Chapter 6 CHEMICAL BONDING

CHEMICAL BOND A mutual electrical attraction between the nuclei and valence electrons of DIFFERENT atoms that BINDS the atoms together. Chemical bonds form to LOWER POTENTIAL ENERGY thus becoming more stable.

Three types of chemical bonding- IONIC- bonding that results from the electrical attraction between cations and anions. COVALENT- bonding that results from the sharing of electron pairs between two atoms. METALLIC- The chemical bonding that results from the attraction between metal atoms and the surrounding sea of electrons.

COVALENT BONDING Non-polar covalent bond- Covalent bond in which the electrons are SHARED EQUALLY by the bonding atoms. Polar covalent bond- covalent bond in which the bonded atoms have an UNEQUAL SHARING of electrons. See figure 6-3 pg 163.

To determine if a bond is ionic or covalent the difference between the electronegativities of the elements is determined. If the difference is from: 0 - 0.3 = nonpolar covalent. 0.4 – 1.7 = polar covalent. over 1.7 = ionic.

Electrons in a polar covalent bond sit closer to the MORE electronegative atom. What type of bond is formed when hydrogen bonds with carbon? Hydrogen with chlorine? Hydrogen with oxygen?

HOMEFUN- PAGE 163 Q 1-4 Page 195 q 1, 2, 4, 5, 33 Ch 6-1 review sheet Electronegativity table is on page 151.

Covalent Bonding and Molecular Compounds Ch 6-2

MOLECULES- A neutral group of atoms held together by covalent bonds MOLECULES- A neutral group of atoms held together by covalent bonds. MOLECULAR COMPOUND- A chemical compound whose simplest units are molecules.

CHEMICAL FORMULAS- show the relative types and numbers of atoms in a chemical compound. (Not limited to covalently bonded atoms.) MOLECULAR FORMULAS- show the types and numbers of atoms found in a single molecule. DIATOMIC MOLECULES- contain only of TWO atoms, of either alike or different elements. Ex: H2, N2, CO

Why do chemical bonds form? To lower the potential energy. When two atoms are separated, they do not affect each other, but as they are drawn together by attractive forces the potential energy decreases. When the attractive force balances the repulsive force, the P.E. is at a minimum. This is the most stable configuration for the atoms. If the atoms move too close the P.E. increases because of the repulsive forces of the like charges.

Characteristics of Covalent Bonds Bond Length- the average distance between two bonded atoms at their lowest P.E. (1 pm= 10-12 m) Bond Energy- the amount of energy it takes to break the bond, which is exactly the same as the amount of energy that is released when the bond is formed. (kJ/mol) Bond lengths and bond energies vary with the types of atoms that have combined, and if the atoms have formed other bonds.

THE OCTECT RULE (finally) Chemical compounds tend to form so that each atom, by gaining, losing or sharing electrons, has an octet of electrons in its highest energy level. This means that both the s and p sublevels at the highest energy level will be filled, giving the atoms a noble- gas configuration.

EXCEPTIONS TO THE OCTET RULE You knew that was coming There are two ways that atoms will not be able to obey the octet rule. These are: 1. Having too few valence electrons to ever obtain an octet. Hydrogen can have at most 2 valence electrons after it shares its electron with another atom. Beryllium will have 4 valence electrons after it has finished bonding. Boron will have 6 valence electrons after it shares its valence electrons with other atoms.

2. Expanding the octet to have 10, 12 or 14 valence 2. Expanding the octet to have 10, 12 or 14 valence electrons instead of 8. elements in periods 3, 4, 5, 6 and 7 can expand their octet to have 10, 12, or 14 valence electrons.

ELECTRON DOT NOTATION – an electron-configuration in which only the valence electrons of an atom are shown and indicated by dots around the element’s symbol. ∙H ∙He∙

A shared pair of electron dots are often represented by a dash A shared pair of electron dots are often represented by a dash. H:H  H-H Lone pair electrons- a pair of electrons that are not involved in bonding and belong exclusively to one atom. Aka unshared pair.

Lewis Structures- formulas in which atomic symbols represent nuclei and inner-shell electrons, dot-pairs or dashes represent electron pairs in covalent bonds, and single dots represent unshared electrons.

Multiple Bonds – double and triple bonds Double bonds- Two pair of electrons are shared. Generally have higher bond energy and shorter bond length than single bonds. Triple Bonds- Three pair of electrons are shared. Generally have higher bond energy and shorter bond length than either double or single bonds . Turn to page 171 in your text Homefun: pg 175 q 1, 3, 4 pg 195 q 6, 7, 14

Ionic Bonding and Ionic Compounds 6-3 Ionic Bonding and Ionic Compounds

Ionic Bonding and Ionic Compounds Ionic Compound – composed of positive and negative ions that are combined so that the number of positive and negative charges are equal. Ex: NaCl Formula Unit – the simplest collection of atoms from which an ionic compound’s formula can be established. The ratio of ions in a formula unit depends on the charges of the ions involved. EX: Calcium (carries a 2+ charge) and fluorine (1- charge) combine to form a neutral compound  CaF2

Characteristics of Ionic Bonding All ionic compounds form crystals, a Crystal Lattice formation that minimizes PE. The distance between the + and – ions in a compound represents a balance between the repulsive and attractive forces, just like in molecular bonding. The arrangement that these ions like to stack into depends on the types and charges of the ions involved. LATTICE ENERGY- the energy released when one mole of an ionic compound is formed from gaseous ions.

