Unit 8. Covalent Bonding

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1 Unit 8 Covalent Bonding

2 The Ionic Bond When sodium and chlorine atoms combine, the sodium atoms give their electrons to chlorine. Both ions now have stable noble gas electron configurations and the oppositely charged atoms form an ionic bond.

3 The Covalent Bond When 2 hydrogen atoms combine to make H 2, the atoms both share their 1 electron so each has a noble gas configuration like helium.

4 The Chemical Bond Why do atoms sometimes take and sometimes share to be come stable? Electronegative Difference (p.290)

5

6 Bond Type Comparing the elctronegativities will help decide who wins, if anyone, in the tug of war for electrons!

7 Bond Types Chlorine is strong enough to take the electron from sodium, but the hydrogen have equal electronegativities and therefore end up sharing the electrons. Cl - Na + H H

8 Determining the Bond Type Electronegativity difference Bond Type Nonpolar Covalent > Polar Covalent Example H H (0) H Cl (0.9) 2.0 Ionic Na + Cl - (2.1) Molecules are usually in this range

9 Bond Type Open to p.405 and let s try a couple. What is the bond type for NaCl? Na = Cl = = 2.1 It is ionic (big surprise)!

10 Bond Type What is the bond type in a water molecule, H 2 O? Water has two H-O bonds! H = 2.1 O = = 1.4 The H-O bonds are Polar Covalent

11 Covalent Vocabulary Covalent bond sharing of a pair of electrons Molecule electrically neutral unit of a substance that retains the properties of the substance and is held together by two or more covalently bonded atoms.

12 Covalent Vocabulary Nonpolar covalent - equal sharing of electrons. Polar covalent - unequal sharing of electrons.

13 Covalent Vocabulary Single covalent bond - sharing of 1 pair of e - between atoms. Cl Cl Cl Cl Lewis Dot Structures Unshared pair - pair of electrons that is not being shared.

14 Covalent Vocabulary Double covalent bond - sharing of 2 pairs of e - between atoms. O O Triple covalent bond - sharing of 3 pairs of e - between atoms. N N

15 Bond Strength and Length Bond length distance between 2 bonding nuclei. - Single bond > double bond > triple bond Bond dissociation energy (bond strength) energy needed to break a covalent bond. - Sum of the bond dissociation values is the amount of chemical potential energy in a molecule - Bond strength is inversely proportional to bond length Triple bond > Double bond > Single bond

16 How do you draw molecules? H-Cl Space filling-model of HCl Lewis dot of HCl For the simple molecules we ll be using, the key to drawing is to find the number of bonds. For octet structures, an easy way to find the number of bonds is to use the N-A=S method!

17 Drawing Covalent Structures 1. First, find the number of bonds (shared e - pairs) using the following formula: N A = S N = Total number of e - needed for all to have a full octet (Remember while most things need 8, H and He need 2) A = Total number of e - available (For polyatomic ions, it equals A - charge) S = Shared electrons S # bonds = 2

18 Drawing Covalent Structures 2. Draw molecule without bonds. - Draw central atom(s). Usually the element there is only one of, otherwise use elements from the following families as the central atom C N B O (never H or halogen) - Draw outer atoms spaced around central atom(s). For SO 3 O S O O

19 Drawing Covalent Structures 3. Draw a single bond between all atoms. - If there are left over ones, then you will need to have multiple covalent bonds! - Multiple bonds will go between carbons (if there are 2) or between central atom and B, N, or O groups(usually O). SO 3 has 4 bonds! O S O O

20 Drawing Covalent Structures 4. Fill in dots (draw them in pairs) so all atoms have full octets. H-Cl 5. If it s a polyatomic ion, put the whole thing in brackets and write the charge in the upper right hand corner. H N + H H + H

21 Let s do a few together: N (needed e - ) A (available e - ) S (shared e _ ) # Bonds Electron dot NH 3 8(1)+2(3) 14 5(1)+1(3) /2 3 O 2 HCN C 2 H 4

22 Polyatomic Ions Polyatomic ions are covalently bonded atoms that have gained or lost electrons to become stable. Because they are covalent, we can use N-A=S to draw these structures. The only difference is that the gain or loss of electrons has effected how we calculate A - Find A like before, then subtract the charge to find the actual value of A. Draw NO 2 - N = A = S = 8(3) = 24 5(1) + 6(2) = 17 - (-1) = = 6

23 Resonance When one Lewis structure does not correctly represent the molecule you have resonance. This occurs when there is an equal choice for placing multiple bonds. Draw all structures connected by double arrows. N N O O O O

24 Non-Octet Structures When you have a structure that doesn t follow the octet rule, you can t use N-A=S!!!! Quickly decide on the number of bonds, and draw the structure!

