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2. Ionic Compounds They are formed by the transfer of one or more valence electrons from one atom to another Electropositive atoms: give up electrons and form cations. Electronegative atoms: accept electrons and form anions Ionic compounds: are composed of positively charged cations and negatively charged anions 2

3. Generally speaking:  within a given horizontal row in the periodic table, the more electropositive elements are those farthest to the left,  and the more electronegative elements are those farthest to the right.  Within a given vertical column, the more electro-positive elements are those toward the bottom,  and the more electronegative elements are those toward the top. 3

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6. Covalent compounds A covalent bond: is formed when two atoms share one or more electron pairs. A molecule consists of two or more atoms joined by covalent bonds Bond energy: is the energy necessary to break a mole (6.022 x 10 23 )of covalent bonds. Bond length: is the average distance between two covalently bonded atoms 6

7. Carbon and the Covalent Bond  carbon is neither strongly electropositive nor strongly electronegative   Carbon atoms have neither a strong tendency to lose all their electrons (and become C 4 + ) nor a strong tendency to gain four electrons  (and become C 4- )  C usually forms covalent bonds with other atoms by sharing electrons. 7

8. Carbon–Carbon Single Bonds  C has ability to share electrons not only with different elements but also with other carbon atoms.  In his case the electrons are shared equally between the two identical carbon atoms  As shown below in the example two C atoms may be bonded to one another, and each of these carbon atoms may be linked to other atoms (  Chains). 8

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10. Polar Covalent Bonds  A polar covalent bond is a covalent bond in which the electron pair is not shared equally between the two atoms.  Example: Hydrogen chloride molecule (HCl) In case of HCl, the shared electron pair is attracted more toward the chlorine, which therefore is slightly negative (partial negative charge) with respect to the hydrogen. 10

11. Polarity of Bonds: depends on the electronegativity difference 11

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13. Multiple Covalent Bonds  In a double bond, two electron pairs are shared between two atoms.  Nonbonding electrons, or unshared electron pairs, reside on one atom.  In a triple bond, three electron pairs are shared between two atoms. 13

14. Valence  The valence of an element is the number of bonds that an atom of the element can form.  The number is usually equal to the number of electrons needed to fill the valence shell.  Oxygen, for example, has six valence electrons but a valence of only 2. 14

16. Formal Charges: are the charges that each atom carries, and can be calculated as follows Formal charge = Valence electrons – bonds – electrons 16

17. Resonance: arises whenever we can write two or more structures for a molecule with different arrangements of electrons but identical arrangement of atoms Example: Carbonate ion The total number of valence electrons in the carbonate ion is 24 (4 from the carbon, 3 x 6 = 18 from the three oxygens, plus 2 more electrons that give the ion its negative charge. Physical measurements tell us that none of the foregoing structures accurately describes the real carbonate ion. In the real carbonate ion, the two formal negative charges are spread equally over the three oxygen atoms.

18. The octet rule  The octet rule states that elements will gain, lose, or share electrons to achieve eight electrons in their outermost (valence) shell.  There are some exceptions to the octet rule. Third row elements (such as sulfur and phosphorus) can hold up to 18 electrons in their outermost valence shell (3s, 3p, and 3d orbitals. 18

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20. If there are two equally long continuous chains, choose the one with most branches

21.  Alkanes are insoluble in water. This is because water molecules are polar, whereas alkanes are nonpolar (all the C !C and C!H bonds are nearly purely covalent).  The OH bond in a water molecule is strongly polarized by the high electronegativity of oxygen  This polarization places a partial positive charge on the hydrogen  atom and a partial negative charge on the oxygen atom  As a result, the hydrogen atoms in one water molecule are strongly attracted to the oxygen atoms  in other water molecules, and the small size of the H atoms allows the molecules to approach each other very closely. This special attraction is called hydrogen bonding.  To intersperse alkane and water molecules, we would have to break up the hydrogen bonding interactions between water molecules, which would require considerable energy  Alkanes, with their nonpolar C H bonds, cannot replace hydrogen bonding among water molecules with attractive alkane–water interactions that are comparable in strength, so mixing alkane molecules and water molecules is not an energetically favored process.

