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Chapter 6 Chemical Bonding
I: Introduction to Chemical Bonding
chemical bond - definition
Elements go from a high potential energy to a lower potential energy in chemical bond formation -- more stable in bond formation that by themselves.
A: Types of Chemical Bonding
bonding involves valence electrons
cations
anions
ionic bonding - definition
Ionic bonding involves one atom giving up one or more electrons and another atom(s) accepting these electrons - this forms the ions and the opposite attraction of the ions is what forms the bond.
covalent bonding - definition
figure 6-1 page 162
metallic bonding
two parts of a pure metal a) valence electrons which are mobile; b) kernel which is the nucleus and the nonvalence electrons
1. Ionic or Covalent?
rarely ionic or covalent -- usually a blend of the two
electronegativity differences
figure 6-2 page 162
0 - 0.3 -- pure covalent -- nonpolar covalent bond -- percent ionic character of 5% or less
0.3 to 1.7 -- polar covalent bond
1.7 or less -- ionic character of 50% or less -- covalent
1.7 to 3.3 -- ionic bond -- percent ionic character of 50% or greater
Polar - definition
figure 6-3 page 163
electron density - closer to more electronegative element
d- and d+ e.g. HCl
Homework 6.1
II: Chemical Bonding and Molecular Compounds
molecule - definition
a unit that can exist on its own
two or more of the same type of atom or two or more different atoms
molecular compound - definition
chemical formula - definition
molecular formula - definition
diatomic molecule - definition
A: Formation of a Covalent Bond
hydrogen-hydrogen bond
electron configuration
orbital notation of valence electrons
how can each element best get an octet?
nuclei of one atom attracted to electron cloud of other atom and visa versa
eventually the nuclei repel each other and so an optimal distance between the nuclei is reached
figure 6-5 page 165
figure 6-6 page 165
B: Characteristics of the Covalent Bond
the electron cloud and, therefore, the orbitals overlap
figure 6-7 page 167
bond length - definition
bond energy - definition
bond lengths and bond energies vary with the type of atoms that have combined
in general, the shorter the bond length, the higher the bond energy
table 6-1 page 168
sharing electrons to get an octet or a noble gas configuration e.g. hydrogen molecule
figure 6-8 page 168
Homework 6.2
C: The Octet Rule
noble gases exist independently in nature
they have a minimum of potential energy
stable electron configuration -- filled s and p sublevel of the outermost energy level -- octet
Octet rule - statement
fluorine molecule F2
figure 6-9a page 169
hydrogen chloride molecule HCl
figure 6-9b page 169
1. Exceptions to the Octet Rule
octet rule applies to most main group elements that form covalent bonds
exceptions: B - three pairs of electrons instead of four e.g. BF3
Some elements can be surrounded by more than four pairs of electrons when bonding with highly electronegative elements fluorine, oxygen, and chlorine -- these are said to have expanded valence involving d as well as s and p orbitals.
2. Electron Dot Notation
electron dot notation - definition
shows paired and unpaired electrons
figure 6-10 page 170
electron dot notation for hydrogen
electron dot notation for nitrogen
3. Lewis Structures
electron dot formula vs electron dot notation
formula refers to molecules/compounds
electron dot formula for fluorine F2
unshared pair or lone pair - definition
Lewis structure replaces the bonding pair of electrons with a dash
Structural formula - definition
e.g. H-H or F-F or H-Cl
Single covalent bond (single bond) - definition
Homework 6.3
4. Multiple Covalent Bonds
double covalent bond - definition
two pairs of side by side dots or two parallel dashes
both atoms share the four electrons (two pairs) involved
C2H4 is the compound ethene
triple covalent bond - definition
three pairs of side by side dots or three parallel dashes
both atoms share the six electrons involved
N2
figure 6-11 page 173
C2H2 is the compound ethyne
multiple bonds - definition
double bonds have higher bond energy and are shorter than single bonds
triple bonds have higher bond energy and are shorter than double bonds
carbon, oxygen, nitrogen are elements that have the possibility of forming multiple covalent bonds with the same element e.g. oxygen with oxygen, nitrogen with nitrogen, carbon with carbon
Homework 6.4
5. Resonance Structures
Some molecules and ions cannot be represented by a single Lewis structure e.g. O3
O = O - O <----> O - O = O diagram page
175
even though the resonance structures indicate the double bond on the left molecule between the first and second oxygen and then a single bond on the right molecule between the first and second oxygen, the actual bond is somewhere between a single and double bond
the two bonds are identical but we don't have a good way to represent it using a single Lewis structure - thus resonance
the two structures are resonance structures or resonance hybrids
Resonance - definition
double headed arrow - only place this is allowed - between the two resonance structures
Homework 6.5
6. Covalent Network Bonding
molecules with covalent bonding usually consist of molecules - molecules held together by forces holding the molecules together
some molecules with covalent bonding consist of continuous three dimensional network of bonded atoms.
