1. Unit 9 Bonding 1

2. The attractive forces between atoms leads to chemical bonds that result in chemical compounds. CH 4 methane gas molecule 2

3. Why do atoms form bonds? Atoms form bonds due to the need to have the most stable configuration for its electrons. Atoms lose, gain, or share valence electrons in order to achieve a lower energy state (stable). 3

4. How atoms bond with each other depends on: ee e I want an electron e Electronegativity Ionization Energy # Valence Electrons 4

5. Metallic bonding http://www.launc.tased.edu.au/online/sciences/PhysSci/pschem/metals/Met als.htm 5

6. Metallic bonds Metallic bonding is the strong attraction between closely packed positive metal ions and a 'sea' of delocalized electrons. http://www.bbc.co.uk/schools/gcsebit esize What are the properties of metals? How does metallic bonding affect the properties of metals? Watch This 6

7. What property of metals is illustrated above? Metallic bonds 7

8. Ionic Bonding an electrostatic force Electrostatic refers to the attraction between opposite charges – Stronger than metallic bonds because of the opposite charges 8

9. Ionic bond – Formed by the transfer of electrons between a metal and a nonmetal – # of e - lost by metal = # of e - gained by nonmetal 9

10. Charge The number of electrons that need to be lost or gained by an atom so it has the same electronic configuration as a noble gas. 10

11. Watch This 11

12. Charge based on copy cat principle Sulfur wants to be like Argon (1s 2 2s 2 2p 6 3s 2 3p 6 ) Sulfur Atom 1s 2 2s 2 2p 6 3s 2 3p 4 16 Protons 16(+) 16 Electrons 16 (-) 0 No charge Sulfur has 6 valence e - and will gain 2 more to complete an octet. The result is a: Sulfur Anion 1s 2 2s 2 2p 6 3s 2 3p 6 16 Protons 16 (+) 18 Electrons 18 (-) 2 - Charge 12

13. Charge of Ions Related to the number of electrons that are lost or gained when an atom becomes an ion; called OXIDATION NUMBER Group2 Group 1 Group 13 Group 14 Group 15 Group 16 Group 17 Group 18 +1 +2 +3 +4+4 -3 -2-2 0 13

14. Ionic Bonds Strong electrostatic force (positive-negative) in ionic compounds makes a strong ionic bond. How does the strong ionic bond affect the properties of ionic compounds? Formula unit – smallest unit of an ionic compound; lowest whole number ratio of ions represented in an ionic compound 14

15. The greater the difference between eN values of 2 atoms the more ionic the bond will be. 0.8- 4.0 = 3.2 very ionic 15

16. Ionic Bond http://www.chm.bris.ac.uk/pt/harvey/gcse/ioni c.htm Na + Cl - 16

17. Writing Lewis Dot Structures Element symbol represents the kernel (core) of the atom (nucleus and inner e-) Dots represent the valence e - www.meta-synthesis.com 17

18. Recall that metals tend to lose e - while nonmetals tend to gain electrons Writing Lewis Dot Structures - Ionic hyperphysics.phy-astr.gsu.edu Ionic bonds 18

19. Recall that metals tend to lose e - while nonmetals tend to gain electrons Writing Lewis Dot Structures – Ionic Bonds chemistry58.wikispaces.com Ionic bonds 19

20. Properties of ionic compounds high melting points and boiling points hard solids good conductors – in aqueous solutions and when molten have a crystal lattice structure Ions are here 20

21. Really, we don’t hate you. 21

22. Type of Compound Elements involve in bonding (metal/non metal) Valence electrons are… Melting /Boiling point Electrical conductivity Other properties Metallic Ionic Covalent Type of Compound Elements involve in bonding (metal/non metal) Valence electrons are… Melting /Boiling point Electrical conductivity Other properties MetallicMetal-metaldelocalizedhighConductor Malleable, ductile, shiny IonicMetal - nonmetal Lost/gainedhighConducts in solutions or molten Brittle, solid at room temperature CovalentNonmetal - nonmetal sharedlowNon- conductor Mostly liquid or gas at room temperature Bonding determines some physical properties 22

23. What is a Covalent Bond? Covalent Bond –formed when two nonmetals share pairs of valence electrons in order to obtain the electron configuration of a noble gas Molecule - formed when two or more atoms bond covalently. (A molecule is to a covalent bond as a formula unit is to an ionic bond.) 23

24. How covalent atoms bond 24

25. Diatomic Molecules HOFBrINCl Share electrons when they bond together 25

26. Polyatomic Ions covalently bonded group of atoms, with a charge Watch this 26 They are listed on your STAAR chart. You will not have a bad time.

