2. Chapter 8 – Covalent Bonding Mr. Samaniego Lawndale High School The unspoken hero: “Covalent Bond”

3. Review of Chapter 7 In Chapter 7, we learned about electrons being transferred (“given up” or “stolen away”) This type of “tug of war” between a METAL and NONMETAL is called an IONIC BOND, which results in a SALT being formed

4. Chapter 8.1 – Molecular Compounds In this chapter, you will learn about another type of bond in which electrons are shared Covalent Bonds are atoms held together by SHARING electrons between NONMETALS

5. Salt versus Molecules A metal cation and nonmetal anion are joined together by an ionic bond called SALT A group of atoms joined together by a covalent bond is called a MOLECULE A Compound is a group of two or more elements bonded together (Ionic or Covalent).

6. Monatomic vs. Diatomic Molecules Most molecules can be monatomic or diatomic Diatomic Molecule is a molecule consisting of two atoms There are 7 diatomic molecules (SUPER 7) – N 2, O 2, F 2, Cl 2, Br 2, I 2, H 2 You can also remember them as: H 2 O 2 F 2 Br 2 I 2 N 2 Cl 2

7. Properties of Molecular Compounds Lower Melting Points than Ionic Compounds (which means that they are weaker than ionic) Liquids or gases at room temperature

8. Molecular Formulas The Molecular Formula is the formula of a molecular compound It shows how many atoms of each element a molecule contains Example H 2 O contains 3 atoms (2 atoms of H, 1 atom of O) C 2 H 6 contains 8 atoms (2 atoms of C, 6 atoms of H)

9. Practice How many atoms total and of each do the following molecular compounds contain? 1. H 2 2. CO 3. CO 2 4. NH 3 5. C 2 H 6 O

10. Practice: True or False 1. All molecular compounds are composed of atoms of two or more elements. 2. All compounds are molecules. 3. Molecular compounds are composed of two or more nonmetals. 4. Atoms in molecular compounds exchange electrons. 5. Molecular compounds have higher melting and boiling points than ionic compounds.

12. Ionic versus Covalent IONICCOVALENT Bonded NameSaltMolecule Bonding TypeTransfer e - Share e - Types of ElementsMetal & Nonmetal Nonmetals Physical StateSolid Solid, Liquid, or Gas Melting Point High (above 300ºC)Low (below 300 ºC) Solubility Dissolves in Water Varies ConductivityGoodPoor

13. Chapter 8.2 – Covalent Bonding Remember that ionic compounds transfer electrons in order to attain a noble gas electron configuration Covalent compounds form by sharing electrons to attain a noble gas electron configuration Regardless of the type of bond, the Octet Rule still must be obeyed (8 valence electrons)

14. Single Covalent Bond A Single Covalent Bond consists of two atoms held together by sharing 1 pair of electrons (2 e - )

15. Electron Dot Structure

16. Shared versus Unshared Electrons A Shared Pair is a pair of valence electrons that is shared between atoms An Unshared Pair is a pair of valence electrons that is not shared between atoms

17. Practice Lewis Dot Structures Chemical Formula # of Valence Electrons Single Line Bond Structure # of Remaining Electrons Lewis Dot Structure Octet Check All Atoms=8 Hydrogen=2 F2F2 H2OH2O NH 3 CH 4

18. Double Covalent Bonds A Double Covalent Bond is a bond that involves 2 shared pairs of electrons (4 e - ) Sometimes atoms attain noble gas configuration by sharing 2 or 3 pairs of electrons

19. Triple Covalent Bond A Triple Covalent Bond is a bond that involves 3 shared pairs of electrons (6 e - )

20. Covalent Bonds

21. Practice Lewis Dot Structure Chemical Formula # of Valence Electrons Single Line Bond Structure # of Remaining Electrons Lewis Dot Structure Octet Check All Atoms=8 Hydrogen=2 O2O2 CO 2 N2N2 HCN

22. Bond Dissociation Energy Bond Dissociation Energy is the energy required to break a bond between two atoms A large bond dissociation energy corresponds to a strong bond which makes it unreactive Carbon has strong bonds, which makes carbon compounds stable and unreactive

23. Chapter 8.3 - Bonding Theories A Molecular Orbital is an orbital that applies to the entire molecule, instead of just one atom So far, the orbitals we have been discussing are atomic orbitals (s, p, d, f) for each atom When two atoms combine, their atomic orbitals overlap and they make molecular orbitals

24. Molecular Orbitals Just as atomic orbitals belong to a particular atom, a molecular orbital belongs to molecules as a whole Each orbital is filled with 2 electrons A Bonding Orbital is an orbital that can be occupied by two electrons of a covalent bond (it’s the space in between the two atoms) There are 2 types of bonding orbitals: sigma and pi

25. Sigma Bond (  ) A Sigma Bond is when 2 atomic orbitals combine to form a molecular orbital that is symmetrical around the axis S orbitals overlapping P orbitals overlapping end-to-end

26. Pi Bond (  ) Pi bonding electrons are likely to be found in a sausage-shape above and below the axis Pi bonds are weaker than sigma bonds because they overlap less Pi bonds are weaker than sigma bonds because they overlap less P orbitals overlapping side-by-side

