Unit 2 Review: Molecular and Ionic Compound Structure and Properties

Topic 2.1 - Types of Chemical Bonds

• Learning Objective: Explain the relationship between the type of bonding and the properties of the elements participating in the bond
• Essential Knowledge:
  • Electronegativity values for the representative elements increase going from left to right across a period and decrease going down a group. These trends can be understood qualitatively through the electronic structure of the atoms, the shell model, and Coulomb’s law.
  • Valence electrons shared between atoms of similar electronegativity constitute a nonpolar covalent bond. For example, bonds between carbon and hydrogen are effectively nonpolar even though carbon is slightly more electronegative than hydrogen.
• Electronegativity: The tendency of an atom to attract a shared pair of electrons towards itself.
  • Increases from left to right across a period
  • Decreases down a group
• Polarity: the degree of electron sharing within a molecule
  • Polar: Two atoms share electrons unequally
  • Non polar: Two atoms share electrons equally

Topic 2.2 - Intramolecular Force and Potential Energy

• Learning Objective: Represent the relationship between potential energy and distance between atoms, based on factors that influence the interaction strength.
• Essential Knowledge:
  • A graph of potential energy versus the distance between atoms is a useful representation for describing the interactions between atoms. Such graphs illustrate both the equilibrium bond length (the separation between atoms at which the potential energy is lowest) and the bond energy (the energy required to separate the atoms).
  • In a covalent bond, the bond length is influenced by both the size of the atom’s core and the bond order (i.e., single, double, triple). Bonds with a higher order are shorter and have larger bond energies.
• Essential Knowledge (cont)
  • Coulomb’s law can be used to understand the strength of interactions between cations and anions.
    • Because the interaction strength is proportional to the charge on each ion, larger charges lead to stronger interactions.
    • Because the interaction strength increases as the distance between the centers of the ions (nuclei) decreases, smaller ions lead to stronger interactions.
• Potential Energy and Separation Distance
  • Illustrates the energy between two atoms as a function of distance
  • The “lowest” part of the curve shows the energy (436 kJ/mol) that would be necessary to break the bond between atoms as well as the ideal bond length (74 pm)
  • Atoms repel at short distances
  • Atoms are attracted to one another at longer distances but eventually are too far apart to experience attractive forces
• Chemical Bonds

• Coulomb’s Law
  • States that like charges repel and opposite charges attract, with a force proportional to the product of the charges and inversely proportional to the square of the distance between them.
  • Expressed mathematically as:
    F = k \frac{q1 q2}{r^2}
  • Where:
    • F is the electrostatic force
    • k is Coulomb's constant
    • q1 and q2 are the magnitudes of the charges
    • r is the distance between the charges

Topic 2.3 - Structure of Ionic Solids

• Learning Objective: Represent an ionic solid with a particulate model that is consistent with Coulomb’s law and the properties of the constituent ions.
• Essential Knowledge:
  • The cations and anions in an ionic crystal are arranged in a systematic, periodic 3-D array that maximizes the attractive forces among cations and anions while minimizing the repulsive forces.
• Structure of Ionic Crystals
  • A “crystal lattice” of alternating cations and anions
  • Highly organized
  • Strong attraction between oppositely charged ions
  • Good conductors when melted or dissolved
  • High melting and boiling points

Topic 2.4 - Structure of Metals and Alloys

• Learning Objective: Represent a metallic solid and/or alloy using a model to show essential characteristics of the structure and interactions present in the substance.
• Essential Knowledge:
  • Metallic bonding can be represented as an array of positive metal ions surrounded by delocalized valence electrons (i.e., a “sea of electrons”).
  • Interstitial alloys form between atoms of different radii, where the smaller atoms fill the interstitial spaces between the larger atoms (e.g., with steel in which carbon occupies the interstices in iron).
  • Substitutional alloys form between atoms of comparable radius, where one atom substitutes for the other in the lattice. (In certain brass alloys, other elements, usually zinc, substitute for copper.)
• Essential Knowledge (cont)
  • Valence electrons shared between atoms of unequal electronegativity constitute a polar covalent bond.
    • The atom with a higher electronegativity will develop a partial negative charge relative to the other atom in the bond.
    • In single bonds, greater differences in electronegativity lead to greater bond dipoles.
    • All polar bonds have some ionic character, and the difference between ionic and covalent bonding is not distinct but rather a continuum.
  • The difference in electronegativity is not the only factor in determining if a bond should be designated as ionic or covalent. Generally, bonds between a metal and nonmetal are ionic, and bonds between two nonmetals are covalent. Examination of the properties of a compound is the best way to characterize the type of bonding.
  • In a metallic solid, the valence electrons from the metal atoms are considered to be delocalized and not associated with any individual atom.
• Organization of a Metallic Solid
  • Metallic solids are composed of metal cations surrounds by a “sea of mobile valence electrons
  • This arrangement gives metals unique properties:
    • Ductility: ability to be drawn into wires
    • Malleability: ability to be flattened into sheets
    • Good conductors of heat and electricity
• Alloys
  • What are alloys?
    • A combination of metals or metals combine with one or more other elements
      • Ex. brass, bronze, sterling silver, steel
  • Two main categories of alloys
    • Interstitial: Form between atoms with different radii
      • Ex. steel in which carbon occupies the interstices in iron
    • Substitutional: Form between atoms with similar radii
      • Ex. In certain brass alloys, other elements, usually zinc, substitute for copper.
  • Substitutional alloy: Atoms of one metal are substituted by atoms of another metal.
  • Interstitial alloy: Different metal occupies interstitial spaces (holes) in the lattice structure.