27. Properties of Covalent Molecules Can exist as gases, liquids, or solids depending on molecular mass and polarity Usually have lower MP and BP than ionic compounds of the same mass Do not usually dissociate (break apart into ions) in water Do not conduct electricity 27

28. How to draw Lewis dot structures for covalent molecules: Write everything 1.Write the formula for the compound. 2.Count the total number of valence electrons. 3.Predict the location of the atoms: a)If there is only 1 atom of an element, it is the central atom. b)If carbon is present, it is ALWAYS the central atom. c)The least electronegative atom is generally the central atom. d)Hydrogen is NEVER the central atom. 4.Place one electron PAIR between the central atom and each ligand (side atom) to “hook” the atoms together. 5. Dot the remaining electrons in pairs around the compound to complete the octet. Start with the ligands. 6.Check that each atom has an octet. (H only needs a pair, not an octet.) 7.Watch ThisWatch This 28

29. Lewis Structures for Molecules Draw the Lewis dot structure for these molecules: – Hydrogen + Bromine (HBr) – Carbon + Chlorine (CCl 4 ) 29

30. Bonding e - Pairs Lone Pairs (nonbonding electrons) Covalent bonds Writing Lewis Dot Structures - Covalent Bonds 30

31. Exceptions to the octet rule: Molecules that have an odd # of valence electrons; ex. NO 2 has 17 total valence electrons and can’t form an exact # of pairs Molecules with fewer than 8 electrons present; ex. BH 3 where B only has and only needs 6 electrons Molecules with an expanded octet; ex. PCl 5 where P forms 5 bonds and SF 6 where S forms 6 bonds 31

32. Number of bonds Single Bonds - when one pair of e- is shared between atoms Double bond – when atoms share 2 pairs of valence electrons; ex. O 2 Triple bond – when atoms share 3 pairs of valence electrons; ex. N 2 32

33. Describing bonds Sigma bond - the first bond between 2 atoms – A single bond is a sigma bond. Pi bond - the second bond between 2 atoms – A double bond consists of a sigma bond and a pi bond. – A triple bond consists of a sigma bond and two pi bonds. 33

34. Carbon can form single, double and triple bonds with itself. 34

35. The shape of a molecule plays a very important role in determining its properties. Why are molecular shapes important? Properties such as smell, taste, and proper targeting (of drugs) are all the result of molecular shape. 35

36. VSEPR Theory also called electron geometry Electron groups around the central atom will be most stable when they are as far apart as possible. We call this valence shell electron pair repulsion theory. – Because electrons are negatively charged, they should be most stable when they are separated as much as possible. The resulting geometric arrangement will allow us to predict the shapes and bond angles in the molecule. 36

37. TO DETERMINE VSEPR SHAPE also known as electron geometry write everything Use VSEPR (valence shell electron pair repulsion) rules: 1) Draw the Lewis dot structure for the molecule 2) Identify the central atom 3) Count the number of electron groups around the central atom. 4) Look up the VSEPR shape on the chart. **shapes with no lone pairs are symmetrical **shapes with lone pairs are assymmetrical 37

38. ONO There are three electron groups on N: Three lone pair One single bond One double bond Electron Groups Each lone pair of electrons constitutes one electron group on a central atom. Each bond constitutes one electron group on a central atom, regardless of whether it is single, double, or triple. 38

39. Two Electron Groups: Linear Electron Geometry When there are two electron groups around the central atom, they will occupy positions on opposite sides of the central atom. This results in the electron groups taking a linear shape. 39

40. Linear Geometry 40

41. Three Electron Groups: Trigonal Planar Electron Geometry When there are three electron groups around the central atom, they will occupy positions in the shape of a triangle around the central atom. This results in the electron groups taking a trigonal planar shape. 41

42. Trigonal Planar Geometry 42

43. Four Electron Groups: Tetrahedral Electron Geometry When there are four electron groups around the central atom, they will occupy positions in the shape of a tetrahedron around the central atom. This results in the electron groups taking a tetrahedral shape. 43

44. Tetrahedral Geometry 44

45. Molecular Geometry The actual geometry of the molecule may be different from the VSEPR shape. Lone pairs repel bonded atoms which distorts the expected shape. 45

46. Bond Angle Distortion from Lone Pairs Electron Geometry Tetrahedral Tetrahedral Tetrahedral MolecularGeometry Tetrahedral Trigonal Pyramidal Bent 46

47. Watch This 47

48. Predicting the Shapes around Central Atoms write everything 1.Draw the Lewis structure. 2.Determine the number of electron groups around the central atom. 3.Classify each electron group as a bonding or lone pair, and count each type. – Remember, multiple bonds count as one group. 4.Look it up 48

49. Practice: Determine the shape. 1. NF 3 2. SiCl 4 3. H 2 O 49

50. Types of BondsTypes of Bonds: Nonpolar covalent equal sharing of electrons between atoms; occurs between the atoms in a diatomic molecule (HOFBrINCl) and between C and H; ex. CH 4 50

51. Polar Covalent unequal sharing of electrons between atoms; occurs between two nonmetals or a nonmetal and a metalloid; ex. H 2 O Electrons 51

52. Ionic complete transfer of electrons; occurs between m/nm, m/PAI, PAI/nm or PAI/PAI; ex. NaCl 52

53. THIS IS A CONTINUUM. IT DESCRIBES THE “IONIC CHARACTER” OF THE BOND. Bond type Non-Polar Covalent Polar Covalent Ionic NPC PC I Difference in electronegativity values Distance between atoms on the periodic table Small medium big 53

54. Practice: What type of bond exists in each of the following? 1. HCl 2. CaO 3. H 2 O 4. Br 2 54

55. Water is a POLAR molecule The more electronegative atom will have a slight negative charge, the area around the least electronegative atom will have a slight positive charge. 55

56. Symmetric molecules tend to be nonpolar Asymmetric molecules with polar bonds are polar 56 Symmetric means there are no lone pair around the central atom.