22. Alkanes have lower boiling points for a given molecular weight than most other organic compounds. This is because they are nonpolar molecules. They are constantly moving, and the electrons in a nonpolar molecule can become unevenly distributed within the molecule, causing the molecule to have partially positive and partially negative ends. The temporarily polarized molecule causes its neighbor to become tem- porarily polarized as well, and these molecules are weakly attracted to each other. Such interactions between molecules are called van der Waals attractions. Because they are weak attractions, the process of separating molecules from one another (which is what we do when we convert a liquid to a gas) requires relatively little energy, and the boiling points of these compounds are relatively low. Since these attractive forces can only operate over short distances between the surfaces of molecules, the boiling points of alkanes rise as the chain length increases and fall as the chains become branched and more nearly spherical in shape

25. Reactions of Alkanes: 1. Oxidation and Combustion; Alkanes as Fuels  Combustion of hydrocarbons is an oxidation in which C-H bonds are replaced by C-O bonds  Exothermic reactions produce heat.

26. 2. Halogenation of Alkanes

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30. Ionic Compounds They are formed by the transfer of one or more valence electrons from one atom to another Electropositive atoms: give up electrons and form cations. Electronegative atoms: accept electrons and form anions Ionic compounds: are composed of positively charged cations and negatively charged anions 30

31. Generally speaking:  within a given horizontal row in the periodic table, the more electropositive elements are those farthest to the left,  and the more electronegative elements are those farthest to the right.  Within a given vertical column, the more electro-positive elements are those toward the bottom,  and the more electronegative elements are those toward the top. 31

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34. Covalent compounds A covalent bond: is formed when two atoms share one or more electron pairs. A molecule consists of two or more atoms joined by covalent bonds Bond energy: is the energy necessary to break a mole (6.022 x 10 23 )of covalent bonds. Bond length: is the average distance between two covalently bonded atoms 34

35. Carbon and the Covalent Bond  carbon is neither strongly electropositive nor strongly electronegative   Carbon atoms have neither a strong tendency to lose all their electrons (and become C 4 + ) nor a strong tendency to gain four electrons  (and become C 4- )  C usually forms covalent bonds with other atoms by sharing electrons. 35

36. Carbon–Carbon Single Bonds  C has ability to share electrons not only with different elements but also with other carbon atoms.  In his case the electrons are shared equally between the two identical carbon atoms  As shown below in the example two C atoms may be bonded to one another, and each of these carbon atoms may be linked to other atoms (  Chains). 36

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38. Polar Covalent Bonds  A polar covalent bond is a covalent bond in which the electron pair is not shared equally between the two atoms.  Example: Hydrogen chloride molecule (HCl) In case of HCl, the shared electron pair is attracted more toward the chlorine, which therefore is slightly negative (partial negative charge) with respect to the hydrogen. 38

39. Polarity of Bonds: depends on the electronegativity difference 39

41. Reactions of Alkenes 41  The most common reaction of alkenes is addition of a reagent to the carbons of the double bond to give a product with a C-C single bond.  What is the difference between addition and substitution reactions? Substitution: Addition:

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43. 43 When an unsymmetric reagent adds to an unsymmetric alkene, the electropositive part of the reagent bonds to the carbon of the double bond that has the greater number of hydrogen atoms attached to it Example 2 for reaction b:

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56. 56 In all meta -directing groups, the atom connected to the ring carries a full or partial positive charge and will therefore withdraw electrons from the ring. All meta-directing groups are therefore ring-deactivating groups. Ortho,para-directing groups in general supply electrons to the ring and are therefore ring activating. Halogens (F, Cl, Br, and I) bring about the only important exception to these rules. Because they are strongly electron withdrawing, the halogens are ring deac-tivating; but because they have unshared electron pairs, they are ortho,para directing.

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