Chapter 8 Molecular Compounds Molecular Compounds Covalent bondsAtoms

Chapter 8 Molecular Compounds

Molecular Compounds • Covalent bonds-Atoms held together by sharing electrons • Molecule- is a neutral group of atoms joined together by covalent bonds. – Ex. Air contains oxygen molecules, each oxygen molecule consist of two oxygen atoms joined by a covalent bond. • Diatomic molecule-is a molecule consisting of two atoms. – Ex. Oxygen molecule is a diatomic molecule

Molecular Compounds • Molecular compound-A compound composed of molecules • Molecular compounds tend to have relatively lower melting and boiling points than ionic compounds. • Many molecular compounds are liquids or gases at room temperature. • Ionic compounds are formed from a metal combined with a nonmetal; most molecular compounds are composed of atoms of two or more nonmetals. – Ex. CO carbon monoxide

Molecular Compounds • Molecular formula- is the chemical formula of a molecular compound • A molecular formula shows how many atoms of each element a molecule contains. – Water H 2 O CO 2 Ethanol C 2 H 6 O

Molecular Compounds • In forming covalent bonds electron sharing usually occurs so that atoms attain the electron configuration of noble gases. • Combinations of atoms of the nonmetallic elements in Groups 4 A, 5 A, 6 A, and 7 A of the periodic table are likely to form covalent bonds. Atoms usually acquire a total of eight electrons. • Single covalent bond-two atoms held together by sharing a pair of electrons – Ex. Hydrogen molecule

Molecular Compounds • An electron dot structure such as H: H represents the shared pair of electrons of a covalent bond by two dots. • Structural formula-represents the covalent bond by dashes and shows the arrangement of covalently bonded atoms. Ex. H-H Fluorine molecule F-F • Unshared pair- A pair of valence electrons that is not shared between atoms – Ex. Water molecule

Molecular Compounds • Atoms form double or triple covalent bonds if they can attain a noble gas structure by sharing two pairs or three pairs of electrons. • Double covalent bonds-A bond that involves two shared pairs of electrons – Ex. O=O • Triple covalent bonds-A bond formed by sharing three pairs of electrons – Ex. N=N

Molecular Compounds • Coordinate covalent bond-is a covalent bond in which one atom contributes both bonding electrons – • CO Polyatomic ion-a tightly bound group of atoms that has a positive or negative charge and behaves as a unit. – SO 3 -2 • Bond dissociation energy-The energy required to break the bond between two covalently bonded atoms • A large bond dissociation energy corresponds to a strong covalent bond • Resonance structure-is a structure that occurs when it is possible to draw two or more electron dot structures that have the same number of electron pairs for a molecule or ion. – Ex. O 3

Exceptions to the Octet Rule • The octet rule cannot be satisfied in molecules whose total number of valence electrons is an odd number. There also molecules in which an atom has fewer, or more, than a complete octet of valence electrons • A few atoms, especially phosphorus and sulfur sometimes expand the octet to include ten or twelve electrons. – Ex. PCl 5 or SF 6

Bonding Theories • Molecular orbitals-orbitals that apply to the entire molecule • Just as an atomic orbital belongs to a particular atom, a molecular orbital belongs to a molecule as a whole. • Bonding orbital- a molecular orbital that can be occupied by two electrons of a covalent bond.

Bonding Theories • Sigma bond-orbitals combine to form a molecular orbital that is symmetrical around the axis connecting two atomic nuclei • Pi bond – the bonding electrons are most likely to be found in sausage shaped regions above or below the bond axis of the orbitals.

Bonding Theories • Orbital hybridization provides information about molecular bonding and molecular shape. • Hybridization-several atomic orbitals mix to form the same total number of equivalent hybrid orbitals. pg. 234

VSEPR Model • Valence shell electron pair repulsion theory VSEPR theory- the repulsion between electron pairs causes molecular shapes to adjust so that the valenceelectron pairs stay as far apart as possible. • Geometric shape of molecule – Tetrahedral angle- 109. 5 degrees All of the H-C-H angles in CH 4 – Linear triatomic-180 degrees CO 2 – Trigonal planar-120 degrees C 2 H 4 – Bent triatomic- 105 degrees H 20 – Pyramidal- 107 degrees NH 3

Polar Bonds and Molecules • Nonpolar covalent bond-Atoms in the bond are shared equally in a bond. • This occurs when identical atoms are bonded together. Ex. H-H or O=O • Polar bond-is a covalent bond between atoms in which the electrons are shared unequally. • The more electronegative atom attracts electrons more strongly and gains a slightly negative charge. The less electronegative atom has a slightly positive charge. – Ex. H-Cl , H 2 O • Table on pg 238, and 177

Polar Bonds and Molecules • Polar molecule-one end of the molecule is slightly negative and the other end is slightly positive. • Dipole -is a molecule that has two poles is called a dipolar molecule. • When polar molecules are placed between oppositely charge plates, they tend to become oriented with respect to the positive and negative plates. – Ex. CO 2 linear form non polar molecule. • Intermolecular attraction are weaker than either ionic or covalent bonds.

Polar Bonds and Molecules • Van der Waals forces-the two weakest attractions between molecules. – 1. Dipole interactions-attachments of polar molecules – 2. Dispersion forces-Caused by motion of electrons, momentary polar force • Hydrogen bonds-attractive forces in which a hydrogen is covalently bonded to a very electronegative atom, and also weakly bonded to an unshared electron pair of another electronegative atom. • Network solids-a solid in which all of the atoms are covalently bonded to each other. – ex. Diamond is a network solid • Melting point very high if at all. Diamond doesn’t melt at all it vaporizes to gas at 3500 degrees C. • Melting a network solid requires breaking covalent bonds throughout the solid. • Table 8. 4 pg 244 Comparison of ionic and covalent compounds.