more on this in a later chapter.
Homework 6.6
III: Ionic Bonding and Ionic Compounds
ionic compound - definition
most ionic compounds are crystalline solids - figure 6-12 page 176
Empirical formula - definition
the ratio of ions in an empirical formula depends on the charges of the ions combined e.g. calcium fluoride v sodium fluoride
A: Formation of Ionic Compounds
electron dot equation for reaction of sodium and chlorine to yield sodium chloride
electron dot equation for the reaction between calcium and fluorine to yield calcium fluoride
1. Characteristics of Ionic Bonding
ions minimize their potential energy by combining in an orderly arrangement known as a crystal lattice
figure 6-13 page 177
attractive forces present in ionic crystal include: a) those between oppositely charged ions; b) those between the nuclei and electrons of adjacent ions
repulsive forces in an ionic crystal include: a) those between like-charged ions; b) those between electrons of adjacent ions
distance between ions and their arrangement in a crystal represent a balance among all these forces
figure 6-14 page 177
each sodium cation surrounded by six chloride anions; each chloride anion is surrounded by six sodium cations
the arrangements of ions and the strengths of attraction between the ions vary with
a) the sizes and charges of the ions and
b) the numbers of ions of different charges
figure 6-15 page 178
lattice energy - definition
used to compare bond strengths in ionic compounds
table 6-3 page 179
negative values indicate energy is released (exothermic)
B: A Comparison of Ionic and Molecular Compounds
The forces between molecules are much weaker than the forces of ionic bonding -- thus different properties.
Melting point, boiling point and hardness depend on how strongly its basic units are attracted to each other.
Many molecular compounds melt at low temperatures while many ionic compounds have higher melting and boiling points.
Ionic cpds do not vaporize as readily at room temperature as molecular cpds do.
Ionic cpds are hard but brittle.
figure 6-17 page 179
In the molten state (melted) or when dissolved in water, ionic cpds conduct electricity but not in the solid state.
Those ionic cpds that are not soluble (dissolve) in water are cpds in which the water molecules cannot overcome the attraction between the ions of the cpd.
C: Polyatomic Ions
Polyatomic ions - definition
As with any ion, these result from the shortage or excess of electrons.
Difference is that they contain more than one type of element.
polyatomic ions - page 180
Homework 6.7
IV: Metallic Bonding
The bonding in metals reflects their properties.
Mobile valence electrons.
A: Metallic Bond Model
In metals, usually, the s sublevel is filled and the three orbitals of the p sublevel are empty.
Also have some vacant d orbitals.
The vacant orbitals ant he atoms' outer energy levels overlap. This overlapping of orbitals allows the outer electrons of the atoms to roam freely throughout the entire metal.
Electrons are said to be delocalized.
Mobile electrons form a sea of electrons.
Metallic bonding - definition
1. Metallic Properties
delocalization of electrons explain the high electrical and thermal conductivity.
because they contain many orbitals separated by extremely small energy differences, metals can absorb a wide range of light frequencies - luster.
Malleability - definition
Ductility - definition
Metallic bonding is the same in all directions throughout the solid. One plane of atoms in a metal can slide past another without encountering any resistance or breaking any bonds, unlike ionic crystals.
