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Ch. 10: Chemical Bonding. Dr. Namphol Sinkaset Chem 152: Introduction to General Chemistry. I. Chapter Outline. Introduction Lewis Dot Structures Ionic Compounds Covalent Compounds Shapes of Molecules Polar Bonds and Polar Molecules. I. Introduction.

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ch 10 chemical bonding

Ch. 10: Chemical Bonding

Dr. Namphol Sinkaset

Chem 152: Introduction to General Chemistry

i chapter outline
I. Chapter Outline
  • Introduction
  • Lewis Dot Structures
  • Ionic Compounds
  • Covalent Compounds
  • Shapes of Molecules
  • Polar Bonds and Polar Molecules
i introduction
I. Introduction
  • Bonding theories allow prediction of how atoms form compounds and the shapes they will take.
  • The 3-D shape of a molecule determines many of its physical properties.
  • With the advent of faster computers, simulations allow screening of drug candidates.
i lewis theory
I. Lewis Theory
  • Lewis theory is simple to apply, but amazingly powerful.
  • It can be used to go from a formula of a compound to its 3-D structure.
  • Lewis theory centers on how valence electrons are used by atoms to form compounds.
ii lewis dot structures
II. Lewis Dot Structures
  • Valence e-’s are the most important e-’s in bonding.
  • Lewis dot structures are a way to depict the valence e-’s of atoms.
  • Lewis dot symbols have two parts:
    • element symbol: represents nucleus and core e-
    • dots around symbol: represent valence e-’s
ii origin of dot structures
II. Origin of Dot Structures
  • Oxygen has 6 valence electrons, so it’s dot structure will have 6 dots.
ii lewis dot structures1
II. Lewis Dot Structures
  • The number of valence e- is given by the element’s group number!!
ii the central idea
II. The Central Idea
  • Lewis realized that noble gases all had 8 valence e-’s (except He).
  • He reasoned that having 8 valence e-’s (or 2 for H and He) leads to stability, or being unreactive.
  • These special configurations of valence e-’s are known as an octet or a duet.
ii the octet rule
II. The Octet Rule
  • In Lewis theory, atoms bond in order to obtain eight valence electrons.
  • The octet rule states that in chemical bonding, atoms transfer or share electrons to obtain outer shells with eight electrons.
  • The octet rule is a main-group concept, generally applying to all except H and Li.
ii sample problem
II. Sample Problem
  • Draw Lewis dot structures for Ca, Ge, Se, and Br.
iii transfer of electrons
III. Transfer of Electrons
  • When metals bond with nonmetals, the metal transfers electrons to the nonmetal to form an ionic compound.
  • The positive charge of the metal cation and the negative charge of the nonmetal anion holds the ionic compound together.
  • We can show this with Lewis theory.
iii sample problem
III. Sample Problem
  • Use Lewis dot structures to depict the compound that forms between:
    • magnesium and nitrogen
    • aluminum and bromine
iv sharing electrons
IV. Sharing Electrons
  • When nonmetals bond with other nonmetals, a covalent compound is formed.
  • Electrons are shared in covalent compounds in order to achieve an octet.
  • Electrons that appear in the space between atoms count towards the octet of both atoms.
iv bonding vs lone pairs
IV. Bonding vs. Lone Pairs
  • Electrons shared between atoms are bonding pair electrons.
  • Electrons that are only on one atom are lone pair or nonbonding electrons.
iv lines as bonds
IV. Lines as Bonds
  • Generally, bonding pair electrons are represented by lines.
  • Note that a line equals two electrons.
iv why some elements exist as diatomics
IV. Why Some Elements Exist as Diatomics
  • If two hydrogens combine, they can satisfy their duet.
  • If two chlorines combine, they can satisfy their octet.
iv higher order bonds
IV. Higher Order Bonds
  • In some compounds, atoms need to share more than one electron pair to reach an octet.
  • Double bond – atoms share 4 electrons
  • Triple bond – atoms share 6 electrons
  • However, higher order bonds are a last resort used by atoms to reach an octet!
iv diatomic oxygen
IV. Diatomic Oxygen
  • Oxygen has 6 valence electrons.
  • Sharing one pair satisfies the octet of one oxygen atom, but not the other.
  • Since there are no more electrons that can be used, higher order bonds must be made.
iv triple bonds
IV. Triple Bonds
  • Sometimes six electrons need to be shared by two atoms.
  • An example is diatomic nitrogen.