27. VSEPR Theory VSEPR Theory predicts the 3D shape of molecules According to VSEPR, the repulsion of electrons causes the shape of the molecule to adjust so that the electrons are far apart

28. A Few VSEPR Shapes

29. Nine possible molecular shapes

30. VSEPR Theory Unshared pairs of electrons are very important in predicting the shapes of molecules Unshared pairs of electrons are very important in predicting the shapes of molecules Each bond (single, double, or triple) or unshared pair is considered a steric number Each bond (single, double, or triple) or unshared pair is considered a steric number VSEPR can only be used with the central atom VSEPR can only be used with the central atom Unshared pairs of electrons are very important in predicting the shapes of molecules Unshared pairs of electrons are very important in predicting the shapes of molecules Use the steric number to predict the molecular geometry Use the steric number to predict the molecular geometry

31. Practice Methane (CH 4 ) – tetrahedral Ammonia (NH 3 ) – pyramidal Water (H 2 O) – bent Carbon Dioxide (CO 2 ) - linear

32. Hybrid Orbitals VSEPR is good at describing the molecular shapes, but not the types of bonds formed In hybridization, several atomic orbitals mix to form hybrid orbitals In hybridization, several atomic orbitals mix to form hybrid orbitals Orbital hybridization provides information about both molecular bonding and molecular shape

33. Bond Hybridization Hybridization Involving Single Bonds – sp 3 orbital Ethane (C 2 H 6 ) Hybridization Involving Double Bonds – sp 2 orbital Ethene (C 2 H 4 ) Hybridization Involving Triple Bonds – sp orbital Ethyne (C 2 H 2 )

34. Chapter 8.4 – Polar Bonds and Molecules There are two types of covalent bonds Nonpolar Covalent Bonds (share equally) Polar Covalent Bonds (share unequally)

35. Polar Covalent A Polar Covalent Bond is unequal sharing of electrons between two atoms (HCl) In a polar covalent bond, one atom typically has a negative charge, and the other atom has a positive charge

36. Nonpolar Covalent Bond A Nonpolar Covalent Bond is equal sharing of electrons between two atoms (Cl 2, N 2, O 2 )

38. Classification of Bonds You can determine the type of bond between two atoms by calculating the difference in electronegativity values between the elements Type of BondElectronegativity Difference Nonpolar Covalent 0  0.4 Polar Covalent 0.5  1.9 Ionic 2.0  4.0

39. Practice What type of bond is HCl? (H = 2.1, Cl = 3.1) Your Turn To Practice N(3.0) and H(2.1) N(3.0) and H(2.1) H(2.1) and H(2.1) H(2.1) and H(2.1) Ca(1.0) and Cl(3.0) Ca(1.0) and Cl(3.0) Al(1.5) and Cl(3.0) Al(1.5) and Cl(3.0) Mg(1.2) and O(3.5) Mg(1.2) and O(3.5) H(2.1) and F(4.0) H(2.1) and F(4.0) Difference = 3.1 – 2.1 = 1.0 Therefore it is polar covalent bond.

40. Dipole No bond is purely ionic or covalent … they have a little bit of both characters When there is unequal sharing of electrons a dipole exists Dipole is a molecule that has two poles or regions with opposite charges Dipole is a molecule that has two poles or regions with opposite charges A dipole is represented by a dipole arrow pointing towards the more negative end A dipole is represented by a dipole arrow pointing towards the more negative end

41. Practice Drawing Dipoles P- Br P = 2.1 Br = 2.8 P –Br  +  - Practice H(2.1) – S(2.5) F(4.0) - C(2.5) C(2.5) - Si(1.8) N(3.0) – O(3.5)

42. Attractions Between Molecules Intermolecular attractions are weaker than ionic, covalent, and metallic bonds Besides ionic, metallic, and covalent bonds, there are also attractions between molecules There are 2 main types of attractions between molecules: Van der Waals and Hydrogen

43. Van der Waals Forces Van der Waals forces consists of the two weak attractions between molecules 1. dipole interactions – polar molecules attracted to one another 2. dispersion forces – caused by the motion of electrons (weakest of all forces)

44. Hydrogen Bond Hydrogen Bonds are forces where a hydrogen atom is weakly attracted to an unshared electron pair of another atom

45. Hydrogen Bond This other atom may be in the same molecule or in a nearby molecule, but always has to include hydrogen This other atom may be in the same molecule or in a nearby molecule, but always has to include hydrogen Hydrogen Bonds have about 5% of the strength of an average covalent bond Hydrogen Bonds have about 5% of the strength of an average covalent bond Hydrogen Bond is the strongest of all intermolecular forces Hydrogen Bond is the strongest of all intermolecular forces

46. Intermolecular Attractions A Network Solid contains atoms that are all covalently bonded to each other A Network Solid contains atoms that are all covalently bonded to each other A few solids that consist of molecules do not melt until the temperature reaches 1000ºC or higher called network solids (Example: diamond, silicon carbide) A few solids that consist of molecules do not melt until the temperature reaches 1000ºC or higher called network solids (Example: diamond, silicon carbide) Melting a network solid would require breaking bonds throughout the solid (which is difficult to do)

47. Classwork Chapter 8 Assessment Page 247 #’s 39-41, 43-46, 51, 53, 54, 57-59, 61, 65, 68, 83, 85, 86, 89, 96, 99, 100