1. Glencoe, Chapters 7 & 8

2. The Process What is a chemical bond? It is where 2 or more atoms are linked to one another through a chemical change… a new substance is formed. What part of the atom is directly involved in chemical bonding? the electrons Natural systems, like a reaction system, tend to seek a decrease in energy. Remember that when new bonds form, energy is released from the system to the surroundings.

3. Determining Bond Types – pp.215-217, 241-242, 263-264 There are two basic types of chemical bonds, covalent and ionic. Nearly all bonds (except those where the two bonded atoms are identical) have characteristics of both types, but usually one type is more predominant.

4. Determining Bond Types – pp.215-217, 241-242, 263-264 Which type is predominant is determined by the electronegativities of the two atoms involved in bonding. [Recall the definition of electronegativity – a measure of the attraction an atom has for a shared pair of electrons.] Each element has an electronegativity value. (See p.265.)

5. Determining Bond Types – pp.215-217, 241-242, 263-264 The calculated difference between the electronegativities of the two atoms involved in bonding, called electronegativity difference, is the best predictor of predominant bond type between the atoms. This continuum is generally used to determine predominant bond type after calculation of electronegativity difference: nonpolar covalent polar covalentionic 0 0.2 1.67 4.0

6. Determining Bond Types – pp.215-217, 241-242, 263-264 An electronegativity difference of 1.67 is the 50/50 breakpoint for covalent/ionic character. [Recall the electronegativity trends we learned for the periodic table. Generally, electronegativity increases across a period and up a group.] Let’s predict bond types. Ex: carbon dioxide calcium chloride hydroiodic acid

7. Covalent Bonds – pp.241-247, 264-265 Covalent bonds are formed as a result of the sharing of a pair of electrons between two atoms. Ex: the formation of diatomic fluorine from 2 fluorine atoms

8. Covalent Bonds – pp.241-247, 264-265 Electron sharing may be virtually equal between the two atoms (e-diff of 0.2 or less) or unequal (e-diff of greater than 0.2 but less than 1.67). In unequal sharing, the atom with the greater electronegativity value has a greater attraction for the shared pair and pulls it closer to it. This inequality creates partial charges. The atom with the greater electronegativity value takes on a partial negative charge; the atom with the lesser electronegativity value takes on a partial positive charge.

9. Covalent Bonds – pp.241-247, 264-265 α+α- Ex: hydrochloric acidH – Cl e-diff = 0.96 2.20 3.16 PC bonds The partial charge differential creates a phenomenon known as a dipole. A bond with a dipole is called a polar bond, thus the polar covalent classification. A covalent bond without a dipole (equal sharing) is called a nonpolar covalent bond.

10. Ionic Bonds – pp.215-217 Ionic bonds are formed as a result of the transfer of one or more electrons from one atom to another. In the broad sense, ionic bonding is the electrostatic attraction between oppositely charged atoms, ions. Ex: the formation of sodium chloride

11. Ionic Bonds – pp.215-217 When electronegativity difference between atoms in a reaction system exceeds 1.67, ions will likely form. Electrons are considered to be transferred, rather than shared, between bonding atoms. The atom with the greater electronegativity value takes on a negative charge, gains the electron(s); the atom with the lesser electronegativity value takes on a positive charge, loses the electron(s).

12. Covalent vs. Ionic Compounds In general, when compared, covalent compounds tend to have lower melting points and boiling points than ionic compounds. Covalent compounds and their solutions tend to be electrically non-conductive, referred to as non- electrolytes. When placed in water, if soluble, covalent compounds tend to simply dissolve, not dissociate. Ex: sucrose (table sugar) C 12 H 22 O 11 (s) H2O C 12 H 22 O 11 (aq)

13. Covalent vs. Ionic Compounds Ionic compounds, when in the molten state or in solution, tend to be electrically conductive, referred to as electrolytes. If soluble, ionic compounds not only dissolve, but also tend to dissociate (break into ions). Ex: sodium chloride (table salt) NaCl (s) H2O Na +1 (aq) + Cl -1 (aq) What seems to be necessary for the conduction of electricity? the presence of mobile charged particles.

14. Multi-bonds – pp.245-246 Covalent bonds may be single, double, or triple. Single bonds result from the sharing of one pair of electrons between two atoms (a sigma bond). Double bonds result from the sharing of two pairs of electrons between two atoms (1 sigma, 1 pi) and triple bonds from the sharing of three pairs (1 sigma, 2 pi). (Examples will be forthcoming.) Triple bonds are stronger and have a shorter bond length than double bonds, and double bonds more than single.

