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thanks a lot for sharing this powerpoint. it is actually for my lecture class.

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BONDS in ORGANIC COMPOUNDS Types of bonds The ionic bond The covalent bond Polar covalent bonds Polarity of organic molecules Coordinate bonds Intermolecular forces Effects in organic molecules-Part I

Slide 2: 

2 When the nucleus↔electron attractions (blue arrows) are greater than the nucleus ↔ nucleus and electron–electron repulsions (red arrows), the result is a net attractive force that holds the atoms together to form a molecule.

Why Do Atoms Bond? : 

3 Why Do Atoms Bond? processes are spontaneous if they result in a system with lower potential energy chemical bonds form because they lower the potential energy between the charged particles that compose atoms the potential energy between charged particles is directly proportional to the product of the charges and inversely proportional to the distance between the charges the more negative the potential energy, the more stable the system becomes a chemical bond forms when the potential energy of the bonded atoms is less than the potential energy of the separate atoms

Energy of Interaction : 

Energy of Interaction Closer together; attraction increases. At 74 pm, attractive forces are at a maximum, energy is at a minimum. Closer than 74 pm, repulsion increases. The distance between nuclei at this minimum energy point is called the bond length.

Types of Bonds : 

5 Types of Bonds

Review of Chemical Bonds : 

Review of Chemical Bonds IONIC — complete transfer of 1 or more electrons from one atom to another (one loses, the other gains) forming oppositely charged ions that attract one another COVALENT — some valence electrons shared between atoms METALIC — holds atoms of a metal together Most bonds are somewhere in between ionic and covalent. There are 3 forms of bonding:

Electronegativity Differenceand Bond Type : 

Identical atoms have the same electronegativity and share a bonding electron pair equally. The bond is a nonpolar covalent bond. When electronegativities differ significantly, electron pairs are shared unequally. The electrons are drawn closer to the atom of higher electronegativity; the bond is a polar covalent bond. With still larger differences in electronegativity, electrons may be completely transferred from metal to nonmetal atoms to form ionic bonds. Electronegativity Differenceand Bond Type

Slide 8: 

F2 H—Br H—F Na+ F- Non-polar covalent Polar covalent Ionic Polar bonds are often depicted using colors to show electrostatic potential (blue = positive, red = negative).

Ionic Bonds : 

The ionic compounds contain ionic bonds. Positive cations and the negative anions are attracted to one another and form a pattern called a crystal lattice Therefore, ionic compounds are usually between metals and nonmetals (opposite ends of the periodic table). Ionic Bonds when metals bond to nonmetals, some electrons from the metal atoms are transferred to the nonmetal atoms metals have low ionization energy, relatively easy to lose an electron nonmetals have high electron affinities, relatively good to accept electrons Li+1 •• Li•

Slide 10: 

10 The noble gases are the least reactive group of elements The alkali metals are the most reactive metals and their atoms almost always lose 1 electron when they react The halogens are the most reactive group of nonmetals and in a lot of reactions they gain 1 electron Metals form cations by losing enough electrons to get the same electron configuration as the previous noble gas (metals are reducing agents, they oxidize) Nonmetals form anions by gaining enough electrons to get the same electron configuration as the next noble gas (nonmetals are oxidizing agents, they reduce) The noble gas electron configuration must be very stable

Covalent Bonds : 

11 Covalent Bonds nonmetals have relatively high ionization energies, so it is difficult to remove electrons from them when nonmetals bond together, it is better in terms of potential energy for the atoms to share valence electrons potential energy lowest when the electrons are between the nuclei shared electrons hold the atoms together by attracting nuclei of both atoms

Bond Order and Bond Length : 

Bond order is the number of shared electron pairs in a bond. A single bond has BO = 1, a double bond has BO = 2, etc. Bond length is the distance between the nuclei of two atoms joined by a covalent bond. Bond length depends on the particular atoms in the bond and on the bond order. Bond Order and Bond Length

Polar Covalent Bonds and Electronegativity : 

Electronegativity (EN) is a measure of the ability of an atom to attract its bonding electrons to itself. EN is related to ionization energy and electron affinity. The greater the EN of an atom in a molecule, the more strongly the atom attracts the electrons in a covalent bond. Polar Covalent Bonds and Electronegativity Electronegativity generally increases from left to right within a period, and it generally decreases from the top to the bottom within a group.

Slide 15: 

Electronegativity Scale MAX MIN It would be a good idea to remember the four elements of highest electronegativity: N, O, F, Cl.

