Electron Dots and VSEPR Theory

1. Electron Dotsand VSEPR Theory (mostly Chapter 9)

2. Metallic Bonding • In metallic bonding the valence electrons are shared between all the atoms in a positive metal crystal.  delocalized “sea” of electrons  metallic bonded materials have good thermal and electrical conduction.

3. Ionic Bonds Occur when the nonmetal takesone or more electrons away from a metal.  The nonmetal becomes a negative ion  metal becomes a positive ion. The atoms are held together by their opposite charges.

4. Ionic Bond Strength Strength of crystal lattice depends on two factors, size and charge transferred. Smaller atoms have stronger ionic bonds. Ex: NaF is stronger than NaCl Atoms transferring more electrons are stronger. Ex: MgCl2 is stronger than NaCl. 2 e- transferred 1 e- transferred

5. Covalent Bond“the bonds of nature” • Shared valence electrons • Complete outer energy levels • Molecule has 2 or more nonmetal atoms covalently bond • Carbohydrates, proteins, fats, DNA, stupendous seven (H2, N2, O2, F2, Cl2, Br2, I2)

6. How do covalent bonds form? p+ & p+  repulsive e- & e-  repulsive distance is too great e- & e-  repulsive p+ & e- attractive repul. = attract Attractive forces balancethe repulsive forces

7. Electronegativity • Electronegativity is the ability of atoms in a molecule to attract electrons to themselves. • On the periodic table, electronegativity increases as you go… • from left to right across a row. • from the bottom to the top of a column.

8. What types of bonds are they? MgO, water, Calcium Carbide, Potassium Oxide, Nitrogen trihydride

9. Electronegativity and Bond Type Find the difference in 2 atoms’ electronegativies to predict bond type… • Ionic Bonds: 1.7 or greater • Polar Covalent Bonds: <1.7 and >0.2 • Pure or Nonpolar Covalent bonds: <0.2

10. Bonding Capacity

11. Electronegativity Table

12. Drawing Lewis Dot Structures • Count the valence electrons. • Predict the location of the atoms • Hydrogen is a terminal atom • The central atom has the smallest electronegativity. • Draw a pair of electrons between the central atom and the surrounding atoms. • Use the remaining electrons to complete the octets of each atom. If there are electrons left over, place them on the central atom. • If the central atom does not have a complete octet then try double or triple bonds. a. If the atom has 1, 2, or 3 valence electrons, it doesn’t require an octet.

13. STEP 1: count the total # of valence e- for all atoms involved in the bonding CCl4 Carbon: 1 carbon with 4 valence electrons (1x4) = 4 CCl4 4+28 =32 Chlorine: 4 chlorine with 7 valence electrons (4x7) = 28

14. STEP 2–place the single atom in the center and other atoms around it evenly spaced Cl CCl4 4+28 =32 e- Cl Cl C Cl

15. STEP 3: place the electrons in pairs between the central atom and each non-central atom CCl4 4+28 =32 -8 =24 Cl Cl C Cl Cl

16. STEP 4: place the remaining electrons around the non-central atom until each has 8 electrons (H atoms have only 2e-) CCl4 4+28 =32 -8 =24 -24 =0 Cl C Cl Cl Cl

17. Step 5: If you run out of electrons before the central atom has an octet, form multiple bonds until it does.Example: HCN Hydrogen- 1 electron Carbon- 4 electrons Nitrogen- 5 electrons TOTAL is 1+4+5 = 10 e- H:C:N .. H:C:N: .. H:C:::N:

18. Drawing Lewis Dot Structures Draw Lewis Dot Structures for: PH3 H2S HCl CCl4 SiH4 CH2Cl2

19. Draw Lewis Dot Structures Cl2 NF3 CS2 BH3 CH4 SCl2 C2H6 BF3

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21. Covalent Bond Strength • Based on proximity (closeness), also called “bond length” • Influenced by atom size and number of shared electrons • Smaller is stronger F2 is stronger than Cl2 is stronger than Br2 F2: 1.43 x 10-10 m single bond O2 1.21 x 10-10 m double bond N2 1.10 x 10-10 m triple bond

22. Bonding Orbitals • When atoms bond together, their valence shell electron orbitals overlap • Overlapping electron orbitals create a bonding orbital  an area with a high probability of finding an electron

23. Types of Bonds • When atoms form a molecule, their orbitals can form different types of bonds: • Sigma Bonds (σ) • Orbitals overlap head-to-head • Form first, there’s only 1 • Pi Bonds (π) • Orbitals overlap side-to-side • Form after sigma bonds Every molecule has one sigma bond, but all subsequent bonds between the same two atoms must have a different way of connecting so they use pi bonds!