15. Bond Energies – pp.246-247 Energy is consumed by the system to break existing bonds. When new bonds form, energy is released by the system. Bond energy is normally expressed as the energy, in kJoules, required to break one mole of bonds or that released when one mole of bonds form. Typically, it takes more energy to separate a triple bond, than a double bond, than a single bond. (See p.275, Table.)

16. Lewis Structures – pp.252-258 Since electrons are the part of the atom directly involved in bonding, Lewis structures (electron dot diagrams) are very useful in predicting bonding behaviors, especially with molecular compounds. Recall how to draw Lewis structures for sulfur, phosphorus, chlorine, magnesium SPClMg Remember, the element symbol represents the atom’s nucleus and all electrons but the valence electrons. Valence electrons are represented around the symbol.

17. Lewis Structures – pp.252-258 Compound bonding and geometry can be predicted through the use of Lewis structures, electronegativity differences, electron-pairing behaviors, and the octet rule. Examples: chlorine gas ammonia

18. Lewis Structures – pp.252-258 When depicting geometries, a shared pair of electrons is represented by a bond line rather than dots. Unshared electrons are still pictured, if they impact geometry.

19. Lewis Structures – pp.252-258 For less straight-forward molecules, there is a way to calculate the numbers of electrons that should be involved in bonding. Example: carbon dioxide electron contributions: C4for each atom involved to have O68 electrons, you'd need 3x8=24 O6electrons. You only have 16. So 16 total24-16=8 electrons must be shared! There are only two bond sites, so there are four electrons at each site. Double bonds are present at each site.

20. Lewis Structures – pp.252-258 You try nitrogen gas:

21. Lewis Structures – pp.252-258 BUT REMEMBER, there are exceptions to the octet rule. Some atoms are satisfied with fewer or more than 8 electrons, especially those in the 3-5 original electrons category. Examples: boron of boron tribromide phosphorus of phosphorus pentachloride hydrogen of hydrogen iodide

22. Lewis Structures – pp.252-258 How do you represent an ionic compound through Lewis structures? You show electron transfer and resulting charges. Ex: cesium sulfide

23. Lewis Structures – pp.252-258 Lewis structures can also be used to represent ions. Ex: the sulfite ion This structure also introduces a special kind of covalent bond, the coordinate covalent bond. When both of the shared electrons in a bond are provided by one of the bonding atoms, a coordinate covalent bond is formed.

24. Resonance Structures – p.256-257 When the structure of a molecule cannot be fully represented by just one Lewis structure, the molecule has resonance structures. Resonance structures are various arrangements of electrons within the molecule which result in changes in bond positions. (Usually double or triple bonds are involved in resonance structures.)

25. Resonance Structures – p.256-257 Ex: the nitrate ion electron contributions: N5for each atom involved O6to have 8 electrons, you'd O6need 4x8=32 electrons. O6 You only have 24. So 32- +1 charge24=8 electrons must be 24 totalbonded! There are only three bond sites, charge so one site must have a double bond.

26. Molecular Geometry – pp.259-262 The Valence Shell Electron Pair Repulsion theory (VSEPR theory), in combination with Lewis structures, aids in the prediction of molecular geometries. The VSEPR theory states that bonded and lone pairs of electrons will be arranged about a central atom as far apart as possible to minimize repulsive forces between electron pairs. Electron pair repulsions vary in strengths: unshared-unshared-shared- unshared >shared> shared repulsionrepulsionrepulsion

27. Molecular Geometry – pp.259-262 If there are only two shared electron pairs in the outer level of the central atom, as in beryllium iodide, they will arrange equidistantly (at 180º angles), on opposite sides of the nucleus… creating a linear geometry:linear BeI 2 Be-I ediff = 1.09 PC bonds

28. Molecular Geometry – pp.259-262 If there are only three shared electron pairs in the outer level of the central atom, as in boron trichloride, they will arrange equidistantly spaced (at 120º angles)… creating a trigonal planar geometry:trigonal planar BCl3 B-Cl ediff = 1.12 PC bonds

29. Molecular Geometry – pp.259-262 If there are only four shared electron pairs in the outer level of the central atom, as in methane(CH 4 ), they will arrange equidistantly spaced (at 109.5º angles)… creating a three-dimensional geometry called tetrahedral (See p.260 for better drawing.):better drawing CH4 C-H ediff = 0.35 PC bonds

30. Molecular Geometry – pp.259-262 The presence of unshared pairs of electrons around the central atom creates more complex geometries, due to the variations in electron-pair repulsions.