25 Non-Octet Structures Non-Octet Structures 1. Less than an octet a. Beryllium Stable with only 4 shared electrons (2 bonds) b. Boron (and sometimes Aluminum) Stable with only 6 shared electrons (3 bonds) 2. Expanded octets (more than an octet) Central atom has more than 4 bonds H Be H PCl 5 and SF 6

26 Non-Octet Structures N (needed e - ) A (available e - ) S (shared e _ ) # Bonds Electron dot BeCl 2 PCl 5 BH 3

27 Molecule Shape Valence Shell Electron Pair Repulsion (VSEPR) Theory - electrostatic repulsion between valence electron pairs surrounding an atom cause these pairs to be oriented as far apart as possible. Two atom molecules are always linear.

28 Molecule Shape Number outside atoms Symmetrical Shape Nonsymetrical Shape Type of molecule 2 3 linear AB 2 Trigonal planar bent Trigonal pyramidal AB 2 E AB 2 E 2 AB 3 AB 3 E 4 tetrahedral AB 4 5 trigonal bypyramidal AB 5 A = number central atom B = number outside atoms E = unshared electron pairs on the central atom 6 octahedral AB 6

29 Hybridization Electrons in an atom are found in atomic orbitals. What happens to these orbitals when atoms share electrons in a molecule? They form hybrid orbitals!

30 Hybridization of CH 4 Carbon electron configuration This carbon needs to share 4 electrons! How does the carbon atom get 4 electrons to share? tetrahedral

31 What does SO 3 do? The sulfur needs to form 3 hybrid orbitals to attach and share with the 3 oxygens! What happens to the unhybridized p orbital? trigonal planar

32 Sigma and Pi Bonds sulfur s unhydridized p orbital oxygen s unhydridized p orbital The bonds that form between the atoms (in the blue) are called Sigma (σ) bonds. All single bonds and 1 of the bonds in a multiple bond are sigma bonds The bond that forms when parallel p orbitals overlap is called a Pi (π) bond. Double bond (1 σ and 1 π) Triple bond (1 σ and 2 π)

33 Shape Total Electron regions Bond Angle Hybridization Example linear sp BeCl 2 CO 2 bent 3 <109.5 sp 2 O 3 bent 4 <109.5 sp 3 H 2 O Trigonal planar sp 2 BF 3 SO 3 Trigonal pyramidal 4 <109.5 sp 3 NH 3 tetrahedral sp 3 CH 4 Trigonal Bipyramidal 5 90 /120 sp 3 d PCl 5 Octohedral 6 90 sp 3 d 2 SF 6

34 Molecule Polarity Just like bonds can be polar or nonpolar, so can molecules. Molecular polarity starts with the bond type. - You can t have a polar molecule unless you have polar bonds! The problem is that polar bonds don t always give you a polar molecule. - Perfect symmetry (symmetrical with all outside atoms of the same element) can neutralize the polarity to give a nonpolar molecule.

35 Molecule Polarity Bond type nonpolar covalent polar covalent polar covalent Perfect Symmetry? yes or no yes no Molecule Type nonpolar nonpolar polar

36 Intermolecular Force Intermolecular Force - forces of attraction between molecules. 1. Dipole-dipole: strong force of attraction between polar molecules. 2. Hydrogen bonding: very strong dipole-dipole caused between polar molecules containing hydrogen bonded to N, O, or F with an unshared pair of electrons. - Examples - HF, NH 3, and H 2 O - Is it strong H-NOF (enough) 3. London dispersion forces: weak force of attraction between nonpolar molecules caused by momentary dipoles of constantly moving electrons.

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