Chapter 6: Chemical Bonding

1. Chapter 6:Chemical Bonding Use the periodic table to infer the number of valence electrons in an atom and draw its electron dot structure. Be able to explain the types of bonds that atoms can form. List the characteristics of the different types of chemical bonds. Define the vocabulary words. Use electronegativity values to classify a bond

2. Valence Electrons • Electrons in the highest occupied energy level of an element’s atoms • For representative elements, the number of valence electrons is the same as the group number of that element (Page 414) • Shown in electron dot structures 3 5 2 1 6 8 4 7 Right, left, top, bottom (1,2,3,4) Then 12 o’clock and counterclockwise (5,6,7,8) Symbol of the element

3. Valence Electrons (cont’d) • Electrons in the highest occupied energy level of an element’s atoms • Can be figured out using the group numbers in the periodic table. • Ex: The elements of Group 1A (hydrogen, lithium, sodium, etc.) all have a valence number of 1, which means there is 1 electron in the highest occupied energy level. The elements of group 7A (fluorine, chlorine, bromine, etc.) have 7 electrons in the outer energy level. • The valence numbers also tell us the likely oxidation state of that element. More on this later.

4. Oxidation States The oxidation state of an atom is the charge it has when it gains or loses electrons to form it’s most stable electron configuration. Valence Number Oxidation State (charge on the ion) 1 +1 2 +2 3 +3 5 -3 6 -2 7 -1 } CATIONS } ANIONS

5. The Octet Rule • Gilbert Lewis used this to explain why atoms form certain kinds of ions and molecule. • In forming compounds, atoms tend to achieve the electron configuration of a noble gas (8 valence e-) • Recall that each noble gas (except He) has 8 electrons in its highest energy level and a general electron configuration of ns2 np6 • Exceptions: Molecules with an odd number of electrons, more than an octet (PCl5), and less than an octet (very rare) Example: NO2 has seventeen valence electrons [Nitrogen contributes five and each oxygen contributes 6 (2 x 6 =12)]

6. The Octet Rule • An atom’s loss of an electron produces a cation, or positively charged ion. The most common cations are those produced by the loss of valence electrons from the metals, since most of these atoms have 1-3 valence electrons. Let’s look at sodium (a 1A metal) as an example: Na 1s22s22p63s1 Na+ 1s22s22p6 -e-

7. Practice Problems • Write the electron dot structure for each of the following: 1. Na 2. Al 3. N 4. S 5. Kr 6. Chloride ion 7. Oxide ion • Refer to pages 414, 417, and 418 for answers.

8. Practice Problems Please write the oxidation numbers of the following: • Na 11) Po • Al 12) Ga • F 13) Cr • Cl 14) N • Mg • P • Ca • Sb • I • Sc

9. Common Polyatomic Ions Hydroxide: OH- Permanganate: MnO4- Bicarbonate: HCO3- Ammonium: NH4+ Carbonate: CO32- Acetate: C2H3O2- Sulfate: SO42- Hydrogen- Sulfite: SO32- Phosphate: HPO42- Phosphate: PO43- Dichromate: Cr2O72- Nitrate: NO3- Nitrite: NO2- Chlorate: ClO3- Cyanide: CN-

10. Chemical Bonding • Chemical energy & potential energy stored in chemical bonds • Atoms prefer a low energy condition • Atoms that are bonded have less energy than free atoms- more stable. • To combine atoms: energy is absorbed • To break a bond: energy is released (AB)

11. Chemical Bonds • Created when two nuclei simultaneously attract electrons • When electrons are donated or received, creating an ion (anion, cation) • In most elements, only valence electrons enter chemical reactions • Atoms of everyday substances are held together by chemical bonds (water, salt anti-freeze)

12. Types of Chemical Bonds 1) Ionic Bond: chemical bonding that results from the electrical attraction between cations and anions where atoms completely give their electron(s) away

13. Types of Chemical Bonds 2) Covalent Bond: chemical bonding that results from the sharing of electron pairs between two atoms. The electrons are “owned” equally by the two atoms.

