Modern Chemistry Chapter 6- Chemical Bonding

1
Modern Chemistry Chapter 6- Chemical Bonding

• chemical bond- a mutual electrical attraction
  between the nuclei and the valence electrons of
  different atoms that binds the atoms together

2

• ionic bond- a chemical bond that results from
  the electrical attraction between anions and
  cations

3

• Na Cl ? NaCl
• sodium loses its one valence electron to form the
  cation Na
• This allows the atom to become like neon with
  eight electrons in it outer energy level.
• chlorine gains that electron to form the anion
  Cl-
• This allows the atom to become like argon with
  eight electrons in its outer energy level.
• Na is attracted to Cl- (opposites attract) and
  an ionic bond forms

4

• covalent bond- a chemical bond that results from
  the sharing of electron pairs between two atoms
• nonpolar covalent bond- a covalent bond in which
  the bonding electrons are shared equally by the
  atoms forming the bond
• -because their electronegativities are
  essentially equal
• polar covalent bond- a covalent bond in which
  the bonded atoms have an unequal attraction for
  the shared electrons
• -because one atom has a greater electronegativity
  than the other

5

• Polar covalent bond
• H F
• -
• Fluorines greater electronegativity causes the
  shared electrons to move closer to it and creates
  areas of slight positive and negative charge
  forming a polar covalent bond

• Nonpolar covalent bond
• H H
• Equal electronegativities of the hydrogen atoms
  cause the pair of electrons to be shared equally
  and a nonpolar covalent bond to form.

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Determining Bond Type

• Find the absolute difference in the
  electronegativities of the bonding atoms.
• The greater the difference, the greater the
• ionic character which makes it more like an
  ionic bond.
• IF the absolute difference is
• lt 0.3 the bond is nonpolar covalent
• gt 0.3 and lt 1.7 the bond is polar covalent
• gt 1.7 the bond is ionic
• Do section review problems 1-4 on page 177.

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Section Review page 177

• 1- Electron pairs are shared in covalent bonds
  and electrons are transferred between atoms in
  ionic bonds.
• 2-The difference in the electronegativities of
  bonding atoms determines the bond type.
• For problem 3 use the electronegativity chart
  on page 161.
• 3a- Li 1.0 F 4.0 4.0-1.0 3.0 ? ionic bond
• 3b- Cu 1.9 S 2.5 2.5-1.90.6 ? polar covalent
• 3c- I 2.5 Br 2.8 2.8-2.50.3 ? polar
  covalent
• 4- c (0.3) lt b (0.6) lt a (3.0)

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Covalent Bonding Molecular Compounds

• molecule- a neutral group of atoms that are held
  together by covalent bonds
• molecular compound- a chemical compound whose
  simplest units are molecules
• chemical formula- indicates the relative numbers
  of atoms of each kind in a chemical compound by
  using element symbols and numerical subscripts
• molecular formula- shows the types and numbers
  of atoms combined in a single molecule of a
  molecular compound
• diatomic molecule- a molecule containing only
  two atoms

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Formation of a Covalent Bond-see figure 5 on
page 179-

• As two atoms come near one another, the nuclei of
  each atom are attracted to the electrons of the
  other atom.
• This causes a decrease in the potential energy of
  the atoms.
• As a bond between the atoms forms, the potential
  energy of the system reaches its lowest point.
  At this point, the two atoms share at least one
  pair of electrons which are then able to move
  freely between the nuclei of the two atoms in
  overlapping orbitals.
• If the atoms get closer to one another, repulsion
  between the nuclei increases and the potential
  energy increases.

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Bond Length Bond Energy

• bond length- the average distance between two
  bonded atoms
• bond energy- the energy required to break a
  chemical bond and form neutral, isolated atoms
• see figure 7 on page 181
• see table 1 on page 182
• What is the general relationship between bond
  length and bond energy (strength) as seen in
  table 1?

