible. • Start with the central atom in a molecule • Mutual Repulsion will force two electron clouds to opposite sides or as far apart as they can get from each other.

42. Basic Shapes • Linear F – Be – F 1800 • Trigonal Planar BF3 1200 • Pyramid with triangular base • Central atom has an unshared pair and is bonded to 3 other atoms • Tetrahedral CH4 109.470 • Four equilateral triangles • Central atom surrounded by four other atoms

43. Unshared Pairs • Unshared (lone) pairs repel more than shared pairs and will distort shapes. • Look at Central Atom • Shape can generally be determined from the number of electrons of the central atom

44. VESPR from http://chemconnections.org/VSEPRUniversity of California — 1111 Franklin St., Oakland, CA 94607 1995-2000

45. Other shapes • Angular • Central atom has 2 unshared pairs and is bonded to 2 other atoms • Trigonal Planar • A central atom (usually IIIA) bonded to 3 other atoms • Trigonal Bi-Pyramid • Central atom bonded to 5 atoms Group VA exception Phosphorus • Octahedon • Central atom bonded to 6 atoms Group VI A exception Sulfur

47. Hybridization • Hybridization occurs when orbitals of different sublevels combine to form new “hybrid” orbitals. • Example: Carbon 1s22s22p2 • We write the dot structure C instead of C • It can form 4 equal bonds because the 2sorbitals and 2p orbitals combine and form 4sp3 orbitals of equal energy greater than the 2s but less than the sp3

48. Network Solids • Network solids are covalent crystals whose covalent bonds extend from one molecule to the next. • Diamonds and Quartz are examples of network solids with carbon and silicon dioxide in the center of the continuous pattern. • They have high melting points and are very hard non conductors.

49. Network Solids The units link

50. Metallic Bonds • Metals have one or more loosely attached electrons that can move freely among the metal ions forming strong attractive forces known as Metallic Bonds. (sea of free roaming electrons) • Because of this free moving electron the metals have high melting points, and are good conductors or heat and electricity, are ductile, malleable, and absorb light and reflect or shine.