Pre AP Chapter 6 Notes

A chemical bond is a mutual electrical attraction between the nuclei and valence electrons of different atoms, and binds those atoms together.

There are three types of bonding: ionic, covalent, and metallic

There are three types of bonding: ionic, covalent, and metallic

1.Ionic bonding results from the transfer of valence electrons from one atom to another.

There are three types of bonding: ionic, covalent, and metallic

1.Ionic bonding results from the transfer of valence electrons from one atom to another.

2.Metallic bonding results from a sharing of very weakly held valence electrons

There are three types of bonding: ionic, covalent, and metallic

1.Ionic bonding results from the transfer of valence electrons from one atom to another.

2.Metallic bonding results from a sharing of very weakly held valence electrons

3.Covalent bonding results from the sharing of valence electrons, whether equally or unequally.

The type of bonding between two atoms can be estimated by calculating the difference in the elements’ electronegativity.

The type of bonding between two atoms can be estimated by calculating the difference in the elements’ electronegativity. To determine the type of bond, if the electronegativity difference is:

0-.3 then the bond is nonpolar covalent

>.3 - 1.7 then the bond is polar covalent

>1.7 then the bond is ionic

Nature favors chemical bonding because most atoms are at lower potential energy when bonded to other atoms than when they are independent particles.

Nature favors chemical bonding because most atoms are at lower potential energy when bonded to other atoms than when they are independent particles. The bond length is the average distance between two bonded atoms.

Nature favors chemical bonding because most atoms are at lower potential energy when bonded to other atoms than when they are independent particles. The bond length is the average distance between two bonded atoms. Bond energy is the energy required to break a chemical bond and form neutral, isolated atoms. The unit for bond energy is in kilojoules per mole (kJ/mol).

An ionic compound is composed of positive and negative ions (cations and anions) that are combined in the right ratio to that the total of the positive and negative ions are zero.

Most ionic compounds are crystalline solids. A crystal is a three-dimensional network of positive and negative ions.

Most ionic compounds are crystalline solids. A crystal is a three-dimensional network of positive and negative ions. A formula unit is the simplest collection of atoms from which the ionic compound’s formula can be established. In an ionic crystal, ions minimize their potential energy by combining in an orderly arrangement known as a crystal lattice.

Most ionic compounds are crystalline solids. A crystal is a three-dimensional network of positive and negative ions. A formula unit is the simplest collection of atoms from which the ionic compound’s formula can be established. In an ionic crystal, ions minimize their potential energy by combining in an orderly arrangement known as a crystal lattice. Lattice energy is the energy released when one mole of an ionic crystalline compound is formed from gaseous ions.

The forces of attraction between compounds and molecules are much weaker than the forces inside holding the compound or molecule together. In addition, the forces holding ionic compounds together are much stronger than the forces holding covalent molecules together. This accounts for the high melting and boiling points of ionic compounds.

Generally ionic compounds do not vaporize at room temperature, and they are hard but brittle. They are poor conductors of electricity in their solid state but are good conductors in their liquid state. Most ionic compounds are easily dissolved and will also be good conductors in their aqueous form.

A polyatomic ion (radical) is a charged group of covalently bonded atoms. The charge of a polyatomic ion results from an overall excess of electrons (negative charge) or a shortage of electrons (positive charge). Polyatomic ions will form ionic bonds with other charged elements or other polyatomic ions to form a neutral compound.

Metallic bonding accounts from the unique properties of metals. Their high conductivity is due to the highly mobile valence electrons (often referred to as a sea of electrons).

Metallic bonding accounts from the unique properties of metals. Their high conductivity is due to the highly mobile valence electrons (often referred to as a sea of electrons). These electrons roam freely throughout the atom and are called delocalized.

Metallic bonding accounts from the unique properties of metals. Their high conductivity is due to the highly mobile valence electrons (often referred to as a sea of electrons). These electrons roam freely throughout the atom and are called delocalized. This also accounts for the luster of metals because the electrons absorb light, jump to a higher energy level, and then fall back down to ground level while emitting the energy in the form of light.

Metallic bonding accounts from the unique properties of metals. Their high conductivity is due to the highly mobile valence electrons (often referred to as a sea of electrons). These electrons roam freely throughout the atom and are called delocalized. This also accounts for the luster of metals because the electrons absorb light, jump to a higher energy level, and then fall back down to ground level while emitting the energy in the form of light. The ability of the metal nuclei to “ride over” the sea of electrons helps account for the ductility and malleability of metals.

The metallic bond strength varies with the nuclear charge of the metal and the number of electrons in the sea. Both of these factors are reflected in the heat of vaporization which is the amount of heat required to vaporize one mole of metal.

Covalent bonding results from the sharing of electrons in pairs between two nuclei. If two atoms share the electron pair equally, the bond is nonpolar covalent. If the two atoms have an unequal sharing it is a polar covalent bond. Again, the type of boding is determined by the difference in the electronegativities of the atoms.

