Chapter 1. The Basics Bonding and Molecular Structure. Table of Contents. 1. Life & the Chemistry of Carbon Compounds

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1 hapter 1 The Basics Bonding and Molecular Structure reated by Professor William Tam & Dr. Phillis hang Table of ontents 1. Life & the hemistry of arbon ompounds 2. Atomic Structure 3. hemical Bonds: The ctet Rule 4. ow to Write Lewis Structures 5. Formal harges and ow to alculate Them 6. Isomers 7. ow to Write and Interpret Structural Formulas 8. Resonance Theory 9. Quantum Mechanics & Atomic Structure 10. Atomic rbitals and Electron onfiguration 11. Molecular rbitals 12. The Structure of Methane and Ethane: sp3 ybridization 13. The Structure of Ethene (Ethylene): sp2 ybridization 14. The Structure of Ethyne (Acetylene): sp ybridization 15. A Summary of Important oncepts That ome from Quantum Mechanics 16. ow to Predict Molecular Geometry: The Valence Shell Electron Pair Repulsion Model 17. Applications of Basic Principles 1. Life & the hemistry of arbon ompounds rganic chemistry is the chemistry of compounds that contain the element carbon. arbon, atomic number 6, is a second-row element If a compound does not contain the element carbon, it is said to be inorganic 1

2 arbon compounds are central to the structure of living organisms and therefore to the existence of life on Earth. We exist because of carbon compounds Although carbon is the principal element in organic compounds, most also contain hydrogen, and many contain nitrogen, oxygen, phosphorus, sulfur, chlorine, or other compounds There are two important reasons why carbon is the element that nature has chosen for living organisms: arbon atoms can form strong bonds to other carbon atoms to form rings and chains of carbon atoms arbon atoms can also form strong bonds to elements such as hydrogen, nitrogen, oxygen, and sulfur Because of these bond-forming properties, carbon can be the basis for the huge diversity of compounds necessary for the emergence of living organisms 2. Atomic Structure ompounds made up of elements combined in different proportions Elements made up of atoms Atoms positively charged nucleus containing protons and neutrons with a surrounding cloud of negatively charged electrons Each element is distinguished by its atomic number (Z) Atomic number = number of protons in nucleus 2A. Isotopes Although all the nuclei of all atoms of the same element will have the same number of protons, some atoms of the same element may have different masses because they have different numbers of neutrons. Such atoms are called isotopes 2

3 Examples (1) (6 protons 6 neutrons) (6 protons 7 neutrons) (6 protons 8 neutrons) (2) ydrogen (1 proton 0 neutrons) Deuterium (1 proton 1 neutron) Tritium (1 proton 2 neutrons) 2B. Valence Electrons Electrons that surround the nucleus exist in shells of increasing energy and at increasing distances from the nucleus. The most important shell, called the valence shell, is the outermost shell because the electrons of this shell are the ones that an atom uses in making chemical bonds with other atoms to form compounds The number of electrons in the valence shell (called valence electrons) is equal to the group number of the atom e.g. arbon is in group IVA arbon has 4 valence electrons e.g. Nitrogen is in group VA Nitrogen has 5 valence electrons e.g. alogens are in group VIIA F, l, Br, I all have 7 valence electrons 3. hemical Bonds: The ctet Rule Ionic (or electrovalent) bonds are formed by the transfer of one or more electrons from one atom to another to create ions ovalent bonds result when atoms share electrons 3

4 ctet Rule In forming compounds, they gain, lose, or share electrons to give a stable electron configuration characterized by 8 valence electrons When the octet rule is satisfied for, N, and F, they have an electron configuration analogous to the noble gas Ne Recall: electron configuration of noble (inert) gas # of e - s in outer shell [1s 2 ] 2 Ne 1s 2 [2s 2 2p 6 ] 8 Ar 1s 2 2s 2 2p 6 [3s 2 3p 6 ] 8 3A. Ionic Bonds Atoms may gain or lose electrons and form charged particles called ions An ionic bond is an attractive force between oppositely charged ions Electronegativity (EN) The intrinsic ability of an atom to attract the shared electrons in a covalent bond Electronegativities are based on an arbitrary scale, with F the most electronegative (EN = 4.0) and s the least (EN = 0.7) 4

