4.1 Chemical Bonds Making Connections

Transcription

1 4.1 Chemical Bonds Making Connections

2 Lewis Structures and the Octet Rule 1.1 Key Topics Ionic Bonding and Ionic Compounds Covalent Bonds and Covalent Compounds Lewis Formulas, VSEPR and Molecular Shapes

3 Introduction Carbon can exist in numerous forms Graphite Diamonds

4 Introduction How carbon atoms arrange and bond with each other explains how they could be so different (yet still the same carbon)

5 Chemical Bonding Chemical bonds are the forces that hold atoms (in molecules) and ions (in crystals) together Chemical bonds determine the shapes of our molecules

6 Chemical Bonding These bonds and molecular shapes then influence chemical and physical behavior

7 How Enzymes Work Enzymes are proteins that serve very important functions in an organism One of their functions is to serve as sites for making new molecules in the controlled environment of the cell

8 Modern Medicine Many medicines (and by extension, toxins) are molecules that serve as inhibitors to these enzymes Ex. Acetaminophen Acetaminophen is an analgesic and antipyretic drug that acts as an inhibitor to cyclooxygenase (an enzyme) or COX. COX is responsible for the synthesis of prostaglandins, which promote inflammatory effects in a person.

9 Modern Medicine Many medicines (and by extension, toxins) are molecules that serve as inhibitors to these enzymes Ex. Also known as Acetaminophen is an analgesic and antipyretic drug that acts as an inhibitor to cyclooxygenase (an enzyme) or COX. COX is responsible for the synthesis of prostaglandins, which promote inflammatory effects in a person.

10 Unfortunately, taking in too much paracetamol leads to the production of a toxic substance (N-acetyl-p-benzo-quinone imine) and can cause acute liver failure

11 We know from before that electrons occupy spaces in the atom based on energy level Also, noble gases, with filled electron shells throughout, are quite unreactive

12 THEORY: stability is related to electronic structure and the elements become less reactive when their shells alter to become like noble gases

13 For example,

14 In contrast,

15 Lewis Symbols We can visualize these valence electrons using Lewis symbols

16 Ex. What are the Lewis symbols for: Na Cl C N O

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20 When Na + and Cl - forms, these are held together by their opposing charges. This is called an ionic bond. Ionic compounds are held by these bonds and are found as crystal lattices

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22 Octet Rule In chemical reactions, notice that atoms gain (or lose) electrons to form a valence shell of 8 electrons. This is known as the octet rule

23 For metals, They tend to rather lose electrons whereas For nonmetals, They tend to get electrons to gain noble gas configuration

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25 Nomenclature For the cations, Charge (+) is the same as group number (besides transition elements of course) Name stays as is (ex. Na + - sodium ion) For the anions, Charge (-) is group number minus 8 Name changes to element suffixed with ide (ex. Cl - - chloride ion) Note that this applies only to ionic compounds

26 Ex. NaF sodium fluoride MgO magnesium oxide NaBr KI H 2 S Al 2 O 3

27 Many transition metals can exhibit more than one charge (we formally call these oxidation states). These are denoted by roman numerals Fe 2+ iron (II) ion Fe 3+ iron (III) Cu 2+ copper (II) Cu + copper (I)

28 Ex. KCl

29 Ex. KBr

30 Ex. MgO

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32 Covalent Compounds While metals and nonmetals usually combine to form ionic compounds, nonmetallic elements tend to share electrons rather than lose them A pair of shared electrons = covalent bond A covalent bond

33 Another example:

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35 Nomenclature For binary compounds, we use prefixes to indicate number of atoms

36 Nomenclature For binary covalent compounds, they have 2 names: First name = prefix + 1 st element (Note: If the first element has only one atom, the prefix mono- is dropped.) Second name = prefix + second element + suffix ide.