Comparison Ionic and Molecular compounds . Comparison Ionic and Molecular compounds Even though the bonds formed in the formation of molecular compounds are relatively strong, the forces between molecules are much, much weaker than the forces between ions in an ionic compound. This means that ionic and molecular compounds have different properties

Ionic compounds tend to have high melting and boiling points. Have you ever tried to melt salt?…good luck. Sugar, on the other hand will melt easily. Ionic compounds are hard but brittle. Since they “stack” they move in rows, causing an increase in repulsive forces across the layer, so the layers separate. Molten ionic compounds conduct electricity because the ions are free to carry the electrical energy. Many ionic compounds conduct electricity when dissolved in water. The water molecules separate the positive and negative ions and can move freely.

POLYATOMIC IONS Polyatomic ions are charged molecules. The group of atoms as a whole carry a charge. Polyatomic ions combine with other ions of opposite charge. Examples of polyatomic ions- NH4+, NO3-, NO2- SO42- , CN- , PO43- Practice: Draw the Lewis structures for these polyatomic ions. Common

Ex: ozone- O3 O=O-O   O-O=O Resonance Structures (aka resonance hybrids)- the bonding in molecules or ions that cannot be represented by a single structure. Ex: ozone- O3 O=O-O   O-O=O (add lone pairs to complete the structures) Scientists have found that the bonds are identical between the oxygen atoms. A double arrow between these structures indicates resonance.

Home FUN: pg 180 q 1-4 pg 197 q. 38, 39, 41, 42

Chapter 6 Section 4 METALLIC BONDING- The chemical bonding that results from the attraction between metal atoms and the surrounding sea of electrons.

The atoms are held in a relatively fixed position while the valence electrons are free to move throughout. Electrons in metals can have vacant p and d orbitals. Since the orbitals between metal atoms can overlap, the valence electrons are free to move within the entire structure. The valence electrons DO NOT belong to any one atom, but rather form a sea of electrons.

Metallic Bond Strength Metallic Properties Metallic bonding accounts for many physical properties of metals, such as malleability, ductility, thermal and electrical conductivity, and luster. Metallic Bond Strength Varies with the nuclear charge of the atom and the number of electrons in the sea. Heat of Vaporization- The amount of energy required to vaporize a metal.

Home Fun Pg 182 q 1-3

Chapter 6 Section 5 Molecular Geometry The three dimensional arrangement of a molecule’s atoms in space.

VSEPR THEORY - a model used for predicting the shapes of individual molecules. Valence shell electron pair repulsion. Pairs of electrons in the valence shell of a central atom repel each other. These pairs of electrons tend to occupy positions in space that minimize repulsions and maximize the distance between them. Unshared pairs also occupy space and must be considered .

ABE method of determining geometry. A- represents the central atom. B- represents the atoms bonded to the central atom. E- represents an unshared electron pair. Determine the Lewis structure of the molecule and count the numbers of each of the above.

B=3 E = 1 Ammonia- NH3 to central atom Atoms bonded Lone pair electrons to central atom

Try this one- H2O # of bonded atoms # of lone to central atom pairs

HYBRIDIZATION Ex:Carbon 1s22s22p2  1s22sp3 s,p  sp s,p,p sp2 The mixing of two or more atomic orbitals of similar energies (same energy level) on the same atom to produce new orbitals of equal energies. Ex:Carbon 1s22s22p2  1s22sp3 s,p  sp s,p,p sp2 s,p,p,p  sp3

Intermolecular Forces- The force of attraction between molecules Intermolecular Forces- The force of attraction between molecules. Generally weaker than covalent, ionic and metallic bonds. 1. dipole-dipole forces 2. hydrogen bonding- a particular type of dipole- dipole. 3. London dispersion forces Boiling point is an indication of the strength of the intermolecular forces between particles in a liquid.

Molecular Polarity DIPOLE-DIPOLE Forces - The negative area of one molecule attracts the positive region of another molecule. How do we show this? A dipole in a molecule is represented by an arrow directed from positive to negative.

This determines if a molecule is polar or non-polar In ammonia, the bond polarities are additive, causing the molecule as a whole to be polar In carbon dioxide, the bond polarities extend equally in different directions, canceling each others effects This determines if a molecule is polar or non-polar

Dipoles can be induced in non-polar molecules by temporarily attracting it’s electrons These induced dipoles are weaker than the dipole-dipole force.

Hydrogen Bonding- a special case of dipole forces. A hydrogen bond is the attractive force between a hydrogen atom that is bonded to a highly electronegative atom and the unshared electrons of an atom in a nearby molecule. Usually the electronegative atom is oxygen, nitrogen, or fluorine . hydrogen bonds link

London Dispersion Forces- intermolecular attractions that result from the constant motion of electrons and the creation of temporary dipoles. the weakest intermolecular force. Act between all atoms and molecules, including noble-gases. a temporary attractive force that results when the electrons in two adjacent atoms occupy positions that make the atoms form temporary dipoles. London forces are the attractive forces that cause nonpolar substances to condense to liquids and to freeze into solids when the temperature is lowered sufficiently.

Because the electrons from different molecules start "feeling" and avoiding each other, electron density in a molecule becomes redistributed in proximity to another molecule Instantaneous dipoles: These occur due to chance when electrons happen to be more concentrated in one place than another in a molecule, creating a temporary dipole. Induced dipoles: These can occur when one molecule with a permanent dipole repels another molecule's electrons, "inducing" a dipole moment in that molecule.

Home fun: pg 193 q 5, pg 198 q 51, 53- 55, 58, 63, 64