5 element (EN) Li (1.0) Na (0.9) K (0.8) Rb (0.8) s (0.7) Be (1.6) Mg (1.2) (2.1).. B (2.0)... (2.5) Si (1.8) N (3.0) P (2.1).... (3.5) S (2.5) F (4.0) l (3.0) Br (2.8) I (2.5) Na 1s 2 2s 2 2p 6 3s 1 give 1 e - to l 1s 2 2s 2 2p 6 3s 2 3p 5 (1 e - in outermost shell) (7 e - in outermost shell) ionic bonding Na + l 1s 2 2s 2 2p 6 1s 2 2s 2 2p 6 3s 2 3p B. ovalent Bonds & Lewis Structures ovalent bonds form by sharing of electrons between atoms of similar electronegativities to achieve the configuration of a noble gas Molecules are composed of atoms joined exclusively or predominantly by covalent bonds Example : l.. l: : : [Ne] 3s 2 3p 5 [Ne] 3s 2 3p 5 : l l: covalent bonding : : : : : : 5

6 eterolysis A B omolysis A B heterolytic bond cleavage homolytic bond cleavage A + B ions A + B radicals Ions, themselves, may contain covalent bonds. onsider, as an example, the ammonium ion N (ammonia) (3 bonds on N) + N (ammonium cation) (4 bonds on N with a positive charge on N) 4. ow to Write Lewis Structures Lewis structures show the connections between atoms in a molecule or ion using only the valence electrons of the atoms involved For main group elements, the number of valence electrons a neutral atom brings to a Lewis structure is the same as its group number in the periodic table If the structure we are drawing is a negative ion (an anion), we add one electron for each negative charge to the original count of valence electrons. If the structure is a positive ion (a cation), we subtract one electron for each positive charge In drawing Lewis structures we try to give each atom the electron configuration of a noble gas 6

7 Examples (1) Lewis structure of 3 Br Total number of all valence electrons: Br x = 14 8 valence electrons Br Br: : : remaining 6 valence electrons (2) Lewis structure of methylamine ( 5 N) N Total number of all valence electrons: N 2 valence electrons left x = valence electrons : N 7

8 4A. Exceptions to the ctet Rule Examples Elements in the 2 nd row in the periodic table usually obey the ctet Rule (Li, Be, B,, N,, F) since they have one 2s and three 2p orbitals available for bonding Elements in the 3 rd row in the periodic table have d orbitals that can be used for bonding and may not obey the ctet Rule l l l P l l (Pl 5 ) F F F Si F F F 2- (SiF 6 ) 2- l l P l l l l l P l Phosphorus pentachloride m.p. = 179 o l l Some highly reactive molecules or ions have atoms with fewer than eight electrons in their outer shell F F B F 5. Formal harges and ow to alculate Them Formal charge = number of valence electrons 1/2 number of shared electrons number of unshared electrons or F = Z - S /2 - U where F is the formal charge, Z is the group number of the element, S equals the number of shared electrons, and U is the number of unshared electrons 8

9 Examples (1) The Ammonium ion (N 4+ ) Formal charge of N: = 5 8/2 0 = +1 N Recall: F = Z - S /2 - U group number of N number of shared electrons number of unshared electrons Formal charge of : = 1 2/2 0 = 0 group number of number of shared electrons number of unshared electrons harge on ion = 4 x 0 +1 = +1 The arithmetic sum of all the formal charges in a molecule or ion will equal the overall charge on the molecule or ion Neutral Nitrogen Forms Three Bonds Nitrogen has one lone pair. If nitrogen does not form three bonds, it is charged. 9

10 Neutral xygen Forms Two Bonds Lewis Structures xygen has two lone pairs. If oxygen does not form two bonds, it is charged. (2) The Nitrate ion (N 3 ) Formal charge of : = 6 2/2 6 = -1 group number of number of shared electrons number of unshared electrons Formal charge of : = 6 4/2 4 = 0 group number of number of shared electrons number of unshared electrons Formal charge of N: = 5 8/2 0 = +1 group number of N number of shared electrons number of unshared electrons harge on ion = 2 x (-1) = -1 10

11 (3) Water ( 2 ) The sum of the formal charges on each atom making up a molecule must be zero 5A. A Summary of Formal harges Formal charge of = 6 4/2 4 = 0 Formal charge of = 1 2/2 0 = 0 harge on molecule = x 0 = 0 Resonance structures resonance hybrid 11

