Unit 3.3: Covalent Bonds and Intermolecular Forces Keep until June 2011! Unit 3.3: Covalent Bonds and Intermolecular Forces

Vocabulary: Covalent bond: “co-” = with; “-valent-” = valence (electrons); bond formed when two atoms share valence electrons Molecule: a neutral group of atoms joined by covalent bonds. Molecular formula: a chemical formula that shows the actual number of each type of atom that participates in the compound. Structural formula: a drawing that represents the bonds of a molecular compound as lines connecting the atoms. Nonpolar: a molecule that has no regions of partial charge Polar: a molecule that has regions of partial charge Dipole: attraction between two polar molecules Hydrogen bond: bond formed between the slightly positively charged hydrogen on one compound to an unshared pair of electrons on a separate molecule

I. Molecular Compounds A molecule (molecular compound) is a neutral group of atoms joined by covalent bonds. A diatomic molecule is composed of two atoms of the same element sharing valence electrons. Br I N Cl H O F series… A molecular compound is a compound formed when two or more different elements share valence electrons. Often, refer to both as simply “molecules”; ex: a molecule of oxygen, a molecule of water

The molecular formula of a compound is the chemical formula that shows the numbers of each type of atom participating in the compound. *Subscripts are not necessarily the smallest whole number ratio like ionic compounds Ex: C6H12O6: 6 atoms of C, 12 atoms of H, and 6 atoms of O, not CH2O A molecular formula does not show arrangement of atoms.

The structural formula of a compound shows the arrangement of the atoms that are participating in the compound in relation to each other.

II. Covalent Bonds A covalent bond forms when two atoms share valence electrons to achieve the stability of a noble gas “co-” = with; “valent” refers to valence electrons Satisfies the octet rule for both atoms (Ionic bonds involve the transfer of an electron between atoms, creating charged atoms that bond through electrostatic attraction.)

**Covalent bonds usually form between nonmetals in groups 4A7A**

When atoms share a single pair of electrons, forms a single covalent bond: Represented by a single line in a structural formula Or by a pair of dots in an electron dot structure

When atoms share two pairs of electrons, forms a double covalent bond: Represented by a double line in a structural formula (Or by 2 pairs of dots in an electron dot structure) Stronger than a single bond, pulls atoms closer together b/c greater interaction between nuclei

When atoms share three pairs of electrons, forms a triple covalent bond: Represented by a triple line in a structural formula (Or by 3 pairs of dots in an electron dot structure) Even stronger than a double bond, pulls atoms even closer together b/c greater interaction between nuclei

Resonance: when there is more than one possible valid dot structure for a compound, the true structure “resonates” between all valid possibilities Ex: In a molecule of ozone (O3), a central oxygen molecule shares a single bond with one oxygen but a double bond with a third oxygen. The double bond is constantly shifting between both external oxygen molecules.

III. Polarity of Bonds Nonpolar covalent bond: occurs when covalent electrons are shared equally in a bond All diatomic molecules are nonpolar (H2, O2, Cl2, etc…) Some compounds are nonpolar as well: CO2, CH4

Polar covalent bonds: electrons are not shared equally between atoms, based on electronegativity Electrons are “pulled” towards more “electronegative” elements, which become slightly negatively charged. Electrons are “pulled” away from less “electronegative” elements, which become slightly positively charged. Partial charges in a molecule are denoted with Greek lower-case “delta” (δ) Can also represent polarity by an arrow pointing to the more electronegative element in a polar molecule

Electronegativity predicts which bond will form: If the difference in electronegativity between the two elements is: 0.0 to 0.4, the bond is nonpolar covalent 0.41 to 1.0, the bond is moderately polar covalent 1.01-2.0, the bond is very polar covalent ≥ 2.01, the bond is ionic

Properties of polar molecules: Compounds with polar bonds are often polar (but not always), have two poles, so they are called “dipoles” Analogy: north/south poles of a magnet – two poles B/c they have regions of charge, they can interact with other charged particles. But if the bonding atoms lie in the same plane, the polar bonds can “cancel” each other out, so the compound itself is nonpolar. Ex: CCl4: Cl is more electronegative, pulls e-’s away from C, but in opposite directions equally Keep until June 2011! Z

IV. Intermolecular Forces (forces between molecules) Van der Waals forces: weakest forces, two different types: Dispersion forces: can even occur between nonpolar molecules b/c as electrons move, they may induce a momentary region of charge that can attract a nearby molecule Sometimes called London forces Dipole-dipole interactions: polar molecules are attracted to each other b/c of opposite partial charges Similar to ionic interactions, but much weaker Dipole Force Dispersion (London) Force

Hydrogen Bonds: dipole interactions that always involve hydrogen atoms Hydrogen bonded to very electronegative atom can bond weakly to unshared pair of electrons of a separate molecule Hydrogen bonds are responsible for keeping water a liquid at room temperature. Hydrogen bonds are critical in many biological molecules Hold the “rungs” of the DNA ladder together Strongest of all intermolecular forces

Intermolecular forces Comparing Bonds Ionic > covalent > hydrogen bonds > dipole> dispersion forces Intramolecular forces Intermolecular forces Characteristic Ionic Compounds Covalent Compounds Representative Unit Formula unit Molecule Bond formation Transfer of one or more electrons between atoms Sharing of electron pairs between atoms Types of elements Metals with nonmetals Nonmetals only Physical state of matter Solid Solid, liquid, or gas Melting Point High (usually greater than 300° C) Low (usually below 300° C) Solubility in water Usually high High to low Electrical conductivity when dissolved in water Good conductor Poor or nonconducting

Electronegativity Values for Selected Elements H 2.1 Li 1.0 Be 1.5 B 2.0 C 2.5 N 3.0 O 3.5 F 4.0 Na 0.9 Mg 1.2 Al Si 1.8 P S Cl K 0.8 Ca Ga 1.6 Ge As Se 2.4 Br 2.8 Rb Sr In 1.7 Sn Sb 1.9 Te I Cs 0.7 Ba Tl Pb Bi