2 Metallic Bond Strength
Varies with the nuclear charge of the metal atoms and the number of electrons in the metal's electron sea.
Both are reflected in the metal's heat of vaporization.
Heat of vaporization - definition
The amount of heat is a measure of the strength of the bonds that hold the metal together.
table 6-4 page 182
V: Molecular Geometry
properties depend on bonding and molecular geometry
polarity of each bond and the geometry of the molecule determines the molecular polarity
molecular geometry - definition
molecular polarity - definition
molecular polarity influences the forces that act between molecules in liquids and solids
the chemical formula does not tell us directly about molecular polarity
two theories are prevalent a) molecular bond angles; b) describe the orbitals that contain the valence electrons
A: VSEPR Theory
figure 6-20 page 183
when there are only two atoms (diatomic) the geometry must be linear
more complicated molecules - consider all electron pairs surrounding the bonded atoms -- VSEPR theory
VSEPR - definition
consider molecules with no unpaired (unshared) electron pairs then those with unpaired (unshared) electron pairs.
first example: BeF2
methodology: a) inert gas electron configuration; b) orbital notation for valence electrons; c) dot structure for each element; d) dot formula for molecule; e) consider effect of electron pairs on geometry
BeF2 dot formula
VSEPR states that electron pairs orient themselves to be as far away from each other as possible
figure 6-21 page 184
all berrylium electrons are involved in bonds; each fluorine has one pair of electrons in a bond and the other three are nonbonding pairs of electrons and so they get as far away from each other as possible
molecule is characterized as an AB2 which is linear and has the central atom with no nonbonding electron pairs and two similar atoms bonded to the central atom, bond angles of 180 degrees, geometry is linear
second example: BF3
methodology: a) inert gas electron configuration; b) orbital notation for valence electrons; c) dot structure for each element; d) dot formula for molecule; e) consider effect of electron pairs on geometry
this is an AB3 molecule in which the three atoms occupy the corners of an equilateral triangle, bond angles of 120 degrees, geometry is trigonal planar (one central atom attached to three other atoms)
third example: CH4
methodology: a) inert gas electron configuration; b) orbital notation for valence electrons; c) dot structure for each element; d) dot formula for molecule; e) consider effect of electron pairs on geometry
carbon uses its four valence electrons to bond with four other atoms; the four atoms occupy the corners of a tetrahedron, bond angles are 109.5 degrees; a tetrahedron consists of four identical triangles put together
other molecules on table 6-5 page 186
keep in mind that if the "B" atoms are not all the same, this will distort the molecule's geometry
Homework 6.8: Review Sample Problem 6-5, page 184, Practice problems, page 185, but do each part using the methodology from above; Section Review, page 193, # 1
1. VSEPR and Unshared Electron Pairs
central atom has unshared electron pairs e.g. water and ammonia
How VSEPR theory handles this:
electron dot formula for ammonia, NH3
lone pair occupies space just as the bonding pairs do
this would not be an AB4 molecule since the lone pair affects things differently than the bonding pair
ammonia is described as an AB3E, where A is the central atom, B represents atoms bonded to the central atom, E represents lone pair(s)
bond angle is 107 degrees, less than the 109.5 degrees of a tetrahedron -- lone pairs repel electrons more than bonding pairs do; geometry is trigonal pyramidal
water molecule: the oxygen has two lone pairs; it is an AB2E2; oxygen is at the center of a modified tetrahedron, with two hydrogen's occupying two of the three corners of the base, one lone pair at the third corner of the base and one lone pair at the top. Geometry is described as bent.
Figure 6-22 page 185
Homework 6.9
B: Hybridization
hybridization - definition
the compound methane, CH4
for carbon: electron configuration and orbital notation; change that results from hybridization that involves the 2s and 2p orbitals to form a new orbital called the sp3 hybridized orbital
figure 6-23 page 188
new electron configuration to show hybridization
hybrid orbital - definition
explain the geometry of molecules formed by Group 15 and 16 elements.