37 Ex. CCl 4 CO 2 PCl 5 XeF 6

38 Bond Polarity We consider electronegativity when we look at a molecule s polarity

39 When the atoms in a molecule have different electronegativities, the bond is actually drawn to the atom with higher EN. The bond exhibits a separation of charge and is called a polar covalent bond

40 We can represent bond polarity on a Lewis structure like the ff.

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42 Nomenclature We could actually use EN to between two bonded atoms to determine what bond exists between them Δ EN Type of Bond < 0.5 Nonpolar covalent Between 0.5 and 2.0 Greater than 2.0 Polar covalent Ionic

43 Polyatomic Ions Polyatomic ions are groups of covalently bonded atoms with a charge Ex. Ammonium, NH 4 +

44 Polyatomic Ions

45 Nomenclature For those with polyatomic ions, simply name the ions in order Ex. (NH 4 ) 2 SO 4 ammonium sulfate

46 Making Lewis Structures 1. Count valence electrons. 2. Sketch a skeletal structure. 3. Place electrons as lone pairs around outer atoms to fulfill the octet rule. 4. Subtract the electrons used so far from the total number of valence electrons. Place any remaining electrons around the central atom. 5. If the central atom lacks an octet, move one or more lone pairs from an outer atom to a double or triple bond to complete an octet.

47 Making Lewis Structures Ex. Ammonia or NH 3

48 Making Lewis Structures Ex. Methanol (CH 3 OH) Skeletal structure: Final:

49 Making Lewis Structures

50 Making Lewis Structures

51 Free Radicals These are basically molecules that have an odd electron without a pair (unpaired electron). These are insanely reactive molecules! (usually only exist for a bit before reacting away) Ex. NO NO 2 ClO 2

52 Free Radicals They are NOT the same as polyatomic ions! Ex. Carbonate ion (follows octet rule) Nitrogen monoxide (does not follow octet)

53 Resonance structures Speaking of the carbonate ion, we can depict it s Lewis structure like this: b However, if we write down the bonding and lone pairs differently among the 3 oxygens, all these Lewis structures are valid: a c

54 Resonance structures No, you re not experiencing double (or triple) vision, rather some molecules exhibit something called resonance This happens to molecules that have adjacent double and single bonds esp. when there is an electronegative atom involved The electrons get delocalized; all these combined make the Lewis structure for carbonate

55 Resonance structures How do you know which one is the best /stable? We look at the formal charges of each atom in the molecule To get the formal charge: FC = # of valence e - - (# of unshared valence e # bonding e - )

56 Resonance structures For example, ozone is quite a pollutant but at the same time responsible for filtering the UV in the atmosphere. It s two resonance structures are: b b a I c a II c FC a = [6 6 ½ (2)] = -1 FC b = [6 2 ½ (6)] = +1 FC c = [6 4 ½ (4)] = 0 FC a = [6 4 ½ (4)] = 0 FC b = [6 2 ½ (6)] = +1 FC c = [6 6 ½ (2)] = -1

57 Resonance structures So, which one?! 1. Smaller formal charges = like 2. Like formal charges on adjacent atoms = dislike 3. A more negative formal charge on electronegative atom Next: cyanate ion, NCO -

58 Molecular Shapes

59 Molecular Shapes The VSEPR Theory Also known as the Valence Shell Electron Pair Repulsion Theory Gist: Electron pairs arrange themselves around a central atom in a way that minimizes repulsion

60 VSEPR Theory The rules: 1. Start with the Lewis structure 2. Identify all of the electron domains (either lone pair, single bond, double bond, etc.) connected to a central atom 3. In 3D space, spread them as far apart as possible (tetrahedral, triangular, linear, etc.) 4. Final molecular geometry is determined by the combination of lone pairs and bonded atoms

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62 Molecular Shapes Other molecules, especially for those with really large central atoms (n=3 or higher valence shell) have interesting geometries

63 Molecular Shapes Ex. PCl 5

64 Molecular Shapes Ex. XeF 4

65 Shapes and Properties: Polarity One property influenced by molecular shape is the polarity of our molecules. In order for a molecule to be polar: 1. It should have polar bonds (recall: bonding and electronegativity) 2. The bonds should be arranged in such a way that charge separation exists (geometry)

66 Notice the charge distribution. While the C-H bond is very slightly polar (recall: electronegativity), since the charges are equally distributed they cancel each other out. Methane (CH 4 ) is a non-polar molecule.

67 Ammonia on the other hand, has a charge distribution where nitrogen, the electronegative atom, is the slightly negative end (notice that the negative end has the lone pair; the hydrogens are the positive end). Ammonia is a polar molecule.

68 Water is even more polar! (notice the charge distribution; oxygen is very electronegative) Think tug of war between the atoms over their electrons :D

69 Polar or nonpolar? 1. Ethanol (CH 3 CH 2 OH) 2. Carbon tetrafluoride (CF 4 ) 3. Acetic Acid (CH 3 COOH) 4. Phosphorus pentachloride (PCl 5 )