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Chapter 6 and 7 Chemical bonding
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12.1 Types of Chemical Bonds Bonds: a force that holds groups of two or more atoms together and makes them function as a unit Bonds: a force that holds groups of two or more atoms together and makes them function as a unit Required 2 e- to make a bond Required 2 e- to make a bond Bond energy: amount of energy required to form or to break the bond Bond energy: amount of energy required to form or to break the bond
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Ionic Bonding Occurs in ionic compound Occurs in ionic compound Results from transferring electron Results from transferring electron Created a strong attraction among the closely pack compound Created a strong attraction among the closely pack compound
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Electron Affinity Electron Affinity (Eea): The energy released when a neutral atom gains an electron to form an anion
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Covalent Bonding Formation of a covalent Bond Formation of a covalent Bond Two atoms come close together, and electrostatic interactions begin to develop Two atoms come close together, and electrostatic interactions begin to develop Two nuclei repel each other; electrons repel each other Two nuclei repel each other; electrons repel each other Each nucleus attracts to electrons; electrons attract both nuclei Each nucleus attracts to electrons; electrons attract both nuclei Attractive forces > repulsive forces; then covalent bond is formed Attractive forces > repulsive forces; then covalent bond is formed
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12.2 Electronegativity Electronegativity (EN): the ability of an atom in a molecule to attract the shared electron in a bond Electronegativity (EN): the ability of an atom in a molecule to attract the shared electron in a bond Metallic elements – low electronegativities Metallic elements – low electronegativities Halogens and other elements in upper right- hand corner of periodic table – high electronegativity Halogens and other elements in upper right- hand corner of periodic table – high electronegativity
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Polarity Polar covalent bonds – the bonding electrons are attracted somewhat more strongly by one atom in a bond Polar covalent bonds – the bonding electrons are attracted somewhat more strongly by one atom in a bond Electrons are not completely transferred Electrons are not completely transferred More electronegative atom: δ-. (δ represents the partial negative charge formed) More electronegative atom: δ-. (δ represents the partial negative charge formed) Less electronegative atom: δ+ Less electronegative atom: δ+
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Relationship Between Electronegativity and Bond Type Predicting bond polarity Predicting bond polarity Atoms with similar electronegativity (Δ EN <0.4) –form nonpolar bond Atoms with similar electronegativity (Δ EN <0.4) –form nonpolar bond Atoms whose electronegativity differ by more than two (Δ EN > 2) – form ionic bonds Atoms whose electronegativity differ by more than two (Δ EN > 2) – form ionic bonds Atoms whose electronegativity differ by less than two (Δ EN < 2) – form polar covalent bonds Atoms whose electronegativity differ by less than two (Δ EN < 2) – form polar covalent bonds
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Polarity
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Examples For each of the following pairs of bonds, choose the bond that will be more polar For each of the following pairs of bonds, choose the bond that will be more polar a.H-P, H-Cb.N-O, S-O
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12.3Polarity and Dipole Moment Dipole moment: Dipole moment: a vector quantity from the center of the positive charge to the center of negative charge a vector quantity from the center of the positive charge to the center of negative charge Represents with an arrow Represents with an arrow E.g Draw the dipole moment for HF, H 2 O, HCl, OF
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13.4 Stable Electron Configurations and Charges on Ions Atoms in stable compounds almost always have a noble gas electron configuration Atoms in stable compounds almost always have a noble gas electron configuration Predicting Formulas of Ionic Compound Predicting Formulas of Ionic Compound Electrons lost by a metal come from the highest- energy occupied orbital Electrons lost by a metal come from the highest- energy occupied orbital Electrons gained by a nonmetal go into lowest- energy unoccupied orbital Electrons gained by a nonmetal go into lowest- energy unoccupied orbital
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Ions configuration
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Examples Predicting formulas of Ionic compound by showing how they loses or gains electrons Predicting formulas of Ionic compound by showing how they loses or gains electrons Ca and O Ca and O Sr and Cl Sr and Cl
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12.5 Ionic Compound Lattice energy (U) – the sum of the electrostatic interaction energies between ions in a solid Lattice energy (U) – the sum of the electrostatic interaction energies between ions in a solid Refer to the breakup of a crystal into individual ions Refer to the breakup of a crystal into individual ions
