Chemical Bonding

1. Chemical Bonding

2. Background Review • Nuclear Charge – This is the total charge present on the nucleus (always positive) It is equal to the number of protons present. • Electrons – In a neutral species, the # of electrons is equal to the number of protons • Ions – species where # of protons is not equal to the # of electrons. (All ions are formed by adding or removing electrons.) • Valence Electrons – All the electrons in the outermost energy level for a given atom or ion. These are involved in chemical bonding. For A families the group number is equal to the number of valence electrons.

3. Lewis Dot Diagrams (Electron Dot Diagrams) • These are used to represent the valence electrons. • Examples – using the elements in period 2

4. TERMS Bond capacity – Is equal to the number of bond pairs. This tells us how many other atoms a particular atom can bond to at one time.

5. Chemical Bonding • Bond – any force that holds atoms (or ions) together. • Compounds – Two or more elements chemically combined • Stable chemical bonds will form if one or both of the following conditions are met. A) The bonded atoms have less energy than the individual atoms. B) Atoms gain or lose electrons to get a configuration like an inert gas.

6. Two models are used to explain chemical bonding: (a) Ionic Bond – this model results from the complete transfer of one or more electrons from one atom to another. (b) Covalent Bond – in this model, the electrons are shared between the bonded atoms. Note: Real bonds have characteristics of both models…We will explain bonding in terms of one or the other. • Electronegativity – tells us how strongly a given atom attracts electrons. The higher the number, the stronger the attraction.

7. Part I The Ionic Bond

8. In this model, electrons are completely transferred from one atom to another. Predicting when to use this model? Two rules allow us to predict this bond type: (A) If the electronegativity difference is 1.7 or greater…..we will use the ionic model. (B) If the elements involved are from group IA and IIA (groups 1 & 2)matched with (bonded to) elements from groups VIA and VIIA (groups 16 & 17), the bond is usually ionic. This results in 2 oppositely charged ions.

9. Ionic compounds can be represented by a chemical formula known as an empirical formula. The empirical formula is based on measurement and observation. It tells us the ratio of positive ions to negative ions in the compound sample. Example: NaCl(s) There is one positive sodium ion (Na+) for each negative chloride ion (Cl-). There is no such thing as a molecule of NaCl(s) . Empirical Formula

10. Lewis Diagrams for Ionic Compounds

11. Sheet 20/8…8

12. REVIEW IONIZATION ENERGY This is the amount of energy required to remove an electron(s) from an atom or ion. (See formation of a positive ion..) IN ANY FAMILY As atomic number increases – ionization energy decreases This is a function of size. The outermost electrons are further and further from the nucleus and are easier to remove. IN ANY PERIOD As atomic number increases – ionization energy increases This is due to increased nuclear attraction on electrons in the same energy level.

13. 3 Step Reaction • The formation of an ionic compound can be represented with 3 steps. • Example: NaCl (aq) • The first step is ENDOTHERMIC. Energy must be supplied in order to remove the electron from the sodium atom. That is: Na + 496 kJ → Na+ + e- • Step 2 is EXOTHERMIC. Energy is released as the electron is accepted by the chlorine atom. That is: Cl + e- → Cl- + 365 kJ

14. 3. Step 3 is also EXOTHERMIC. Energy is again released as the ions move together and take up positions in the forming crystal. That is: Na+ + Cl-→ NaCl + 773 kJ

15. Adding the reactions to get a net ionic equation

17. Characteristics of Ionic Compounds • Solids with high melting points and boiling points. • Non conductors in solid state. • Conductors in aqueous solution. • Conductors in molten (melted) state.

18. Ionic Properties (Ionic Character) • Relative ionic character can be determined by comparing electronegativity differences. • The greater the difference – the greater the ionic character. • Example; which of the following has the greatest ionic character? Li2O(s) , KBr(s) , H2O(l), AlCl3 (s)

19. Types of Covalent Bonds

20. Predicting Molecular Shape V S E P R T Valence Shell Electron We use this theory to predict the shape of molecules. Pair Repulsion Theory The shape that we are describing may be the shape around a specific atom or it may be the overall shape of the molecule.