Bond Polarity : 

Bond Polarity HCl is POLAR because it has a positive end and a negative end. (difference in electronegativity) Cl has a greater share in bonding electrons than does H. Cl has slight negative charge (-d) and H has slight positive charge (+ d)

Bond Polarity : 

Bond Polarity covalent bonding between unlike atoms results in unequal sharing of the electrons one atom pulls the electrons in the bond closer to its side one end of the bond has larger electron density than the other the result is a polar covalent bond the end with the larger electron density gets a partial negative charge the end that is electron deficient gets a partial positive charge d+ d- HF molecules align in an electric field

Slide 19: 

the larger the difference in electronegativi-ty, the more polar the bond

Slide 20: 

This is why oil and water will not mix! Oil is nonpolar, and water is polar, the two will repel each other. Bond Polarity “Like Dissolves Like”: Polar dissolves Polar Nonpolar dissolves Nonpolar Water molecules in cage around a carbon chain

Electronegativity and Bond Polarity : 

21 Electronegativity and Bond Polarity If difference in electronegativity between bonded atoms is 0, the bond is pure covalent (equal sharing) If difference in electronegativity between bonded atoms is 0.1 to 0.4, the bond is nonpolar covalent If difference in electronegativity between bonded atoms 0.5 to 1.9, the bond is polar covalent If difference in electronegativity between bonded atoms larger than or equal to 2.0, the bond is ionic “100%”

Polar Molecules : 

22 Polar Molecules Entire molecules can be polar if electrons are attracted more strongly to one part of the molecule than to another. Molecules polarity is due to the sum of all individual bond polarities and lone-pair contribution in the molecule. Polarity has a dramatic effect on the physical properties of molecules, particularly on melting points, boiling points, and solubility.

Slide 23: 

23 Dipoles or polarity can be represented by an arrow pointing to the negative end of the molecule with a cross at the positive end resembling a + sign.

Slide 24: 

24 Just because a molecule has polar covalent bonds does not mean that the molecule is polar overall. Carbon dioxide and tetrachloromethane molecules have no net polarity because their symmetrical shapes cause the individual bond polarities to cancel each other out.

Ionic Bonding vs Covalent Bonding : 

25 Ionic Bonding vs Covalent Bonding ionic compounds are soluble in water and have high melting points and boiling points (because the attractions between ions are strong, breaking down the crystal requires a lot of energy) MP generally > 300°C all ionic compounds are solids at room temperature molecular compounds have low melting points and boiling points MP generally < 300°C molecular compounds are found in all 3 states at r.t°. melting and boiling involve breaking the weak attractions between the molecules, but not the bonds between the atoms (the covalent bonds inside the molecules are strong) the polarity of the covalent bonds influences the strength of the intermolecular attractions

Characteristics of Molecular Compounds : 

26 Characteristics of Molecular Compounds The bond type has a dramatic effect on the physical properties of molecules, particularly on the state of matter, melting points, boiling points, conductivity and solubility.

Coordinate Covalent Bonds : 

28 Coordinate Covalent Bonds Coordinate covalent bond: The covalent bond that forms when both electrons are donated by the same atom.

Slide 29: 

29 The ammonium ion, is an example of a species with a coordinate covalent bond. Coordination compounds are an entire class of substances based on the ability of transition metals to form coordinate covalent bonds with nonmetals. Essential metal ions are held in enzyme molecules by coordinate covalent bonds.

Questions : 

Why do some solids dissolve in water but others do not? Why are some substances gases at room temperature, but others are liquid or solid? What gives metals the ability to conduct electricity, what makes non-metals brittle? The answers have to do with … Intermolecular forces Questions

Overview : 

Overview There are 2 types of attraction in molecules: intramolecular bonds & intermolecular forces We have already looked at intramolecular forces that hold together the atoms within a molecule Intermolecular forces (IMF) have to do with the attraction between molecules (≠ from the attraction between atoms in a molecule) IMFs are of six types: 1) ionic, 2) dipole - dipole, 3) H-bonding, 4) London forces, 5) covalent (network solids), 6) metallic Intermolecular forces

Intermolecular Forces : 

Intermolecular Forces The amount of forces between molecules causing one molecule to influence another may be assesed Heats of vaporization/fusion or melting or boiling points give a measure of the strength of attractions present between molecules (the energy required to separate molecules when changing from liquid to gas state) Also the solubility is dependent on intermolecular forces

Slide 33: 