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The Octet Rule

• octet rule- chemical compounds tend to form so
  that each atom, by gaining, losing, or sharing
  electrons, has an octet of electrons in its
  outermost energy level
• EXCEPTIONS
• hydrogen atoms are complete with two electrons
  (H2)
• boron atoms are complete with 6 electrons (BF3)
• some elements show expanded valence involving d
  orbitals (PF5 SF6)

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Electron-Dot Notation

• electron-dot notation- is an electron
  configuration notation in which only the valence
  electrons of an atom of a particular element are
  shown, indicated by dots placed around the
  element symbol
• F
• .

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Lewis Structures

• Lewis structures- are formulas in which atomic
  symbols represent nuclei and inner shell
  electrons, dot pairs adjacent to a single atom
  represent unshared electron pairs, and dashes
  between the atomic symbols represent covalent
  bonds between two atoms
• F-F

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Rules for Drawing Lewis Structures

• 1- Determine the types and numbers of atoms in
  the
• molecule.
• 2- Arrange the atoms to form a skeleton
  structure for the
• molecule. If carbon is present, it is the
  central atom.
• Otherwise, the least electronegative atom
  (except for
• hydrogen) is central. Add the electron dot
  structure for
• each atom in the molecule.
• 3- Connect the atoms with lines to represent
  covalent bonds
• between shared electron pairs.
• 4- Add unshared pairs of electrons so each atom
  (other than
• hydrogen) has eight electrons.

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• structural formula- indicates the kind, number,
  arrangement, and bonds but not the unshared pairs
  of electrons of the atoms in a molecule
• F-F

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Covalent Bonds

• single bond- is a covalent bond in which one pair
  of electrons is shared between two atoms
• H H ? H-H
• -Do practice problems 1-4 on page 186

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Drawing Lewis Structures

• page 186 1
• NH3 N (5 electrons) 3 H (1 electron each)
• H H
• H N H ? H N H

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Drawing Lewis Structures

• Page 186 2 H2S
• H S ? H S
• H H

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Page 186 3 4

• 3- SiH4 H
• H Si H
• H
• 4- PF3
• F P F
• F

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• multiple covalent bonds- are double (two shared
  pairs of electrons) or triple (three shared pairs
  of electrons) bonds
• OO NN
• -Do practice problems 1-2 on page 188.
• -Do Section Review problems 2, 4, 5 on page
  189.

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Drawing Lewis Structures

• page 188 1 CO2
• O C O O C O

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Drawing Lewis Structures

• page 188 2 HCN
• H C N H C N

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Section review page 189

• 2- State the octet rule.
• Atoms will gain, lose, or share electrons so
  that each atom has an octet of electrons.
• 4- a) I Br c) H C C Cl
• b) H d) Cl
• H C H Cl Si Cl
• Br Cl
• e) F O
• F

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Section review page 189

• 5- H N N H
• H H
• H N N H

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Lewis Structure Practice

• Draw the Lewis structures of the following
  molecular compounds. Also use the model kits to
  build the molecule. The highlighted formulas
  represent molecules that contain multiple bonds.
• NH3 CO2 N2 O2
• HBr H2CO3 C2H6 C2H4
• C2H2 PF3

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• NH3 ? H N H
• H
• HBr ? Br H
• C2H2 ? H C C H

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• CO2 ? O C O
• H2CO3 ? H C O O H
• O
• PF3 ? F P F
• F

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• N2 ? N N
• C2H6 ? H H
• H C C H
• H H
• O2 ? O O
• C2H4 ? H C C H
• H H

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Drawing Lewis Structures- review

• 1- Determine which element in the formula will be
  the central atom(s) of the structure.
• 2- Make a probable skeleton arrangement of the
  atoms.
• 3- Put the correct number of dots to equal the
  valence electrons of each atom.
• 4- Draw lines between single electrons of
  adjacent atoms.
• 5- If there are extra dots around adjacent atoms,
  draw multiple bond lines.
• 6- Make sure each atom has 8 electrons either in
  unshared pairs or shared bonds. Remember,
  hydrogen has just two electrons.