31. Molecular Geometry – pp.259-262 If there are three shared pairs and one unshared pair of electrons in the outer level of the central atom, as in ammonia, they will arrange with the shared pairs closer to one another and the unshared pair more isolated… creating a three-dimensional geometry called trigonal pyramidal (107.3° bond angles):trigonal pyramidal NH3 N-H ediff = 0.84 PC bonds

32. Molecular Geometry – pp.259-262 If there are two shared pairs and two unshared pairs of electrons in the outer level of the central atom, as in water, they will arrange with the shared pairs closer to one another and the unshared pairs more isolated from them and each other… creating a geometry known as bent or angular (104.5° bond angles):angular H2O H-O ediff= 1.24 PC bonds

33. Intermolecular Forces – pp.228-229, 264-266 How is it that uncharged (neutral) atoms or molecules, like water, can be attracted enough to one another to come together as liquids or solids? There are attractive forces between atoms and molecules known as intermolecular forces. They are not "real" bonds, like ionic or covalent bonds. They are often referred to as weak or secondary forces because even the strongest type has only 1/10 the strength of a true bond. We'll discuss five types of intermolecular forces… the metallic bond, London forces, dipole-dipole interactions, hydrogen bonding, and electrostatic attractions.

34. Intermolecular Forces – pp.228-229, 264-266 The metallic bond (pp.228-229) - is what holds metallic atoms together, as in a piece of solid iron. Since each atom has the same number of electrons in its outer shell, low electronegativity, and no impetus to form ions, the electrons become delocalized. Electrons move freely from one atom to the next, creating a "sea of electrons" within the material. These free flowing electrons give metallic substances their luster (shine), ductility, malleability, and thermal/electrical conductivity.free flowing electrons

35. Intermolecular Forces – pp.228-229, 264-266 {London forces and dipole-dipole interactions are grouped as van der Waals forces, accounted for by Johannes van der Waals} London forces - are the weakest of the intermolecular forces. They occur between nonpolar molecules or uncharged non-metallic atoms (for example, the forces between CO 2 molecules as they solidify or between Ne gas molecules as they liquefy). London forces result from momentary dipoles that form as electrons move about the nucleus (or the molecule) and become concentrated to one side, and are otherwise known as dispersion forces.dispersion forces

36. Intermolecular Forces – pp.228-229, 264-266 Dipole-dipole interactions - occur between polar molecules, those with molecular dipoles. Electronegativity difference tells you whether or not bonds are polar, but it doesn’t tell the whole story with respect to molecule polarity. Geometry and direction of pull must also be considered.molecule polarity

37. Intermolecular Forces – pp.228-229, 264-266 Of course, if no polar bonds are present, a molecule is nonpolar. But, does the presence of polar bonds in a structure always result in a polar molecule? Yes, UNLESS the polar bonds are symmetrically arranged, equal in force, and opposite in direction, such that they cancel out. Recall that the more electronegative element in a bond pulls harder on the shared pair creating partial charges, and a resulting dipole.

38. Intermolecular Forces – pp.228-229, 264-266 Analyze both bond polarity and molecule polarity of the following: Hydrochloric acid Beryllium hydride

39. Intermolecular Forces – pp.228-229, 264-266 Water Methyl chloride (CH 3 Cl)

40. Intermolecular Forces – pp.228-229, 264-266 Hydrogen bonding – is the strongest of the intermolecular forces. When hydrogen is bonded to a very electronegative element, hydrogen is left with a very strong partial positive charge and no other electrons to shield the attractive force of the nucleus. In essence, from a charged particle viewpoint, it is reduced to a proton! It is therefore “attracted” to another molecule’s very electronegative element. Hydrogen bonding Hydrogen bonding is most prevalent in molecules where hydrogen is bonded to oxygen, nitrogen, or fluorine. Hydrogen bonding has 1/10 the strength of a true bond, yet greatly influences physical and chemical properties of substances. Water is an excellent illustration of hydrogen bonding:Water

41. Intermolecular Forces – pp.228-229, 264-266 Electrostatic Attractions – the intermolecular forces between ions in an ionic compound