PAP Chemistry

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

A Lewis Structure is formula in which atomic symbols represent nuclei and inner-shell electrons, dot-pairs or dashes between two atomic symbols represent electron pairs in covalent bonds, and dots adjacent to only one atomic symbol represent unshared electrons.

Ionic, Covalent, and Metallic Bonding

Ionic Bonding – NaCl

1.  Ionic bonds are electrostatic attractions between large numbers of cations and anions. Cations (+ particles) adhere to anions (- particles) because of opposite charges. One of the particles exhibits enough electronegativity to “pull” an electron from another particle. Thus cations and anions are formed and an ionic bond is created.

This type of bond is found in all salts and a crystalline lattice structure is made when the following bonds are created:

§  A metallic cation and a non-metallic anion

§  A metallic cation and a polyatomic anion

§  A polyatomic cation (ammonium is the only one we know) and a non-metallic anion

§  A polyatomic cation and a polyatomic anion

  1. Covalent bonds are formed when atoms share the electron pair between the two atoms. This occurs because neither of the particles is electronegative enough to be able to actually take an electron away from the other (they are similar in their ability to pull bonded electrons toward themselves). Therefore the electrons are simply placed between the two atoms and are shared. Because of differences in electronegativity values the sharing of electrons will not always be equal.

Covalent Bonding

Metallic Bonding

  1. Metallic bonds are found only in pure metal elements and consist of positively-charged ions surrounded by a “sea” of electrons. This “sea” of electrons is free to move throughout the entire sample. This freedom of movement of the “sea” of electrons gives metals their basic physical characteristics such as malleability, ductility, electrical, and heat conductivity.

Bond Polarity

The degree to which bonding between atoms of two elements is ionic or covalent can be estimated by calculating the difference in the elements’ electronegativities. The electronegativity table is on page 151 in your textbook.

Use the electronegativity table values to determine bond polarity. See chart below.

Electronegativity difference / Type of bond
> 1.7 / Ionic
1.7 – 0.3 / Polar covalent
<0.3 / Non-polar covalent

Electronegativity Table

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
1.5 / Si
1.8 / P
2.1 / S
2.5 / Cl
3.0
K
0.8 / Ca
1.0 / Sc
1.3 / Ti
1.5 / V
1.6 / Cr
1.6 / Mn
1.5 / Fe
1.8 / Co
1.9 / Ni
1.9 / Cu
1.9 / Zn
1.6 / Ga
1.6 / Ge
1.8 / As
2.0 / Se
2.4 / Br
2.8
Rb
0.8 / Sr
1.0 / Y
1.2 / Zr
1.4 / Nb
1.6 / Mo
1.8 / Tc
1.9 / Ru
2.2 / Rh
2.2 / Pd
2.2 / Ag
1.9 / Cd
1.7 / In
1.7 / Sn
1.8 / Sb
1.9 / Te
2.1 / I
2.5
Cs
0.7 / Ba
0.9 / La-Lu
1.0-1.2 / Hf
1.3 / Ta
1.5 / W
1.7 / Re
1.9 / Os
2.2 / Ir
2.2 / Pt
2.2 / Au
2.4 / Hg
1.9 / Tl
1.8 / Pb
1.9 / Bi
1.9 / Po
2.0 / At
2.2
Fr
0.7 / Ra
0.9 / Ac
1.1 / Th
1.3 / Pa
1.4 / U
1.4 / Np-No
1.4-1.3

Sample Problem 1:

Determine the bond polarity between sulfur and the following elements, hydrogen, H; cesium, Cs; and chlorine, Cl. In each pair, which atom will be more negative?
Bonding between Electronegativity Bond More-negative
sulfur and difference Type atom
hydrogen
Cesium
Chlorine

Practice Problem 1a:

Use electronegativiy differences to classify bonding between chlorine, Cl, and the following elements: calcium, oxygen, and bromine. Indicate the more-negative atom in each pair.
Bonding between Electronegativity Bond More-negative
Chlorine and difference type atom
Calcium
Oxygen
bromine

Using the diagrams above, discuss the how polarity affects physical properties such as boiling point and solubililty.