24. Multiple Covalent Bonds – Double 6 valence electrons 6 valence electrons 12 valence electrons Octet satisfied More stable and stronger 1 sigma bond 1 pi bond (lines represent bonded pairs of e-)

25. Multiple Covalent Bonds – Triple 5 valence electrons 5 valence electrons 10 valence electrons Octet satisfied More stable and stronger 1 sigma bond 2 pi bonds

26. Molecular Shapes • The shape of a molecule plays an important role in its reactivity. • Look at bonding and non-bonding electron pairs • You can predict the shape of the molecule!

27. What Determines the Shape of a Molecule? • Electron pairs repel each other. • Assuming electron pairs are placed as far as possible from each other, we can predict the shape of the molecule. Valence Shell Electron Pair Repulsion Theory (VSEPR) “The best arrangement of a given number of electron domains is the one that minimizes the repulsions among them.”

28. Molecular Shape Chart

29. Molecular Polarity • Molecules can be polar and non-polar. • Imagine you are turning over the 3D models on the table. Are they still the same when you flip them over? • If yes, then the molecule is non-polar (symmetrical)

30. Molecular Polarity • Non-bonding electron pair = polar • The free pair pushes the other atoms away • Non-polar molecule has equal pull from the same atoms Non-bonding electron pair

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32. 2sp3 Bonding Orbital Hybridization • Electron orbitals mix to make a new set of bonding orbitals (hybrids) • These have different shapes than regular atomic orbitals • Requires energy but the energy is returned during bond formation Bonds can form here This occurs to allow more bonds! new hybridized orbital

33. Hybrid Orbitals L Consider beryllium: • In its ground state, it would not be able to form bonds because it has no singly-occupied orbitals.

34. Hybrid Orbitals L But by promoting an electron from the 2s to the 2p orbital, it can now form two bonds. This new hybridized orbital is called 2sp J J 2s 2sp orbitals

35. Group 2A elements make sp hybridizedorbitals Group 3A elements make sp2hybridized orbitals

36. Group 4A elements have 4 valence electrons - need 4 bonds to make an octet - they will have sp3 hybridization.

37. Endothermic and Exothermic Reactions • Endothermic Reactions – the energy needed to break the bonds is greater than the energy that is released, energy is used • They feel cool • Exothermic Reactions – the energy needed to break the bonds is less than the energy released, energy given off • They feel warm

38. Questions • Why do some solids dissolve in water but others do not? • Why are some substances gases at room temperature, but others are liquid or solid? • What gives metals the ability to conduct electricity, what makes non-metals brittle? • The answers have to do with … Intermolecular forces

39. Intermolecular forces (also called Van der Waal’s forces) 2 types of attraction in molecules: Intramolecular bonds: (Covalent and ionic) attraction between atoms in a molecule Intermolecular forces (IMF): the attraction betweenmolecules • 1) dipole-dipole • 2) hydrogen bonding • 3) London forces

40. + – H Cl + – + – + – + – Dipole - Dipole attractions • Dipoles: a separation of charge • This happens in both ionic and polar covalent bonds • Oppositely charged dipoles (+δ and –δ) are attracted to each other in a molecule

41. Hydrogen Bonding H-bonding is a special type of dipole - dipole attraction that is very strong (5x stronger) • Happens when N, O, or F are bonded to H • Due to the high electronegativity difference between the H and the other atom • Compounds containing these bonds are important in biological systems (special!)

42. London forces • Named after Fritz London, sometimes called dispersion forces • London forces are due to small dipoles that exist in non-polar molecules • Random movement of electrons can sometimes form temporary dipoles • The resulting tiny dipoles cause attractions between atoms/molecules This is how non-polar molecules can form solids and liquids!

43. London forces Induced dipole: Instantaneous dipole: Sometimes the random arrangement of electrons forms tiny dipoles A random dipole forms in one atom or molecule, inducing a dipole in the other

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45. IMF Strength and Molar Mass • The size of a molecule (molar mass) affects the strength of intermolecular forces (IMFs) • Larger size = stronger forces • Because the large molecule has more area and electrons available for intermolecular attractions such as London Forces • (this is opposite of covalent bond strength) Stronger IMFs Weaker IMFs

46. IMF Strength and Molar Mass • Consider the halogens (group 7A) as an example • F2 and Cl2 are gases, Br2 is liquid, I2 is a solid • Liquids and solids form when IMFs are stronger • Since they are further down the group, the atoms are bigger • Larger mass = stronger IMFs

47. What about these? (such as H2O) Boiling Point and IMFs • Boiling (liquid  gas) occurs when there is enough energy to overcome intermolecular attractions • Boiling point tends to increase down a group, as size of atoms in molecules increases This is because the larger atoms/molecules have stronger IMFs so it takes more energy to break those attractions  higher boiling point!

48. Hydrogen Bonds and Boiling Point • H2O, HF, and NH3 have particularly high boiling points • This is because of hydrogen bonds! • Because they are the strongest IMF, they require more heat energy to break the attraction  higher boiling points

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50. ***Hints for IMF Lab*** • Activity 1, Question 3 asks you to draw the Lewis Dot structures for acetone and ethanol. • Here are their shapes to help you… ethanol acetone