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12.6 Lewis Structures represents how an atom’s valence electrons are distributed in a molecule represents how an atom’s valence electrons are distributed in a molecule Show the bonding involves (the maximum bonds can be made) Show the bonding involves (the maximum bonds can be made) Try to achieve the noble gas configuration Try to achieve the noble gas configuration
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Rules Duet Rule: sharing of 2 electrons Duet Rule: sharing of 2 electrons E.g H 2 E.g H 2 H : H H : H Octet Rule: sharing of 8 electrons Octet Rule: sharing of 8 electrons Carbon, oxygen, nitrogen and fluorine always obey this rule in a stable molecule Carbon, oxygen, nitrogen and fluorine always obey this rule in a stable molecule E.g F 2, O 2 E.g F 2, O 2 Bonding pair: two of which are shared with other atoms Bonding pair: two of which are shared with other atoms Lone pair or nonbonding pair: those that are not used for bonding Lone pair or nonbonding pair: those that are not used for bonding
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12.7 Lewis Structures of Molecules with Multiple Bonds Recall: Elements typically obey the octet rule; they are surrounded by eight electrons Recall: Elements typically obey the octet rule; they are surrounded by eight electrons single bond: involves two atoms sharing one electron single bond: involves two atoms sharing one electron Double bond: involves two atoms sharing two pair of electrons Double bond: involves two atoms sharing two pair of electrons Triple bond: involves two atoms sharing 3 pair of electrons Triple bond: involves two atoms sharing 3 pair of electrons Use 6N + 2 Rule Use 6N + 2 Rule N = number of atoms other than Hydrogen N = number of atoms other than Hydrogen
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Dots Lewis Structure If If Total valence – (6N + 2) = 2 Total valence – (6N + 2) = 2 1 double bond 1 double bond Total valance e- - (6N + 2) = 4 Total valance e- - (6N + 2) = 4 two double bonds or 1 triple bond two double bonds or 1 triple bond
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Rules for Wring Dot Lewis structure Draw a dot Lewis structure of ClO 4 - Draw a dot Lewis structure of ClO 4 - Calculate the total number of valence electrons of all atoms in the molecule Calculate the total number of valence electrons of all atoms in the molecule Cl – Valence e- = 7 Cl – Valence e- = 7 O – Valence e - = 6 x 4 = 24e- O – Valence e - = 6 x 4 = 24e- ClO 4 - => total valence e- = 7 + 24 +1 ( -1 charge) = 32 e- ClO 4 - => total valence e- = 7 + 24 +1 ( -1 charge) = 32 e-
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Rules Create a skeletal structure using the following rules: Create a skeletal structure using the following rules: a.Hydrogen atoms (if present) are always on the “outside” of the structure. They form only one bond a.Hydrogen atoms (if present) are always on the “outside” of the structure. They form only one bond b.The central atom is usually least electronegative. It is also often unique (i.e,. the only one atom of the element in the molecule). Remember, there might be no “central” atom. b.The central atom is usually least electronegative. It is also often unique (i.e,. the only one atom of the element in the molecule). Remember, there might be no “central” atom. c.Connect bonded atoms by line (2-electron, covalent bonds c.Connect bonded atoms by line (2-electron, covalent bonds
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Rules Place lone pairs around outer atoms (except hydrogen) so that each atom has an octet Place lone pairs around outer atoms (except hydrogen) so that each atom has an octet
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Rules Calculate the number of electrons you haven’t used. Subtract the number of electrons used so far, including electrons in lone pair and bonding pairs, from the total in Step 1. Assign any remaining electrons to the central atom as lone pair Calculate the number of electrons you haven’t used. Subtract the number of electrons used so far, including electrons in lone pair and bonding pairs, from the total in Step 1. Assign any remaining electrons to the central atom as lone pair Cl-O bonds = 4 x 2e- = 8 e- Cl-O bonds = 4 x 2e- = 8 e- O – 4 x 6e- = 24 e- O – 4 x 6e- = 24 e- Total used = 8 + 24 = 32 e- Total used = 8 + 24 = 32 e-
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Rules If the central atom is B (boron) or Be (beryllium), skip this step If the central atom is B (boron) or Be (beryllium), skip this step If the central atom has an octet after step 4, skip this step If the central atom has an octet after step 4, skip this step If the central atom has only 6 electrons, move a lone pair from an outer atom to form a double bond between outer atom and the central atom If the central atom has only 6 electrons, move a lone pair from an outer atom to form a double bond between outer atom and the central atom If the central atom has only 4 electrons, do Step 5a to two different outer atoms (i.e, form two double bonds) or twice to one outer atom (i.e., form one triple bond) If the central atom has only 4 electrons, do Step 5a to two different outer atoms (i.e, form two double bonds) or twice to one outer atom (i.e., form one triple bond)
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Examples Give the Lewis structure for the following Give the Lewis structure for the following NaO NaO H 2 ONH 4 + H 2 ONH 4 + CF 4,BeF 2 CF 4,BeF 2 CO2 NO 3 -, CO2 NO 3 -,
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