A noble gas electronic configuration is highly desirable energetically

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1 Chemistry 261 Fall 2018 Lecture 1 Emphasis 1. Organic chemistry is based on the chemistry of carbon, Z = 6, 1s 2 2s 2 2p 2. As you will recall, the 1s 2 electrons are buried, leaving the 2s 2 and 2p 2 electrons as valence (outer) electrons. However, owing to hybridization (section 1.12) carbon forms 4 bonds. 2. Atoms differ principally in the number of protons in the nucleus, and achieve electrical neutrality by matching the number of protons by the same number of electrons distributed according to the aufbau principal. A noble gas electronic configuration is highly desirable energetically Variations in the number of neutrons may occur and give rise to isotopes. These may be stable ( 12 C, 13 C) or radioactive ( 11 C, 14 C) depending on n:p, but will react and behave in most instances just like one another 3. Ionic bonds (section 1.2(b)) are opposite charge attractions between oppositely charged species 1

2 4. Covalent bonds (section 1.2(c)) occur by the sharing of valence electrons, generally to obtain a noble gas electron configuration. The shared electrons are shown as a line and it is incredibly important for you to see the line between atoms in a Lewis structure as a pair of valence electrons 5. If a single valence electron of one non-metal atom enters into a bond with a single valence electron of another non-metal atom, a neutral molecule will result. It may or may not be polar, but it will not act as an electrolyte. Recall that when 2 atomic orbitals overlap, they give rise to 2 molecular orbitals, one of which places the electron density between the contributing atoms the bonding molecular orbital and one which decreases the electron density between the atoms the antibonding molecular orbital (section 1.11) While molecular orbital considerations are important, a great deal of day to day organic chemistry is conducted using the [more simplified] Lewis structure Lewis Structures (section 1.4), an Abbreviated How To 1 : A. Draw the Lewis dot structure for each element, showing the valence electrons B. Using those atoms that can form more than 1 bond, pair unpaired electrons to form a single bonds between the atoms, bearing in mind geometric restrictions (e.g. 3 membered rings are difficult to form) 1 See recommendations on page 8 as well 2

3 C. If the number of single electrons remaining = the number of hydrogens and halogens (which only require 1 bond to achieve a noble gas configuration) add them, forming bonds to complete the structure D. If there are more unpaired electrons then hydrogen or halogens, pair them to generate double or triple bonds until # of single electrons remaining = # of H and halogens. Finish as in (a) Example: Draw the Lewis structure for C2H4O 6. When considering Lewis structures, it is important to bear in mind the possibility of charged covalent species by the determination of formal charge. By comparing the valence depiction of an isolated atom (group number or table 3.5 above) to the bonding configuration of an atom under a particular circumstance, assigning formal charge is a straightforward matter. All you have to do is remember to count one of the 2 electrons in a bond (they are shared after all) and count all non-bonded electrons for the atom in question. As an example, oxide has no covalent bonds and 4 non-bonded (or lone) pairs and thus bears a -2 charge (oxygen is group 6), hydroxide has one covalent bond and 3 non-bonded pairs, thus a -1 charge, water does not bear a formal unit charge since there are 2 covalent bonds and 2 lone pairs, and the hydronium ion has a +1 charge given 3 covalent bonds and only a single lone pair remaining 3

4 7. To expand the concept of building neutral Lewis structures by bringing single valence electrons together and then investigating for the possibility of having a charged species, Brønsted-Lowry acid-base theory is highly useful. After all, an acid base reaction between 2 neutral molecules does give rise to a salt, no? Recall in a Brønsted-Lowry acid base reaction the equivalent of a proton, H +, is being transferred from acid to base. Neutral molecules based on the period 2 non-metals (C, N, O, F) have covalent bonds equal to the number of single valence electrons based on Lewis dot depictions. If one of the period 2 hydrides acts like an acid it will have one fewer bond, one additional pair of non-bonded electrons and a negative charge (having lost H + ). If one of the period 2 hydrides acts as a base (excluding C) it will have an additional covalent bond, one fewer pair of non-bonded electrons, and a positive charge (having gained H + ). Which of these forms are most likely may be deduced by considering electronegativities, as they all have approximately the same Van der Waal s radii. +H + +H + +H + +H + CH4 NH3 H2O HF -H + -H + -H + -H + Notice that it is difficult to place a positive charge on an electronegative atom; similarly it is difficult to place a negative charge on an electropositive atom. It is much easier to place a negative charge on an electronegative atom (fluoride ion) and a positive charge on an electropositive atom (sodium cation) 4

5 Question: An extraordinarily useful reducing agent that may be used in an aqueous environment is sodium borohydride (NaBH4). Draw the Lewis structure for sodium borohydride beginning with the electron dot depictions, noting that it may be viewed as an addition product between BH3 and NaH (NaH may be conveniently viewed as a [nearly] ionic compound between Na and H) 8. Exceptions to the octet rule (section 1.2(c)) occur when there are not enough electrons available to fulfill the octet rule (BH3) or there are available d orbitals which electrons may be promoted into (H2SO4, H3PO4). We can still use the same method for constructing Lewis structures, only you have to promote one of the paired electrons into a vacant orbital to maximize bonding potential. Recall sulfuric acid is H2SO4, and phosphoric acid is H3PO4. Spread the S or P electrons out first, add in the oxygens, form double bonds where necessary, and add in hydrogens, making certain there are no single ( radical ) electrons remaining Problem: Draw the Lewis structure for sulfuric acid and phosphoric acid, beginning with the electron dot symbols Problem: Draw the Lewis structure for the important polar, non-protic solvent dimethyl sulfoxide, (CH3)2SO (hint: spread out only one of sulfur s non-bonded pairs of electrons) 5

6 9. Electronegativity is the tendency of an atom to attract an electron in a covalent bond; the most usual Pauling scale runs from 0.8 to 4.0 EN 0.5 is considered polar though this is a continual scale o Hydrocarbon EN = 0.4. These act as the standard for non-polar organic solvents (oil and water do not mix) o The common organic solvent dichloromethane (DCM, methylene chloride has EN = 0.5 and will not dissolve in H2O ( EN = 1.4) even though considered weakly polar The concept of polarity and its influence on physical properties and reactivity is one of the most important in all of organic chemistry you absolutely must develop a sense for which atoms are polarizing to an organic skeletal structure and in what direction o The concept of polarity is so important we will consider variants of it and their influence such as polarizability of large atom electron clouds and influence on properties that results 6

Essential Organic Chemistry. Chapter 1

Essential Organic Chemistry. Chapter 1 Essential Organic Chemistry Paula Yurkanis Bruice Chapter 1 Electronic Structure and Covalent Bonding Periodic Table of the Elements 1.1 The Structure of an Atom Atoms have an internal structure consisting