Intermolecular forces pin gigantic molecules like enzymes, proteins, and DNA into the shapes required for biological activity. Myoglobin

Ionic and Polar Covalent Compounds : 

Ionic and Polar Covalent Compounds The force of attraction in ionic compounds is the electrostatic force between ions A relatively strong force Hvap /100 kJ/mol The forces of attraction in polar covalent compounds (that present a separation of charge, i.e. dipoles) are the dipole-dipole attractions Dipole-dipole forces are smaller than ion-ion forces Hvap .20 kJ/mol

Slide 36: 

+ + - - Molecules are electrically neutral overall but organize themselves by attractions of head to tail dipole orientation The greater the EN, the greater the separation of charge in dipoles

H - bonding : 

H - bonding H-bonding is a special type of dipole - dipole attraction that is very strong It occurs when N, O, or F are bonded to H The high EN of NH, OH, and HF bonds cause these to be strong forces (about 5x stronger than normal dipole-dipole forces) They are given a special name (H-bonding) because occur in compounds containing H bonds; They are very important in biological systems and biochemical reactions

Hydrogen Bonding: A special dipole-dipole interaction : 

Hydrogen Bonding: A special dipole-dipole interaction :X-H :X-H - + - + The energy of the H-bond depends on the electronegativity of the X-atom F > O > N , Cl ......... Calculate the EN for HCl and H2O HCl = 2.9-2.1 = 0.8 H2O = 3.5-2.1 = 1.4

Heats of Vaporization : 

non-polar molecule Heats of Vaporization H-bonding H-bonding

London forces : 

London forces Non-polar molecules do not have dipoles like polar molecules. How, then, can non-polar compounds form solids or liquids? London forces are named after Fritz London (but also called van der Waals forces) London forces are due to small dipoles that exist in non-polar molecules Because electrons are moving around in atoms there will be instants when the charge around an atom is not symmetrical The resulting tiny dipoles cause attractions between atoms/molecules

London forces : 

London forces Instantaneous dipoles: Induced dipoles: Eventually electrons are situated so that tiny dipoles form A dipole formed in one atom or molecule, inducing a dipole in the other

Slide 42: 

Hvap increases with increasing of the electrons number

Principles of Solubility : 

Principles of Solubility Solubility is dependent on intermolecular forces “like dissolves like” liquids with similar structures (similar type & magnitude intermolecular forces) will be soluble in each other in all proportions. hexane pentane Both are held together by London Forces When mixed together, there is no significant environment change

Oil Slicks : 

Oil Slicks Non-polar substances have little water solubility Water molecules are held together by H-bonds Non-polar oil molecules are held together by London Forces H-bonds must be broken to dissolve appreciable quantities of non-polar oil or other substances in water Non-Polar & Slightly Polar Substances Most soluble in solvents of low polarity Least soluble in H bonding solvents

Water Solubility of Polar Molecules : 

Water Solubility of Polar Molecules Water will dissolve some polar molecules like short chain alcohols –methyl and ethyl alcohol, that are capable of forming H-bonds Intermolecular forces between these alcohols and water are similar to those forces in pure alcohol and pure water. Solubility decreases as length of carbon chain increases As the chain gets longer, more H-bonds in the water must be broken to make room for the alcohol. Not enough H-bonds can be reformed to compensate

Solid-Liquid : 

Solid-Liquid Solids always have limited solubility in liquids: due to differences in the magnitudes of intermolecular forces in solid vs. liquid state at 25oC a solid has much stronger intermolecular forces than a liquid The closer a solid is to its mp, the better its intermolecular forces will match up with a liquid Typically, solubility increases as the temperature increases Low mp solids tend to exhibit greater solubility than high mp solids


EFFECTS IN ORGANIC MOLECULESTHE INDUCTIVE AND THE RESONANCE EFFECTS The terms "resonance" and "induction" refer to the electronic effects that atoms or functional groups may have within a compound. These effects are defined below and are dependent on the valence, bonding order, position within a structure and electronegativity of atoms, as well as the molecular geometry. All the effects in organic molecules may be “electron-withdrawing” or “electron-donating” effects, and their sign, - or + accordingly assigned.

