Chemical bonding

1
Chapter 6 and 7

• Chemical bonding

2
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
• Required 2 e- to make a bond
• Bond energy amount of energy required to form or
  to break the bond

3
Ionic Bonding

• Occurs in ionic compound
• Results from transferring electron
• Created a strong attraction among the closely
  pack compound

4
Electron Affinity
Electron Affinity (Eea) The energy released
when a neutral atom gains an electron to form an
anion
5
Covalent Bonding

• Formation of a covalent Bond
• Two atoms come close together, and electrostatic
  interactions begin to develop
• Two nuclei repel each other electrons repel each
  other
• Each nucleus attracts to electrons electrons
  attract both nuclei
• Attractive forces gt 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
• Metallic elements low electronegativities
• 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
• Electrons are not completely transferred
• More electronegative atom d- . (d represents the
  partial negative charge formed)
• Less electronegative atom d

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Relationship Between Electronegativity and Bond
Type

• Predicting bond polarity
• Atoms with similar electronegativity (? EN lt0.4)
  form nonpolar bond
• Atoms whose electronegativity differ by more than
  two (? EN gt 2) form ionic bonds
• Atoms whose electronegativity differ by less than
  two (? EN lt 2) form polar covalent bonds

9
Polarity
10
Examples

• For each of the following pairs of bonds, choose
  the bond that will be more polar
• a. H-P, H-C b. N-O, S-O

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12.3 Polarity and Dipole Moment

• Dipole moment
• a vector quantity from the center of the
  positive charge to the center of negative charge
• Represents with an arrow
• E.g Draw the dipole moment for HF, H2O, HCl, OF

12
13.4 Stable Electron Configurations and
Charges on Ions

• Atoms in stable compounds almost always have a
  noble gas electron configuration
• Predicting Formulas of Ionic Compound
• Electrons lost by a metal come from the
  highest-energy occupied orbital
• Electrons gained by a nonmetal go into
  lowest-energy unoccupied orbital

13
Ions configuration
14
Examples

• Predicting formulas of Ionic compound by showing
  how they loses or gains electrons
• Ca and O
• 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
• Refer to the breakup of a crystal into individual
  ions

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12.6 Lewis Structures

• represents how an atoms valence electrons are
  distributed in a molecule
• Show the bonding involves (the maximum bonds can
  be made)
• Try to achieve the noble gas configuration

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Rules

• Duet Rule sharing of 2 electrons
• E.g H2
• H H
• Octet Rule sharing of 8 electrons
• Carbon, oxygen, nitrogen and fluorine always obey
  this rule in a stable molecule
• E.g F2, O2
• Bonding pair two of which are shared with other
  atoms
• 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
• single bond involves two atoms sharing one
  electron
• Double bond involves two atoms sharing two pair
  of electrons
• Triple bond involves two atoms sharing 3 pair of
  electrons
• Use 6N 2 Rule
• N number of atoms other than Hydrogen

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Dots Lewis Structure

• If
• Total valence (6N 2) 2
• ? 1 double bond
• Total valance e- - (6N 2) 4
• ? two double bonds or 1 triple bond

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Rules for Wring Dot Lewis structure

• Draw a dot Lewis structure of ClO4-
• Calculate the total number of valence electrons
  of all atoms in the molecule
• Cl Valence e- 7
• O Valence e - 6 x 4 24e-
• ClO4- gt total valence e- 7 24 1 ( -1
  charge) 32 e-

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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
• 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

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Rules

• 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 havent
  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-
• O 4 x 6e- 24 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 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 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
• Na O
• H2O NH4
• CF4, BeF2
• CO2 NO3-,