iv single double triple bonds
IV. Single, Double, Triple Bonds
  • Higher order bonds mean more stability and shorter internuclear distances.
iv steps for drawing lewis structures
IV. Steps for Drawing Lewis Structures
  • Determine total # of valence e-. (Cation, subtract e-’s for charge; anion, add e-’s for charge).
  • Place atom w/ lower Group # (lower electronegativity) as the central atom.
  • Attach other atoms to central atom with single bonds.
  • Fill octet of outer atoms. (Why?)
  • Count # of e- used so far. Place remaining e- on central atom in pairs.
  • If necessary, form higher order bonds to satisfy octet rule of central atom.
  • Allow expanded octet for central atoms from Period 3 or lower.
iv sample problem
IV. Sample Problem
  • Draw correct Lewis structures for NF3, CO2, SeCl2, CCl4, and H2CO.
iv exceptions to the octet rule
IV. Exceptions to the Octet Rule
  • Lewis theory is too simple to cover all bonding possibilities.
  • Some exceptions exist to the octet rule:
    • e- deficient atoms like Be and B, e.g. BeCl2 and BF3.
    • Compounds w/ odd # of e-’s: free radicals. Examples include NO and NO2.
    • Expanded valence – when d orbitals are used to accommodate more than an octet.
v vsepr theory
V. VSEPR Theory
  • From a correct Lewis structure, we can get to the 3-D shape using this theory.
  • VSEPR stands for valence shell electron pair repulsion.
  • The theory is based on the idea that electron groups (lone pairs, single bonds, or multiple bonds) repel each other.
v linear geometry
V. Linear Geometry
  • CO2 has two electron groups.
  • The two double bonds try to get as far away from each other as possible.
v trigonal planar
V. Trigonal Planar
  • Formaldehyde has three electron groups around the central atom.
  • They form 120° angles to get away from each other.
v tetrahedral
V. Tetrahedral
  • Methane has four electron groups.
  • A bond angle of 109.5° keeps them furthest apart.
v tetrahedral based shapes
V. Tetrahedral-Based Shapes
  • Other shapes are based on tetrahedral with bonding groups being replaced by lone pair electrons.
  • The electron geometry is the shape based on all electron types (bonding and lone pair).
  • The molecular geometry is the shape based on just atoms.
v ammonia
V. Ammonia
  • The central atom has three bonds and one lone pair (4 electron groups).
  • EG = tetrahedral
  • When drawing MG, lone pairs are left off!
v trigonal pyramidal
V. Trigonal Pyramidal
  • Ammonia has a trigonal pyramidal molecular geometry.
v water
V. Water
  • The central atom has two bonds and two lone pairs (4 electron groups).
  • EG = tetrahedral
  • Again, lone pairs omitted when drawing MG!
v bent
V. Bent
  • Water has a bent molecular geometry.
v drawing w perspective
V. Drawing w/ Perspective
  • We use the conventions below to depict a 3-D object on a 2-D surface.
v practice problem
V. Practice Problem
  • Draw the molecular shapes for ClO2-, BF3, and NF3. Indicate the name of the molecular and electronic geometries for each as well.
vi sharing electrons
VI. Sharing Electrons
  • It is not reasonable to assume that all atoms will share electrons equally.
  • Some atoms will pull electrons closer to them.
  • Electronegativity is the ability of an element to pull electrons in a covalent bond closer.
vi effect of electronegativity
VI. Effect of Electronegativity
  • O is more electronegative than H, so in an O-H bond, the bonding electrons are more likely found around the O.
vi dipole moment
VI. Dipole Moment
  • The unequal sharing leads to a partial charge separation in the bond called a dipole moment.
  • Covalent bonds with a dipole moment are polar covalent bonds.
  • The greater the difference in electronegativity, the greater the polarity.
vi bonding type
VI. Bonding Type
  • Electronegativity difference can be used to determine the type of bond.
vi sample problem
VI. Sample Problem
  • Calculate the difference in electronegativity for the following pairs of atoms and determine of the bond between them is pure covalent, polar covalent, or ionic.
    • I and I
    • Cs and Br
    • P and O
vi polar molecules
VI. Polar Molecules
  • Just because a molecule has polar bonds doesn’t mean that the molecule is polar overall.
  • The degree of polarity in each bond and the orientation of those polar bonds determines whether the molecule is polar.
vi polarity vectors
VI. Polarity Vectors
  • We can use vectors to represent dipole moments and analyze the resultant.