To indicate the positive end of a bond, “δ+” is used and for the negative end of the bond, “δ-” is used. For example, in the bond between H and Cl, the bond is polar covalent because hydrogen’s electronegativity is 2.1 and chlorine’s is 3.0, so the difference is 0.9. We would show this as:

δ+ δ-

H - Cl

A molecule is a neutral group of covalently bonded atoms. It is capable of existing on its own (does not have a crystal lattice). A chemical formula or molecular formula indicates the relative numbers of atoms of each kind in a chemical compound by using symbols for the elements and subscripts. A molecule of only two elements is called a diatomic molecule.

H2OH2O2NO2

N2O5

To draw compounds and molecules, the valence electrons act as a guide using the Octet Rule. Elements form compounds and molecules with gaining, losing, or sharing electrons to have an octet of valence electrons (eight), filling up the s and p sublevels to give a noble gas configuration and stability.

For ionic compounds show the electrons being lost from the metal and gained by the nonmetal. Place a bracket around the newly formed cations and anions with the charge outside.

Examples:

K+ [ O ]2- K+ and [ I ]- Ba2+ [ I ]-

Notice it is the negative nonmetal that gets all the dots!

.. .. ..

.. .. ... . .. ... .. . . .

For covalent molecules, there is a simple method that can be used to determine how many bonds and how many lone pairs should be around each atom. This method is the W - A = S method

W = the total number of electrons all the atoms want

(each atoms wants 8, but hydrogen only wants 2)

A = the total number of electrons all atoms have available

(count the dots off the periodic table, add for each - charge and subtract for each + charge)

S = the total number of electrons that need to be shared

(divide by 2 for the number of bonds)

Examples:SO2

PO33-

C2H2

Limitations to the octet rule (cases in which the octet rule may not apply):

1.Most covalent compounds of Be

2.Most covalent compounds of Group 3A (13)

3.Compounds in which the central atom must share more than 8 valence electrons to accommodate all the attached atoms (like SF6)

4.Compounds containing d or f metals

5.Species with odd numbers of electrons (like NO with A = 11 electrons)

Unshared electron pairs are called lone pairs and belong to one atom only. Lewis structures are these structures in which atomic symbols represent nuclei, dashes represent shared pairs, and dots represent lone pairs.

Unshared electron pairs are called lone pairs and belong to one atom only. Lewis structures are these structures in which atomic symbols represent nuclei, dashes represent shared pairs, and dots represent lone pairs. Structural formulas indicate the kind, number and arrangement, and bonds of the atoms, but not the unshared lone pairs.

A single bond is produced by sharing one pair of electrons between two atoms. A double bond is produced by sharing two pairs of electrons between two atoms. A triple bond is produced by sharing three pairs of electrons between two atoms.

Double and triple bonds are referred to as multiple bonds. Resonance can occur with multiple bonds, and refers to situations where more than one structure can be drawn.

Example:SeO2

In real life molecules form three dimensional structures that are not as simple as the two dimensional structures above. Thus it takes a new theory - Valence Shell, Electron Pair Repulsion, or VSEPR to tell us how to draw the structures three dimensionally.

In real life molecules form three dimensional structures that are not as simple as the two dimensional structures above. Thus it takes a new theory - Valence Shell, Electron Pair Repulsion, or VSEPR to tell us how to draw the structures three dimensionally. This theory states that the repulsion between the sets of valence level electrons surrounding the atom causes these sets to be arranged as far apart as is possible.

In real life molecules form three dimensional structures that are not as simple as the two dimensional structures above. Thus it takes a new theory - Valence Shell, Electron Pair Repulsion, or VSEPR to tell us how to draw the structures three dimensionally. This theory states that the repulsion between the sets of valence level electrons surrounding the atom causes these sets to be arranged as far apart as is possible. This could lead to molecular polarity where there is an uneven distribution of molecular charge.

There are several forces that act between compounds and molecules to hold them together (not referring to the ones inside the compounds and molecules). These forces are called Intermolecular forces (as opposed to Intramolecular forces).

There are three types of intermolecular forces:

1.Hydrogen bonding - the attraction of a hydrogen on one polar molecule to the lone pair of a highly electronegative atom (N, O, F) of a second polar molecule.

Click here for a demo of hydrogen bonding in water.

2.Dipole-dipole - a dipole is created by equal but opposite charges that are separated by the short distance of a bond. The positive end of one dipole molecule will be attracted to the negative end of a second dipole molecule, and vise versa.

3.London dispersion forces - resulting from the constant motion of electrons and the creation of instantaneous dipoles.

Intramolecular Forces

Covalent Network

Ionic

Metallic

Intermolecular Forces

Hydrogen bonding – polar with hydrogen

Dipole-dipole –polar onlyLondon –nonpolar