Lewis Structures

To effectively discuss bonding, we need to draw Lewis Structures

§  The Lewis Dot Structures show valence electrons. Valence electrons are the s and p cloud electrons. Groups 1 (alkali metals) and 2 (alkaline earth metals) will show only s cloud electrons. Groups 13 through 17 will show s cloud and p cloud electrons. The Noble Gases have full octet of valence electrons and therefore do not form bonds. Xenon is an exception to this rule and actually forms compounds with fluorine (the most highly electronegative element).

§  Elements in period 3 have the capacity to shift electrons into the d orbital and are able to form bonds exceeding the normal octet. The octet consists of 8 valence electrons which form the bond between two atoms or ions. Elements in period 3, p block may form bonds with 10 or more electrons when they exceed the octet. This allows for very interesting shapes of molecules.


Write formulas for the compounds above. Note the periodic table location of the central atoms and how this affects the bonding ability of the atoms. Note how location on the periodic table helps to determine the number of bonding sites and thus shapes of the molecules.

Draw Lewis Dot Structures (electron-dot notation) for the following atoms:

1.  lithium 4. hydrogen 7. magnesium

2.  boron 5. neon 8. carbon

3.  oxygen 6. fluorine 9. nitrogen

Electron-dot notation can also be used to represent molecules.

Hydrogen molecule, H2
Fluorine molecule, F2

The pair of dots between atoms represents the shared electron pair. The unshared pair, also called lone pairs, is a pair of electrons that is not involved in bonding and that belongs exclusively to one atom. A single covalent, or a single bond, is a covalent bond produced by the sharing of one pair of electrons between two atoms.

Lewis Structures are formula in which atomic symbols represent nuclei and inner-shell electrons, dot-pairs or dashes between two atomic symbols represent electron pairs in covalent bonds, and dots adjacent to only one atomic symbol represent unshared electrons. A structural formula indicates the kind, number, arrangement, and bonds but not the unshared pars of the atoms in a molecule.

For example:
N-N H-Cl

Sample Problem 2: Determine number of bonds needed in Lewis Structure also.

Draw the Lewis structure for iodomethane, CH3I.
Step 1. Determine the type and number of atoms in the molecule.
The formula shows one carbon atom, one iodine atom, and three hydrogen atoms.
Step 2. Write the electron-dot notation for each type of atom in the molecule.
Step 3. Determine the total number of valence electrons in the atoms to be combined.
Step 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.
Step 5. Add unshared pairs of electrons so that each hydrogen atom shares a pair of electrons and each other nonmetal is surrounded by eight electrons.
Step 6. Count the electrons in the structure to be sure that the number of valence electrons used, equal the number available.
Draw the Lewis structure for
ammonia, NH3
hydrogen sulfide, H2S

Multiple Covalent Bonds

Atoms of some elements, especially carbon, nitrogen, and oxygen, can share more than one electron pair. A double covalent bond, or simply double bond, is a covalent bond produced by the sharing of two pairs of electrons between two atoms. A triple bond is a covalent bond produced by the sharing of three pairs of electrons between two atoms.

Elemental nitrogen / C2H2

Double bonds in general have higher bond energies and are shorter than single bonds. Triple bonds are even stronger and shorter.

Sample Problem 3:

Draw the Lewis structure for methanal, CH2O, which is also known as formaldehyde.
Step 1. Determine the number of atoms of each element present in the molecule.
Step 2. Write the electron-dot notation for each type of atom.
Step 3. Determine the total number of valence electrons in the atoms to be combined.
Step 4. Arrange the atoms to form a skeleton structure for the molecule, and connect the atoms by electron-pair bonds.
Step 5. Add unshared pairs of electrons so that each hydrogen atom shares a pair of electrons and each other nonmetal is surrounded by eight electrons.
Step 6a. Count the electrons in the Lewis structure to be sure that the number of valence electrons used equals the number available.
Step 6b. If too many electrons have been used, subtract one or more lone pairs until the total number of valence electrons is correct. Then move one or more lone electron pairs to existing bonds between non-hydrogen atoms until the outer shells of all atom are completely filled.

Practice Problem 3.

