Chapter 6

 Chemical bonds are formed when atoms combine to become more stable.  Types of bonds  Ionic  Molecular  Metallic

 The level of bonding depends on how strongly the atoms of each element attract electrons.  This is the electronegativity-ability to attract electrons.  The difference (subtract) in electronegativity between the 2 atoms determines if the compound is ionic, strongly polar/molecular, moderately polar/molecular, or nonpolar molecular.

Ionic Bonds Cations lose electrons. Anions gain electrons. Both want a noble gas configuration. Electronegativity difference is high because one atom holds electrons strongly (anion) while other does not(cation). One completely gains e-; other completely loses e-. When atoms gain or lose electrons, they become ions. Forces between ions form ionic bonds.

Properties of Ionic Bonds Ionic bonds form between a cation and an anion. (metal and nonmetal). They are formed by the transfer of electrons. Cations lose electrons while anions gain electrons to form a stable octet.  Ionic bonds are strong. Nearly all ionic compounds are crystalline solids, with orderly structures.  High melting and boiling point because hard to break apart.  Total positive charge is balanced by total negative charge, so electrically neutral (crystal lattice.)  Easily dissolves in water.  In solution, conducts electricity.

The resulting compound is unique both chemically and physically from its parent atoms. compound y/flash_viewer.php?oid=1349&mid=55

Ionic compounds are solids due to forces between molecules Each sodium ion is attracted equally to all of its neighboring chlorine ions, and likewise for the chlorine to sodium attraction.ions The concept of a single molecule becomes blurred in ionic crystals because the solid exists as one continuous system.crystals Forms crystal lattice with high mp and bp.

Covalent Bonding *Occurs when 2 or more elements share electrons. *Usually between 2 or more nonmetals. *Both want to gain electrons so they will share to fill their outer shell.   They are not held together as strongly, so they are often found as liquids and gases.  Unlike ionic, found as a true molecule.

IONICCOVALENT  Metal and nonmetal  Crystal lattice  High melting and boiling point.  Solids  Completely gain and lose electrons  Conduct electricity in water  High electronegativity difference.  2 nonmetals  Molecule  Low melting and boiling point.  Liquids and gases (or brittle solids)  Share electrons  Do not conduct electricity.  Similar electronegativities

POLARNONPOLAR  Uneven distribution of charges.  Bonded atoms have an unequal attraction for the shared electrons.  The compound has a slight positive end and slight negative end. (dipoles)  Due to electronegativity differences.  Electrons are shared equally among the atoms.  Equal attraction of electrons.  Not charged.  Similar electronegativities.  Examples: diatomic molecules.

Lewis Dot Structures Lewis dot structures are a shorthand way to represent the valence electrons of an atom. valence electronsatom The structures are written as the element symbol surrounded by dots that represent the valence electrons. elementelectrons

 The bonding electrons are placed between the atoms and can be represented by a pair of dots or a dash (each dash represents one pair of electrons, or one bond). Lewis structures for H 2 and O 2 are shown below.electrons atoms electrons

Bond Length—The average distance between 2 bonded atoms. Longer bonds tend to be weaker. Bond Energy—The energy required to break a chemical bond and form neutral isolated atoms. More energy needed if shorter bond length.

 1. Determine the type and number of atoms in the molecule. (CH 3 I)  2. Write the electron-dot notation for each type of atom in the molecule.  3. Determine the total number of valence electrons available in the atoms to be combined.

 4. Arrange the atoms to form a skeleton structure for the molecule. If carbon is present, it is the central atom. Otherwise, the least- electronegative atom is central (except for hydrogen, which is never central). Then connect the atoms by electron-pair bonds.  5. Add unshared pairs of electrons to each nonmetal atom (except H) such that each is surrounded by eight electrons.  6. Count the electrons in the structure to be sure that the number of valence electrons used equals the number available. Be sure the central atom and other atoms besides H have an octet.

 1. NH 3  2. H 2 S  3. SiH 4  4. PF 3

 Double bonds—share 2 pairs of electrons.  C 2 H 4  Triple bonds—share 3 pairs of electrons.  N 2 or C 2 H 2

 Bonding in molecules or ions that cannot be correctly represented by a single Lewis Structure.  See handout  More Practice: p. 189 #4

 The chemical bonding that results from the attraction between metal atoms and the surrounding sea of electrons.  There are usually a lot of empty d orbitals in metals.  The electrons are delocalized, which means they do not belong to any one atom but move freely about the metal’s network of empty orbitals.  Sea of electrons form a crystal lattice.

Because of the sea of electrons in metals: They have luster.(shiny-reflect light) They conduct electricity. They are malleable.(hammered or beaten into sheets) They are ductile.(drawn into wires) When struck, one plane of atoms in a metal can slide past another without encountering resistance or breaking bonds.

 Repulsions between the sets of valence-level electrons surrounding an atom causes these sets to be oriented as far apart as possible.

 1. Dipole-Dipole Forces—between polar molecules because of their uneven charge distribution.  One dipole’s positive end is attracted to another’s negative end.  A polar molecule can induce a dipole in a nonpolar molecule by temporarily attracting its electrons. (weak force)

 2. Hydrogen Bonding—A hydrogen atom that is bonded to a highly electronegative atom is attracted to an unshared pair of electrons of an electronegative atom in a nearby molecule.  3. London Dispersion Force—Results from the constant motion of electrons and the creation of instantaneous dipoles.  In all atoms and molecules (even noble gases)