Topic 2.5 - Lewis Diagrams

• Learning Objective: Represent a molecule with a Lewis diagram.
• Essential Knowledge:
  • Lewis diagrams can be constructed according to an established set of principles

Topic 2.6 - Resonance and Formal Charge

• Learning Objective: Represent a molecule with a Lewis diagram that accounts for resonance between equivalent structures or that uses formal charge to select between nonequivalent structures.
• Essential Knowledge:
  • In cases where more than one equivalent Lewis structure can be constructed, resonance must be included as a refinement to the Lewis structure. In many such cases, this refinement is needed to provide qualitatively accurate predictions of molecular structure and properties.
  • The octet rule and formal charge can be used as criteria for determining which of several possible valid Lewis diagrams provides the best model for predicting molecular structure and properties.
  • As with any model, there are limitations to the use of the Lewis structure model, particularly in cases with an odd number of valence electrons.
• Resonance Structure
  • A set of two or more Lewis structures that can be used to describe a single molecule or polyatomic ion.
    • The movement of delocalized electrons
    • Arrangement of atoms does not change as it would for an isomer
  • Not all resonance structures are equally stable
  • The most stable resonance structure:
    • Has all or most of the atoms with a formal charge of zero
    • Puts the negative formal charge on the most electronegative atom
• Determining Formal Charge
  • Formal Charge = Valence Electrons - Non-Bonding Val Electrons - (Bonding Electrons/2)

Topic 2.7 - VSEPR and Bond Hybridization

• Learning Objective: Based on the relationship between Lewis diagrams, VSEPR theory, bond orders, and bond polarities:
  • Explain structural properties of molecules.
  • Explain electron properties of molecules.
• Essential Knowledge:
  • VSEPR theory uses the Coulombic repulsion between electrons as a basis for predicting the arrangement of electron pairs around a central atom.
  • Both Lewis diagrams and VSEPR theory must be used for predicting electronic and structural properties of many covalently bonded molecules and polyatomic ions, including the following:
    • Molecular geometry
    • Bond angles
    • Relative bond energies based on bond order
    • Relative bond lengths (multiple bonds, effects of atomic radius)
    • Presence of a dipole moment
• Essential Knowledge (cont)
  • The terms “hybridization” and “hybrid atomic orbital” are used to describe the arrangement of electrons around a central atom. When the central atom is sp hybridized, its ideal bond angles are 180°; for sp2 hybridized atoms the bond angles are 120°; and for sp3 hybridized atoms the bond angles are 109.5°.
  • Bond formation is associated with overlap between atomic orbitals. In multiple bonds, such overlap leads to the formation of both sigma and pi bonds. The overlap is stronger in sigma than pi bonds, which is reflected in sigma bonds having greater bond energy than pi bonds. The presence of a pi bond also prevents the rotation of the bond and leads to structural isomers.
• Molecular Shape and Polarity
  • Regardless of differences in electronegativity values, some molecular shapes will always be nonpolar because their bond dipoles cancel out.
  • For an ABn molecule, the following shapes will always be nonpolar
    • Tetrahedral, Trigonal Planar, Linear
• Types of Hybridization
  • Hybridization: the mixing of atomic orbitals into new hybrid orbitals suitable for the pairing of electrons to form chemical bonds
    • The arrangement of electrons around a central atom.
  • We are responsible for the three main types of hybridization: sp, sp2, and sp3
  • sp hybridization - one s and one p orbital hybridizes
    • Present in atoms with only two electron regions
    • Bond angle of 180o
  • sp2 hybridization - one s and two p orbitals hybridize
    • Present in atoms with three electron regions
    • Bond angle of 120o
  • sp3 hybridization - one s and three p orbitals hybridize
    • Present in atoms with four electron regions
    • Bond angle 109.5o
• Sigma and Pi Bonding
  • The first bond will always be a sigma bond (σ)
    • Head to head overlap of orbitals
  • Every bond after that will be a pi bond (π)
    • Side to side overlap of orbitals
  • Single bond: one sigma bond
  • Double bond: one sigma bond, one pi bond
  • Triple bond: one sigma bond, two pi bonds