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Lewis Structure Quiz

• Draw the Lewis structures for the following
  compounds.
• 1- H2O
• 2- PF3
• 3- SiO2
• 4- SeBr2
• 5- CS2

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Ionic Bonding Ionic Compounds

• ionic compound- is composed of positive and
  negative ions that are combined so that the
  number of positive and negative charges are equal
• formula unit- is the simplest collection of
  atoms from which an ionic compounds formula can
  be established
• eg. NaCl

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Formation of Ionic Bonds

• An atom of an element with low electronegativity
  approaches another with high electronegativity.
• The highly electronegative atom then transfers an
  electron from the atom with low
  electronegativity.
• This creates an anion and a cation.
• The attraction between the ions forms an ionic
  bond.
• Na Cl ? Na Cl -

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Ionic Crystals

• Ionic compounds tend to form an orderly
  arrangement known as a crystal lattice which then
  forms crystals.
• see figure 14 on page 191

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Comparing Ionic Molecular Compounds

• ionic compound molecular compound
• high melting point lower melting point
• high boiling point lower boiling point
• extreme hardness lower hardness (usually)
• brittle less brittle

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Polyatomic Ions

• polyatomic ion- a covalently bonded group of
  atoms with a positive or a negative charge
• Review the list of polyatomic ions given to you
  by the teacher.

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Metallic Bonding

• metallic bond- a chemical bond resulting from
  the attraction between metal atoms and the
  surrounding sea of electrons
• The ability of the electrons to move freely
  between the nuclei of the metal atoms accounts
  for the unique properties of metals.
• This accounts for their being good conductors of
  heat and electricity.

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Properties of Metals

• luster- the ability to absorb light energy and
  immediately re-emit it at the same or similar
  frequency which makes them reflective.
• malleable- the ability of a substance to be
  hammered or beaten into thin sheets
• ductility- the ability of a substance to be
  drawn, pulled, or extruded through a small
  opening to produce a wire

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Molecular Geometry

• VSEPR theory (valence shell electron pair
  repulsion)- allows us to predict the shape of
  molecules. It states that repulsion between the
  sets of valence level electrons surrounding an
  atom causes these sets to be oriented as far
  apart as possible

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VSEPR Theory

• When determining the shape of a molecule using
  VSEPR, use the following steps
• Draw the Lewis structure of the compound.
• Find the central atom(s). Use the letter A for
  this atom.
• Count the number of atoms bonded to the central
  atom. Use the letter B and a subscript for the
  number of atoms bonded to the central atom A.
• Count the number of unshared electron pairs
  around the central A atom.
• Use the letter E and a subscript for the number
  of unshared electron pairs around the central
  atom
• Use this ABE designation to find the molecular
  shape using table 5 on page 200 of the textbook.
• Do practice problems 1 2 on page 201.

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Practice problems page 201

• 1a- F-S-F S A F B 2 E 2
• AB2E2 ? bent or angular
• 1b- Cl-P-Cl P A B Cl 3 E 1
• Cl
• AB3E ? trigonal-pyrimidal

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• H
• H-C-N-H
• H H
• This molecule has two central atoms (C N) so it
  has two molecular shapes that are combined.

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Hybridization

• hybridization- is the mixing or two or more
  atomic orbitals of similar energies on the same
  atom to produce new hybrid atomic orbitals of
  equal energies
• hybrid orbitals- orbitals of equal energy
  produced by the combination of two or more
  orbitals on the same atom
• Hybridization explains the unique qualities of a
  carbon atom with its sp3 orbitals.

