Chemistry Chapter 7: Chemical Bonding and Molecular Geometry

Chemical Bonding and Molecular Geometry

Ionic Bonding

• Ionic bonds are formed when a cation (positive ion) and an anion (negative ion) are attracted to each other through electrostatic forces.
• Cations are formed when a neutral atom loses one or more electrons, and anions are formed when a neutral atom gains one or more electrons.
• Ionic compounds are composed of ions and have a crystalline structure, are rigid and brittle, and have high melting and boiling points.
• Ionic compounds are poor conductors of electricity in the solid state but are excellent conductors when dissolved or melted.
• The formation of ionic compounds can be explained by the periodic properties of the elements.

The Formation of Ionic Compounds

• Binary ionic compounds are composed of just two elements: a metal (which forms the cations) and a nonmetal (which forms the anions).
• The formula of an ionic compound represents the simplest ratio of the numbers of ions necessary to give identical numbers of positive and negative charges.
• The attractive forces between ions are isotropic, meaning they are the same in all directions.

Electronic Structures of Cations

• When forming a cation, an atom of a main group element tends to lose all of its valence electrons, assuming the electronic structure of the noble gas that precedes it in the periodic table.
• The charge of a cation formed by the loss of all valence electrons is equal to the group number minus 10.
• Exceptions to the expected behavior involve elements toward the bottom of the groups, which can form ions with a charge of 1+, 2+, or 3+.

Electronic Structures of Anions

• Most monatomic anions form when a neutral nonmetal atom gains enough electrons to completely fill its outer s and p orbitals, thereby reaching the electron configuration of the next noble gas.
• The charge on a negative ion is equal to the number of electrons that must be gained to fill the s and p orbitals of the parent atom.

Covalent Bonding

• Covalent bonds are formed when two atoms share electrons equally between each other.
• Covalent bonds are formed between two atoms when both have similar tendencies to attract electrons to themselves.
• Covalent compounds exhibit different physical properties than ionic compounds, such as lower melting and boiling points.
• Covalent compounds are generally poor conductors of electricity.

Formation of Covalent Bonds

• Nonmetal atoms frequently form covalent bonds with other nonmetal atoms.
• The potential energy of two separate atoms decreases as they approach each other, and the single electrons on each atom are shared to form a covalent bond.

Pure vs. Polar Covalent Bonds

• If the atoms that form a covalent bond are identical, the electrons in the bond are shared equally, forming a pure covalent bond.
• If the atoms are different, the bonding electrons are shared, but not equally, forming a polar covalent bond.
• In a polar covalent bond, the bonding electrons are more attracted to one atom than the other, giving rise to a shift of electron density toward that atom.

Electronegativity

• Electronegativity is a measure of the tendency of an atom to attract electrons (or electron density) towards itself.
• Electronegativity determines how the shared electrons are distributed between the two atoms in a bond.
• The more strongly an atom attracts the electrons in its bonds, the larger its electronegativity.Here are the study notes for the text:

Electronegativity and Bond Type

• Electrons in a polar covalent bond are shifted towards the more electronegative atom, resulting in a partial negative charge.
• The greater the difference in electronegativity, the more polarized the electron distribution and the larger the partial charges of the atoms.
• Electronegativity values increase from left to right across a period and decrease down a group in the periodic table.
• Nonmetals (upper right of the periodic table) tend to have the highest electronegativities, with fluorine being the most electronegative element (EN = 4.0).
• Metals tend to be less electronegative elements, with group 1 metals having the lowest electronegativities.
• The absolute value of the difference in electronegativity (ΔEN) provides a rough measure of the polarity of a bond and the bond type.

Bond Type and Electronegativity Difference

• Bonds with small or zero electronegativity difference are covalent and nonpolar.
• Bonds with large electronegativity difference are polar covalent or ionic.
• Examples of bond type and electronegativity difference:
  • H-H: 0 (nonpolar)
  • H-Cl: 0.9 (polar covalent)
  • Na-Cl: 2.1 (ionic)

Lewis Symbols and Structures

• Lewis symbols are used to describe valence electron configurations of atoms and monatomic ions.
• A Lewis symbol consists of an elemental symbol surrounded by one dot for each of its valence electrons.
• Lewis symbols can be used to show the formation of cations and anions from atoms.
• Lewis structures are used to indicate the formation of covalent bonds in molecules and polyatomic ions.
• The octet rule states that main group atoms tend to form enough bonds to obtain eight valence electrons.

Writing Lewis Structures

• Steps to write Lewis structures:
  1. Determine the total number of valence electrons in the molecule or ion.
  2. Draw a skeleton structure of the molecule or ion, arranging the atoms around a central atom.
  3. Distribute the remaining electrons as lone pairs on the terminal atoms (except hydrogen) to complete their valence shells with an octet of electrons.
  4. Place all remaining electrons on the central atom.
  5. Rearrange the electrons of the outer atoms to make multiple bonds with the central atom in order to obtain octets wherever possible.

Examples of Lewis Structures

• SiH4: 8 valence electrons, with 4 electrons on the central Si atom and 4 electrons on the H atoms.
• CHO2-: 18 valence electrons, with 6 electrons on the C atom, 6 electrons on the O atoms, and 2 electrons on the H atom.
• NO+: 10 valence electrons, with 5 electrons on the N atom and 5 electrons on the O atom.
• OF2: 20 valence electrons, with 6 electrons on the O atom and 14 electrons on the F atoms.### Chemical Bonding and Molecular Geometry

Writing Lewis Structures

• Steps to write Lewis structures:
  • Calculate the number of valence electrons
  • Draw a skeleton structure
  • Distribute electrons to terminal atoms
  • Place remaining electrons on the central atom
  • Rearrange electrons to form multiple bonds
• Examples: HCN, H3CCH3, HCCH, NH3

Exceptions to the Octet Rule

• Molecules with central atoms that do not have eight electrons:
  • Odd-electron molecules (free radicals)
  • Electron-deficient molecules
  • Hypervalent molecules
• Examples: NO, BeH2, BF3, PCl5, SF6

Formal Charges and Resonance

• Calculating formal charge:
  • Formal charge = # valence shell electrons (free atom) - # lone pair electrons - 1/2 # bonding electrons
• Guidelines for predicting molecular structure:
  • A molecular structure with all formal charges zero is preferred
  • The arrangement with the smallest nonzero formal charges is preferred
  • Lewis structures are preferable when adjacent formal charges are zero or of the opposite sign
  • The structure with the negative formal charges on the more electronegative atoms is preferred
• Examples: CO2, CNS-, NCS-, CSN-

Resonance

• If two or more Lewis structures with the same arrangement of atoms can be written for a molecule or ion, the actual distribution of electrons is an average of that shown by the various Lewis structures
• Resonance forms: individual Lewis structures
• Resonance hybrid: the actual electronic structure of the molecule (the average of the resonance forms)
• Examples: NO2-