Draw the Lewis structure for carbon dioxide.
Draw the Lewis structure for hydrogen cyanide, which contains one hydrogen atom, one carbon atom, and one nitrogen atom.

Resonance Structures – Resonance refers to bonding in molecules or ions that cannot be correctly represented by a single Lewis structure. To indicate resonance, a double-headed arrow is placed between a molecule’s resonance structures.

Resonance

With certain molecules or ions, given the atomic geometry, it is possible to satisfy the octet rule with more than one bonding arrangement. The proceeding structures are called resonance structures:

Molecules or ions that have two or more resonance structures are said to exhibit resonance. The actual bonding in such molecules or ions is thought to be an average of the bonding present in the resonance structures and (for the above example) might be represented as

The stability of molecules or ions exhibiting resonance is found to be higher than that anticipated for any single resonance structure.

Ionic Bonding: An ionic compound is composed of positive and negative ions that are combined so that the numbers of positive and negative charges are equal. A formula unit is the simplest collection of atoms from which an ionic compound’s formula can be established. For example, one formula unit of sodium chloride, NaCl, is one sodium cation plus one chloride anion.

The ratio of ions in a formula unit depends on the charge of the ions combined. For example, to achieve electrical neutrality in the ionic compound calcium fluoride, two fluoride anions, F-, each with a charge of 1-, must balance the 2+ charge of each calcium cation, Ca2+. Therefore formula of calcium fluoride is CaF2.

Draw the Lewis dot notation for the calcium atom reacting with the fluorine atoms to form calcium fluoride.

Crystal Lattice – 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.

Polyatomic Ions – A charged group of covalently bonded atoms. Draw the Lewis Structures for the following polyatomic ions.

Ammonium ion nitrate ion
Hydronium ion phosphate ion
Sulfate ion

Molecular Geometry

VSEPR – “Valence Shell Electron Pair Repulsion” referrs to the repulsion between pairs of valence electrons of the atoms in a molecule. The VSEPR Theory states that repulsion between the sets of valence-level electrons surrounding an atom causes these sets to be oriented as far apart as possible.

Molecular geometry / Formula Type / Ideal bond angles (with lone pair effect) / Examples
Linear / AX, AX2 / 180 o
Bent / AX2E, AX2E2 / 105 o / H2O
Trigonal planar / AX3 / 120 o / BCl3
Trigonal pyramidal / AX3E / 109.5 o (< 109.5 o) / NH3
T-shaped / AX3E2 / 90o (< 90 o) / ClF3
Tetrahedral / AX4 / 109.5 o / CH4
Seesaw shaped / AX4E / 90 o (< 90 o),
120 o (< 120 o) / SF4
Square planar / AX4E2 / 90 o / XeF4
Square pyramidal / AX5E / 90o (<90 o) / BrF5
Octahedral / AX6 / 90 o / SF6

Sample Problem 4.

Use VSEPR theory to predict the molecular geometry of aluminum trichloride, AlCl3.
Step 1. Determine the total number of valence electrons in the Lewis structure.
(hint, Al does not fulfill the octet rule)
What is the formula type of molecule?
Geometry?

Practice 4a.

Use VSEPR theory to predict the molecular geometry of the following molecules:
HI
CBr4
AlBr3
CH2Cl2

VSEPR and Unshared Electron Pairs: VSEPR theory postulates that the lone pair occupies space around the nitrogen atom just as the bonding pairs do. Lone pairs occupy more space than bonding pairs of electrons occupy.

Sample Problem 5

Use VSEPR Theory to predict the shape of a molecule of:
carbon dioxide
Chlorate ion
Water

Practice Problem 5a.

Use VSEPR theory to predict the molecular geometries of:
sulfur difluoride.
Phosphorus trichloride
Phosphorus pentachloride
Sulfur hexafluoride

Hybridization is the mixing of two or more atomic orbitals of similar energies on the same atom to produce new orbitals of equal energies.

Methane, CH4, has 4 valence electrons, two in the 2s orbital and two in the 2p orbitals.

C ______

1s 2s 2px 2py 2pz

To achieve 4 equivalent bonds, carbon’s 2s and 2p orbitals hybridize to form 4 new, identical orbitals called sp3 orbitals.