43
Intermolecular Forces

• intermolecular forces- the forces of attraction
  between molecules
• dipole-dipole forces- forces of attraction
  between polar molecules
• hydrogen bonding- intermolecular force in which
  a hydrogen atom bonded to a highly
  electronegative atom is attracted to an unshared
  pair of electrons of an electronegative atom in a
  nearby molecule
• London Dispersion forces- intermolecular
  attraction resulting from the constant motion of
  electrons and the creation of instantaneous
  dipoles

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Chapter 6 Review

• Do the following review problems from pages
  209-211 of the textbook.
• 6, 15, 19, 21, 34,
• 43, 48

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End of Chapter Practice

• 6 H (2.1) I (2.5) ? 0.4 polar covalent
• S (2.5) O (3.5)? 1.0 polar covalent
• K (0.8) Br (2.8)? 2.0 ionic
• Si (1.8) Cl (3.0)? 1.2 polar covalent
• K (0.8) Cl (3.0) ? 2.2 ionic
• Se (2.4) S (2.5) ? 0.1 nonpolar covalent
• C (2.5) H (2.1) ? 0.4 polar covalent
• 15 H 1 F 7 Mg 2 O 6
• Al 3 N 5 C 4

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End of Chapter Practice

• 19 Li Ca Cl
• O C P
• Al S

47
End of Chapter Practice

• F
• 21 F C F H Se H
• F
• Br
• I N I Br Si Br
• I Br
• Cl
• H C H
• H

48
End of Chapter Practice

• 34 AB2 linear
• AB3 trigonal planar
• AB4 tetrahedral
• AB5 trigonal bipyramidal
• AB6 octahedral

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End of Chapter Practice

• 43 AB3E trigonal pyramidal
• AB2E2 bent or angular
• AB2E bent or angular

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Honors Chemistry Chapter 6 Test Review

• 40 multiple choice questions
• valence electrons, chemical bonds (how they
  occur)
• atoms potential energy leading to stability and
  bond formation
• polar nonpolar covalent bonds
• difference in electronegativity ionic
  character
• using electronegativities, determine if a bond is
  ionic, polar or nonpolar covalent
• definition of molecule, molecular formula
  (examples), bond length, octet octet rule
• elements meeting octet rule naturally
• how to draw a Lewis structure, identifying a
  Lewis structure, bonding in Lewis structures

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Honors Chemistry Chapter 6 Test Review

• formula of an ionic compound represents
• lattice energy, crystal lattice
• compare properties of ionic molecular compounds
  and the strength of their bonds
• electrons charge of polyatomic ions
• metallic bonds their electrons
• properties of metals the cause
• properties of ionic crystals
• VSEPR definition use
• intermolecular forces (dipole-dipole, hydrogen
  bonding, London dispersion), their relative
  strength and properties

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Chemistry Chapter 6 Test Review

• 25 multiple choice questions
• definitions and functions of valence electrons
  chemical bonds
• bonding potential energy
• polar nonpolar covalent bonds their
  characteristics
• use difference in electronegativity to determine
  bond type
• define molecule, molecular formula, octet
• which elements satisfy the octet rule by
  themselvs
• which elements form multiple covalent bonds
• what is necessary to draw a Lewis structure
  recognize a correct Lewis structure
• properties of ionic vs. covalent compounds
• excess (or deficit) electrons in polyatomic ions
• valence electrons in metallic bonds
• properties of metals and why they occur
• definition of VSEPR
• intermolecular forces and why they occur,
  especially dipole-dipole forces
• polar molecules

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VSEPR Lab

• H2O H O AB2E2
• H
• CO2 O C O AB2
• H AB4
• CH3NH2 H C N H
• H H AB3E

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VSEPR Lab

• H2CO H C O AB3
• H
• CH4 H
• H C H AB4
• H

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VSEPR Lab

• C2H6 H H AB4
• H C C H
• H H AB4
• C2H2 H C C H AB2
• AB2

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VSEPR Lab

• HCOOH H C O H AB3
• AB2E2
• O