12 6. Isomers Isomers: different compounds that have the same molecular formula onstitutional isomers: isomers that have the same molecular formula but different connectivity their atoms are connected in a different order Examples Molecular Formula l Butane l 1-hloropropane onstitutional Isomers and and 2-Methylpropane l 2-hloropropane Examples Molecular Formula 2 6 onstitutional Isomers and 3 3 Ethanol Methoxymethane 7. ow to Write and Interpret Structural Formulas and 3 Butanoic acid Methyl propanoate 12

13 7A. More About Dash Structural Formulas Atoms joined by single bonds can rotate relatively free with respect to one another Equivalent dash formulas for propyl alcohol 7B. ondensed Structural Formulas 7. Bond-Line Formulas 13

14 7D. Three-Dimensional Formulas A dashed wedge ( ) represents a bond that projects behind the plane of the paper A solid wedge ( ) represents a bond that projects out of the plane of the paper An ordinary line ( ) represents a bond that lies in the plane of the paper 7D. Three-Dimensional Formulas Br R R Ethane R Br Bromomethane Br etc. etc. Br l Examples of bond-line formulas that include three-dimensional representations N 2 Br 8. Resonance Theory In a Lewis structure, we draw a welldefined location for the electrons in a molecule. In many molecules and ions (especially those containing bonds), more than one equivalent Lewis structure can be drawn which represent the same molecule An example involving trigonal planar geometry An example involving linear geometry 14

15 We can write three different but equivalent structures, becomes 2 Structures 1 3, although not identical on paper, are equivalent; all of its carbon oxygen bonds are of equal length becomes 3 Resonance structures, then, are not real structures for the actual molecule or ion; they exist only on paper It is also important to distinguish between resonance and an equilibrium In an equilibrium between two or more species, it is quite correct to think of different structures and moving (or fluctuating) atoms, but not in the case of resonance (as in the carbonate ion). ere the atoms do not move, and the structures exist only on paper. An equilibrium is indicated by and resonance by 15

16 8A. ow to Write Resonance Structures We are only allowed to move electrons in writing resonance structures These are resonance structures This is not a proper resonance structure of 1 and 2 because a hydrogen atom has been moved All of the structures must be proper Lewis structures This is not a proper resonance structure of methanol (hypervalent carbon having 5 bonds!!) 16

17 The more stable a structure is (when taken by itself), the greater is its contribution to the hybrid 8B. The Use of urved Arrows: ow to Write Resonance Structures urved arrows show the direction of electron flow in a reaction mechanism point from the source of an electron pair to the atom receiving the pair always show the flow of electrons from a site of higher electron density to a site of lower electron density never show the movement of atoms. Atoms are assumed to follow the flow of the electrons Examples NT Examples (1) Benzene N 3 + NT N 3 + (2) arboxylate ion (R - ) R (3) zone ( 3 ) 17

18 8. ow to Decide When ne Resonance Structure ontributes More to the ybrid The more covalent bonds a structure has, the more stable it is harge separation decreases stability Resonance structures for formaldehyde Four covalent bonds more stable less stable Three covalent bonds Structures in which all the atoms have a complete valence shell of electrons (i.e., the noble gas structure) are more stable 9. Quantum Mechanics & Atomic Structure Wave mechanics & quantum mechanics Each wave function ( ) corresponds to a different energy state for an electron Each energy state is a sublevel where one or two electrons can reside 18

19 Erwin Schrödinger Nobel Prize in 1933 Found the probability of finding an electron in an atom, like flies to a c andle. 10. Atomic rbitals and Electron onfiguration rbitals 3py 2py 3pz 2pz 1s 2s 2px 3d 3s 3px 10A. Electron onfigurations The relative energies of atomic orbitals in the 1 st & 2 nd principal shells are as follows: Electrons in 1s orbitals have the lowest energy because they are closest to the positive nucleus Electrons in 2s orbitals are next lowest in energy Electrons of the three 2p orbitals have equal but higher energy than the 2s orbital rbitals of equal energy (such as the three 2p orbitals) are called degenerate orbitals 19