14. Relative Forces of Attraction • Ability of a nucleus to hold its valence electrons (Group 7A has a greater ability to hold on to its valence electrons than Group 1A) • Ionization energy: energy required to lose an electron (As atomic number increases down a group, the most loosely bound electrons are more easily removed, so ionization energy decreases. For the most part, it increases along each period.)

15. Electron affinity – tendency to gain an electron (Energy is released) • Electronegativity – measure of the electron attracting power of an atom when it bonds with another atom * Fluorine (4.0) is the highest * Cesium (0.7) is the lowest – least ability to attract bonding electrons and thus the greatest tendency to lose an electron * Noble gases are not assigned electronegativities because these elements do not generally form bonds (inert)

16. The periodic trend of the electronegativities is the same as that of the ionization energies. Thus, as the atomic number increases along a period, the electronegativity increases. As the atomic number increases down a group, the electronegativity decreases. • In general, metals have a low electronegativity and nonmetals have a high electronegativity

17. Electronegativity and Bond Types Covalent Bonds: bonding between elements with an electro- negativity difference of 1.7 or less. Nonpolar-Covalent Bonds: covalent bond in which electrons are shared evenly by the bonded atoms with an electronegativity difference of 0 – 0.3. Polar-Covalent Bonds: covalent bond in which the bonded atoms have unequal attraction of the shared electrons, and have an electronegativity difference of 0.4 – 1.7 Ionic Bonds: bonding due to difference in electric charge of two elements due to loss/gain of electrons. Must have an electro- negativity of 1.8 – 4.0.

18. Electronegativity and Bond Types Water is a polar molecule, because the electrons are not shared evenly by the hydrogen and oxygen. Ionic Bonds: bonds in which electrons are donated from one atom to another and have an electronegativity difference of 1.8 or higher.

19. Electronegativity and Bond Types Using the electronegativity values found on page 161 of your book, predict the types of bonds the following will form. 1) O2 2) NaCl 3) N2 4) Knowing that the electronegativity of sulfur is 2.5, what type of bond will sulfur form with: a) hydrogen b) cesium c) chlorine

20. Electronegativity Electronegativity is a measure of how strongly an element can remove an electron from another element.

21. Ionic (Electrovalent) Bonds • The strongest chemical bond Complete transfer of electron(s) from one element to another • Generally formed when metals combine with nonmetals (Groups 1-2a w/ 5-7a) • Coulombic forces – electrostatic force in which two oppositely charged ions are mutually attracted • Usually occurs when the difference in electronegativities is 1.8 or greater

22. NaCl – Ionic Bond • Draw

23. NaCl – Ionic Bond • Draw : . . : Na Cl :

24. NaCl – Ionic Bond • Draw : . . : Na Cl :

25. NaCl – Ionic Bond • Draw : . . : Na Cl :

26. NaCl – Ionic Bond • Draw : . . : Na Cl :

27. NaCl – Ionic Bond • Draw : : : Na Cl :

28. Writing Ionic Compounds • Beryllium fluoride • Calcium oxide • Scandium sulfide • Aluminum chloride

29. Ionic Solids • Form crystal lattice (orderly, repeating, three-dimensional pattern) • The charges and relative sizes of the ions determines the crystal structure • The number of ions of opposite charge that surround the ion in a crystal is called the coordination number of the ion.

30. Poor conductors of electricity (no free electrons) • High melting point • High boiling point • Brittle and break easily under stress • Liquid or aqueous: good conductors of electricity but ionic bond is dissolved

31. The Normal Arrangement of an Ionic Crystal Opposite charges attract - + - + + - + - - + - + + - + - - + - +

32. Arrangement when Stress is Applied - Adjacent to ions with same charge (repulsion) + + - - + - + + - + - - + - + + - - +

33. Crystal Lattice is Destroyed - + + - - + Crystal melts, vaporizes, or dissolves in water (ions free to move about) Cleavage – splitting along a definite line - + + - + - - + - + + - - +

34. Covalent Bonding • Electrons are shared • One atom does not have enough pull on the electron to take it completely from the other atom • Occurs when electronegativity difference is 1.7 or less • Covalently Bonded Solids: 1. Softness 2. Poor conductor of electricity and heat 3. Low melting point

35. Lewis Structures • Single covalent bond – one shared pair of electrons: • H· + ·H H H or H H • Double covalent bond – two shared pairs of electrons O + O O O or O O • Triple covalent bond – three shared pairs of electrons N + N N N or N N Note: all of these obey the octet rule : : : : : : : . . : : : : : : : : . . : : . . . . : : : : : : : . .