Slide 50: 

Any atom or group, if attracts electrons more strongly than hydrogen, it is said to have a -I effect (electron-attracting or electron-withdrawing), e.g. NO2, Cl, Br, I, F, COOH OCH3, etc., while if atom or group attracts electrons less stronlgy than hydrogen it is said to have +I effect (electron-donating/repelling or electron-releasing) e.g. CH3, C2H5, Me2CH and Me3C groups. The important atoms or groups which cause -I or +I (inductive) effect are arranged below in the order of decreasing effect. -I (Electron-attracting) groups: +I (Electron-donating) groups: THE INDUCTIVE EFFECT (±I)

Slide 51: 

Inductive effects refer to those electronic effects of an atom or functional group that can contribute through single bonds such as saturated (sp3) carbon atoms to the reactivity of the bond. The contribution of electronegativity, bonding order and position toward induction is as follows: THE INDUCTIVE EFFECT (±I) The inductive effect is simply the deplacement of the electrons of a s bond in agreement with the difference in electronegativity of the participant atoms in that bond. The C─H bond is considered as nonpolar if C is an sp3 C

Slide 52: 

Electronegativity: Atoms or functional groups that are electronegative relative to hydrogen such as the hal X, O, N, etc. may have a negative inductive effect (-I), depending on their bonding order.Thus these atoms withdraw electron density through the single bond structure of a compound and can assist in the stabilization of negative charge that may form in reactions. The electronic effect is being "induced“ through single bonds. Atoms or functional groups that are electron donating (hydrocarbons, anions) have a positive inductive effect (+I). These groups can help stabilize positive charges in reactions such as protonation of bases.

Slide 53: 

Bonding order and charge: As mentioned above, it is important to consider both the electronegativity and bonding order when analyzing the inductive potential of an atom. For example, oxygen in a hydroxyl group (OH) is electron withdrawing by induction (-I) because the oxygen atom is relatively electronegative and is uncharged in that bonding arrangement. However, oxygen in an "alkoxide" (O-) structure is electron donating (+I) by induction because in this bonding order (a single bond to oxygen) it has an "excess" of electron density. Thus an OH group would help to stabilize a negative charge within a structure, while it's ionized form, the alkoxide, would stabilize a positive charge!

Slide 54: 

Bonding position: The strength of the inductive effect produced by a particular atom or functional group is dependent on it's position within a structure. For example, the further from the site of ionization, the lower the inductive effect. This is illustrated in the example below where the acid with the chlorine atom positioned on a carbon atom nearer the reaction site (OH) is more acidic that the acid where the chlorine atom is positioned further away

Slide 55: 

Examples of –I effect In a molecule of the type S—CH3, the substituent S may be more electronegative than C or less electronegative then C, which generates two different polarizations of the S—C simple bond: S←CH3 or S→CH3 which may be written also using partial charges d+ d- d- d+ S—CH3 or S—CH3 We may also say that in the first case S has a - I effect and in the second S has a +I effect. d+ d- - I + I

Slide 56: 

The different effect of S may be proven by the VERY DIFFERENT RESULTS of the reaction of S—CH3 with water. In such reactions, the polarization of the HO bond in water has also to be considered: + I + I - I - I

Slide 57: 

Another example is concerning the ionization of acids: HA + H2O H3O+ + A- This equilibrium is strongly influenced (shifted to the right) by the capacity of HA to donate it proton, H+, and at its turn, this capacity is influenced by internal factors that enhance the H+ expulsion, or, on the contrary, oppose to this expulsion; the degree of ionization is reflected by the measurable strengths of the acid (Ka) or (pKa) Examples of –I effect

Slide 58: 

Consider the above case of these acids - acetic acid, chloroacetic acid and trichloroacetic acid – that all can ionize (loss of proton from the carboxyl OH), but the easyness of such loss is influenced by the presence/absence of Cl, (an electronegative atom, with a strong –I effect) and by the number of such atoms. Thus they can help stabilize the negative charge in the resulted acid anion, and thus enhance the ionization of the acid. Note the Ka differences between all these acids. Furthermore, the more Cl atoms present, the greater the total -I effect and the greater the ease of ionization (greater Ka or lower pKa).

Slide 59: 

Examples of +I effect Atoms or functional groups that are electron donating (hydrocarbons, anions) have a positive inductive effect (+I). These groups can help stabilize positive charges in reactions such as protonation of bases. Besides metal atoms, alkyl groups are considered to have an electron releasing effect (+I). The greater the volume of the alkyl group, the more intense the +I effect of such group over the neighbouring sequence of the molecule

Slide 60: 

General considerations over the Inductive effects The Inductive effect –I is directly proportional to the electronegativity of the atom that has such an effect, directly proportional with the number of such atoms decreasing with the distance from the reaction centre The effect has to be evaluated as a resut of the difference in electronegativity between the partner atoms of a bond

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