21. *** Shape is always determined by drawing lines from the center of one specific atom to the center of all atoms bonded to it.***

22. Exercise: Bond each element in period II with as many atoms as possible of an element (‘X’). Element X has a bond capacity of one. Group IA Group IIA Group IIIA Group IVA

23. Group IA Group IIA Group IIIA Group IVA Green (1) black (4) Group VA Group VIA Group VIIA Group VIIIA

24. Molecular Polarity • Some molecules have a polar nature. • This means that one part of the molecule, often an end, is positively charged and one part is negatively charged. • This polarity is the result of uneven sharing of electrons – that is the presence of polar bonds. • Examples: H – Cl Polar molecule H – H Non – Polar molecule

25. We use vectors to determine if a molecule is polar or not. • Procedure: • Replace all polar bonds with vectors. (point the arrow toward the atom with the higher elecronegativity) • Starting at a given point, place all the vectors end to end. • Now: If you end up where you started – the molecule is non – polar. If you end up anywhere else – the molecule is polar. Examples:

26. Types of Forces • Intermolecular forces are attractions between molecules. • Intramolecular forces are attractions within the molecule. (Covalent Bond)

27. Intermolecular attractions • These are forces that exist between covalently bonded molecules. • Example: from one water molecule to another • We will look at three of these intermolecular forces. • In effect, these are a weak, a medium and a stronger force.

28. I. Van der Waals (London Dispersion) • These attractive forces are present between the molecules of ALL samples of matter. • They result from the attraction between the electrons of one molecule and the nuclei of another. • Example: • RULE: the greater the number of electrons, the greater the Van der Waals forces and therefore the higher the melting point.

30. II. Dipole – Dipole • These forces exist, in addition to Van der Waals. • Generally, they are a little stronger. • They consist of the attraction between the positive end of one polar molecule and the negative end of another. • Example:

31. III. Hydrogen Bonding • This is a special type of dipole – dipole. • Generally these are stronger forms. • This will ALWAYS involve molecules that have a Hydrogen atom covalently bonded to a Nitrogen, N, Oxygen, O, or a Fluorine, F. • Hydrogen bonds are generally responsible for the unique characteristics of water.

33. Put the following in order of increasing melting points.CH4, NH3, NCl3, BeO, H2O, HBr, CCl4, LiF,

34. NaCl, C3H8, H2S, OF2, PI3, N2, CaCl2, H2, O2

35. C2H6, MgO, HBr, HF, Cu, NCl3, CO2, CaS, C4H10, PBr3

36. Typical Metallic Solid and Its “Sea of Electrons”

38. Trends in atomic size

40. Electron Dot Structures Symbols of atoms with dots to represent the valence-shell electrons 1A 2A 3A 4A 5A 6A 7A 8A H He:      Li Be  B  C  N  O : F : Ne :            Na  Mg Al Si P  S :Cl  :Ar:    

41. Lewis Diagram for Ionic Compounds      Na + Cl  Na[ Cl ]-           

42. Oxidation/Reduction Half Reactions • Oxidation the loss of electrons by metal. • Reduction the gain of electrons by non-metals. • LEO says GER

43. Oxidation • Loss of electrons (LEO) • Na(s) Na+ + 1e-

44. Reduction • Gain of electrons (GER) • Cl2(g)+ 2e-  2 Cl-

45. Balance • The balanced equation must show a balance between the two half reactions of the equation. Na(s) Na+ + 1e- Cl2(g)+ 2e-  2 Cl- () x 2 2Na(s) + Cl2(g) 2NaCl(aq)

46. Forces

47. Van Der Waals Forces • These are the weakest of the intermolecular forces. There are two types: • London Dispersion • Dipole Interactions

48. London Dispersion • Attraction of electrons to the other molecules positive nucleus. • Weakest of all • the greater the number of electrons in the atoms of a molecule the greater are the dispersion forces

49. Dipole Interactions • When polar molecules are attracted to each other • Electrostatic attractions occur between the opposite charged regions of molecular dipoles. • Similar to ionic bonding but much weaker

50. Hydrogen Bonding • Strongest intermolecular forces. • Weak bond of hydrogen atoms to the unshared electron pair of an electronegative atom in a nearby molecule. • Due to the fact that the valence electrons are not shielded from the nucleus.