36. Coordinate covalent bond – one atom contributes both bonding electrons • NH3 + H+ [NH4]+ ammonia hydrogen ion ammonium ion H + H N H + H+ H N H H H The structural formula shows an arrow that points from the atom donating the electrons to the atom receiving them. Refer to page 444 : : : : : : :

37. How to Construct Lewis Structures Step 1: Determine the type and # of atoms in molecule CH3I has 1 Carbon, 3 Hydrogens and 1 Iodine Step 2: Write electron dot notation for each type of atom CH· I Step 3: Determine the total # of electrons available in the atoms to be combined. C 1 x 4e- = 4e- I 1 x 7e- = 7e- H 3 x 1e- = 3e- 14 e- . . . . . . . . . . .

38. How to Construct Lewis Structures Step 4: Arrange the atoms to form a skeleton structure for the molecule. Then connect the atoms by electron-pair bonds. H C I Step 5: Add unshared pairs of electrons to each non- metal atom so that each is surrounded by 8. HC I H . . . . . . . . H H . . . . . . . . . . . . . . H

39. Electron Dot Practice: Compounds • H2O 5) CCl2H2 • H2O2 6) NH3 • HCN 7) N2 4) AlF3 8) CO2

40. e- e- e- e- e- e- e- e- e- e- • A single water molecule is a good example of covalent bonding between atoms. The hydrogen atoms “share” their electrons with the larger oxygen atom so that oxygen now has a full outer level with 8 electrons and each hydrogen has a full outer level with 2 electrons. Oxygen has a higher electronegativity than hydrogen, so there is actually an uneven sharing of electrons, resulting in a polar molecule. More on this later. 8p+ 8n0 shared electrons shared electrons 1p+ 1p+

41. Bond dissociation energy: total energy required to break the bond between two covalently bonded atoms (remember that energy is measured in joules or kilojoules) H–H + 435 kJ H + H • Resonance Structures: refers to bonding in molecules or ions that cannot be correctly represented by a single Lewis structure. . .

42. Bond Length vs. Bond Energy There is a correlation between bond length and the amount of potential energy stored in that bond. For example: Bond Bond Length (pm) Bond energy (Kj/mol) C C 154 346 C C 134 612 C C 120 835 C N 147 305 C N 132 615 C N 116 887 N N 145 163 N N 125 418 N N 110 945

43. Molecular orbitals – when two atoms combine and their atomic orbitals overlap • Sigma bond - molecular orbital that is symmetrical along the axis connecting two atomic nuclei In both of these examples, the p orbitals are overlapping and sharing electrons.

44. pi bond – weaker than sigma bond; usually sausage-shaped regions above and below the bond axis (Page 445)

45. H3C – CH3 H3C – CH3 Examples of Sigma and Pi bonds H2C = CH2 – HC CH – –

46. VSEPR Theory (page 200) VSEPR Theory (Valence Shell Electron-Pair Repulsion theory): states that repulsion between the sets of valence-level electrons surrounding an atom causes these sets to be oriented as far away from each other as possible, thus determining the shape of molecules.

47. 1 : O . 2 : . VSEPR Theory So then why is H2O bent, but BeF2 is linear? The answer is the free electron pairs. Oxygen has 2 pairs, beryllium has none. H2O BeF2 Be . .

48. VSEPR Shapes Linear Trigonal-Planer Bent/Angular Tetrahedral Trigonal-Pyramidal Trigonal-Bipyramidal Octahedral (#s 3, 5 and 7 are coordinate covalent bonds!)

49. VSEPR Theory . . These free electron pairs repel each other because they have a negative charge, and so they force those atoms that are covalently bonded to be pushed as far away as possible. . .

50. Hybridization – several atomic orbitals mix to form the same total number of equivalent hybrid orbitals (CH4 – Page 457) *Note: An sp3 orbital is an example of a hybrid.