20 Aufbau principle rbitals are filled so that those of lowest energy are filled first Pauli exclusion principle A maximum of two electrons may be placed in each orbital but only when the spins of the electrons are paired und s rule When we come to orbitals of equal energy (degenerate orbitals) such as the three p orbitals, we add one electron to each with their spins unpaired until each of the degenerate orbitals contains one electron. (This allows the electrons, which repel each other, to be farther apart.) Then we begin adding a second electron to each degenerate orbital so that the spins are paired 11. Molecular rbitals 20

21 The Bonding in 2 We cannot simultaneously know the position and momentum of an electron An atomic orbital represents the region of space where one or two electrons of an isolated atom are likely to be found A molecular orbital (M) represents the region of space where one or two electrons of a molecule are likely to be found Waves an Reinforce Each ther; Waves an ancel Each ther A bonding molecular orbital ( molec ) results when two orbitals of the same phase overlap 21

22 An antibonding molecular orbital ( * molec ) results when two orbitals of opposite phase overlap Atomic rbitals ombine to Form Molecular rbitals Side-to-Side verlap of In-Phase p rbitals Forms a π Bond 22

23 The p rbital of the More Electronegative Atom ontributes More to the Bonding M 12. The Structure of Methane and Ethane: sp 3 ybridization Ground state of acarbonatom 1s 2s 2p x 2p y 2p z ybridization sp 3 covalent bond (line bond structure) sp 3 hybridized carbon 23

24 ybridization sp 3 (Lewis structure) The arbon in Methane is sp o Tetrahedral structure arbon with (3-D stucture) 4 bonds arbon is tetrahedral. The tetrahedral bond angle is B. The Structure of Ethane 24

25 13. The Structure of Ethene (Ethylene): sp 2 ybridization sp 2 ~120 o ~120 o 1 bond + 1 bond sp 2 hybridized carbon ~120 o sp 2 3-Dimensional View -bond ( -orbitals overlap) Planar structure arbon with (3 + 1 ) bonds 25

26 An sp 2 -hybridized carbon atom A model for the bonding molecular orbitals of ethene formed from two sp 2 -hybridized carbon atoms and four hydrogen atoms 26

27 13A. Restricted Rotation and the Double Bond There is a large energy barrier to rotation associated with groups joined by a double bond ~264 kj/mol (strength of the bond) To compare: rotation of groups joined by - single bonds ~13-26 kj/mol 13B. is -Trans Isomerism Stereochemistry of double bonds R identical to R 3 identical to i Pr i Pr 3 27

28 3 different from 3 (trans) Restricted rotation of = 3 3 (cis) e.g. is-trans System o Useful for 1,2 disubstituted alkenes (1) l Br trans -1-Bromo- 2-chloroethane vs. l Br cis-1-bromo- 2-chloroethane (2) trans -3-exene vs. cis -3-exene 14. The Structure of Ethyne (Acetylene): sp ybridization sp 1 bond + 2 bond (3) Br Br trans -1,3- Dibromopropene vs. Br Br cis -1,3- Dibromopropene 180 o sp 2 hybridized carbon Linear structure arbon with (2 + 2 ) bonds 28

29 sp orbital 50% s character, 50% p character sp 2 orbital 33% s character, 66% p character sp 3 orbital 25% s character, 75% p character 29

30 14A. Bond Lengths of Ethyne, Ethene, & Ethane The carbon carbon triple bond of ethyne is shorter than the carbon carbon double bond of ethene, which in turn is shorter than the carbon carbon single bond of ethane Reasons: The greater the s orbital character in one or both atoms, the shorter is the bond. This is because s orbitals are spherical and have more electron density closer to the nucleus than do p orbitals The greater the p orbital character in one or both atoms, the longer is the bond. This is because p orbitals are lobe-shaped with electron density extending away from the nucleus 16. ow to Predict Molecular Geometry: The Valence Shell Electron Pair Repulsion Model Valence shell electron pair repulsion (VSEPR) model: 30

31 16B. Ammonia 16. Water A tetrahedral arrangement of the electron pairs explains the trigonal pyramidal arrangement of the four atoms. The bond angles are 107 (not ) because the nonbonding pair occupies more space than the bonding pairs 16D. Boron Trifluoride 16E. Beryllium ydride 31

32 16F. arbon Dioxide 17. Applications of Basic Principles pposite harges Attract Like harges Repel Nature Tends toward States of Lower Potential Energy END F APTER 1 rbital verlap Stabilizes Molecules 32