Presentation on theme: "Ch. 10: Chemical Bonding Dr. Namphol Sinkaset Chem 152: Introduction to General Chemistry."— Presentation transcript:
Ch. 10: Chemical Bonding Dr. Namphol Sinkaset Chem 152: Introduction to General Chemistry
I. Chapter Outline I.Introduction II.Lewis Dot Structures III.Ionic Compounds IV.Covalent Compounds V.Shapes of Molecules VI.Polar Bonds and Polar Molecules
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. HIV-protease Inhibitors
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 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: 1)element symbol: represents nucleus and core e- 2)dots around symbol: represent valence e-’s
II. Origin of Dot Structures Oxygen has 6 valence electrons, so it’s dot structure will have 6 dots.
II. Lewis Dot Structures The number of valence e- is given by the element’s group number!!
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 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 Draw Lewis dot structures for Ca, Ge, Se, and Br.
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. Formation of Potassium Chloride
III. Formation of Sodium Sulfide
III. Sample Problem Use Lewis dot structures to depict the compound that forms between: magnesium and nitrogen aluminum and bromine
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. Lewis Structure for Water
IV. Octet/Duets Satisfied
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 Generally, bonding pair electrons are represented by lines. Note that a line equals two electrons.
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 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 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. Double Bond Formation
IV. Triple Bonds Sometimes six electrons need to be shared by two atoms. An example is diatomic nitrogen.
IV. Single, Double, Triple Bonds Higher order bonds mean more stability and shorter internuclear distances. Bond TypeInternuclear Distance (pm) Single148 Double121 Triple110
IV. Steps for Drawing Lewis Structures 1)Determine total # of valence e-. (Cation, subtract e-’s for charge; anion, add e-’s for charge). 2)Place atom w/ lower Group # (lower electronegativity) as the central atom. 3)Attach other atoms to central atom with single bonds. 4)Fill octet of outer atoms. (Why?) 5)Count # of e- used so far. Place remaining e- on central atom in pairs. 6)If necessary, form higher order bonds to satisfy octet rule of central atom. 7)Allow expanded octet for central atoms from Period 3 or lower.
IV. Sample Problem Draw correct Lewis structures for NF 3, CO 2, SeCl 2, CCl 4, and H 2 CO.
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. BeCl 2 and BF 3. Compounds w/ odd # of e-’s: free radicals. Examples include NO and NO 2. Expanded valence – when d orbitals are used to accommodate more than an octet.
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 CO 2 has two electron groups. The two double bonds try to get as far away from each other as possible.
V. Trigonal Planar Formaldehyde has three electron groups around the central atom. They form 120° angles to get away from each other.
V. Tetrahedral Methane has four electron groups. A bond angle of 109.5° keeps them furthest apart.
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 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 Ammonia has a trigonal pyramidal molecular geometry.
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 Water has a bent molecular geometry.
V. Geometry Summary Electron Groups Bonding Groups Lone Pairs Electron Geometry Ideal Angle Molecular Geometry 220Linear180°Linear 330Trigonal planar 120°Trigonal planar 321Trigonal planar 120°Bent 440Tetrahedral109.5°Tetrahedral 431 109.5°Trigonal pyramidal 422Tetrahedral109.5°Bent
V. Drawing w/ Perspective We use the conventions below to depict a 3-D object on a 2-D surface.
V. Practice Problem Draw the molecular shapes for ClO 2 -, BF 3, and NF 3. Indicate the name of the molecular and electronic geometries for each as well.
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 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 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. Electronegativity Values
VI. Bonding Type Electronegativity difference can be used to determine the type of bond.
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 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 We can use vectors to represent dipole moments and analyze the resultant.