We will consider hybridization only as it applies to carbon
Homework 6.10
C: Intermolecular Forces
boiling point is a good measure of the force of attraction between particles of a substance i.e. attractive forces between molecules
intermolecular forces - definition
usually they are weaker than chemical bonds
comparing boiling points of metals and ionic compounds with boiling points of molecular substances table 6-7 page 190
1. Molecular Polarity and Dipole-Dipole Forces
strongest intermolecular forces exist between polar molecules
polar molecules act as tiny dipoles
dipole - definition
The direction of the dipole is from the positive to the negative end of the molecule.
Indicated by an arrow pointing toward the negative end of the molecule. The tail of the arrow is crossed and is at the positive end of the molecule.
dipole-dipole forces - definition
these are short range forces
figure 6-25 page 191
For molecules containing more than two atoms, molecular polarity depends on both the polarity and the orientation of each bond.
e.g. water
figure 6-26 page 191
e.g. ammonia
In some molecules, individual bond dipoles cancel one another, causing the resulting molecular polarity to be zero.
E.g. carbon tetrachloride and carbon dioxide figure 6-26 page 191
A polar molecule can induce a dipole in a nonpolar molecule by temporarily attracting its electrons. The short range intermolecular force is somewhat weaker than the dipole-dipole force.
The force of an induced dipole accounts for the solubility of nonpolar oxygen in water.
Figure 6-27 page 192
2. Hydrogen Bonding
This is a particularly strong type of dipole-dipole force.
Occurs between the hydrogen of one molecule and the F, O, or N of an adjacent molecule.
examples are HF, water and ammonia
hydrogen bonding - definition
represented by dotted lines connecting the hydrogen of one molecule to the highly electronegative atom (fluorine, oxygen or nitrogen) of an adjacent molecule.
effect can be seen by comparing boiling points of phosphine and ammonia, and hydrogen sulfide and water on table 6-7 page 190
3. London Dispersion Forces
noble gases and nonpolar molecules still experience a weak intermolecular attraction
because of the random motion of the electrons, the distribution of the electrons may become uneven
the momentary unequal distribution of the electrons creates a temporary dipole in the molecule; this temporary dipole can induce a dipole in an adjacent molecule and the two molecules are attracted for an instant then the effect disappears
London dispersion forces - definition
These forces operate between all atoms and molecules but they are the only intermolecular forces acting among noble gas atoms, nonpolar molecules, and slightly polar molecules. Notice the low bp's of the noble gases, etc. on table 6-7 page 190.
Because these forces depend on the motion of the electrons, their strength increases with the number of electrons in the interacting atoms or molecules. i.e. the forces increase with increasing atomic or molar mass. Note the BP's of helium and argon, hydrogen and oxygen, chlorine and bromine.
Homework 6.11
end of notes
London dispersion forces are the intermolecular attractions resulting from the constant motion of electrons and the creation of instantaneous dipoles. back
Hydrogen bonding is the intermolecular force in which 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. Back
Dipole-dipole forces are the forces of attraction between polar molecules. Back
A dipole is created by equal but opposite charges that are separated by a short distance. Back
Intermolecular forces are the forces of attraction between molecules. Back
Hybrid orbitals are orbitals of equal energy produced by the combination of two or more orbitals on the same atom. Back
Hybridization is a mixing of two or more atomic orbitals of similar energies on the same atom to produce new orbitals of equal energies.
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.
VSEPR stands for valence shell electron pair repulsion. Back
Molecular polarity is the uneven distribution of molecular charge. Back
Molecular geometry is the three dimensional arrangement of a molecule's atoms in space. Back
Heat of vaporization is the heat necessary to convert a metal from the solid state to individual metal atoms in the gaseous state. Back
Ductility is the ability of a substance to be drawn, pulled, or extruded through a small opening to produce a wire. Back
Malleability is the ability of a substance to be hammered or beaten into thin sheets. Back
Metallic bonding is the bonding that results from the attraction between metal atoms and the surrounding sea of electrons.back
A polyatomic ion is a charged group of covalently bonded atoms. Back
Lattice energy is the energy released when one mole of an ionic crystalline compound is formed from gaseous ions. Back
An empirical formula (formula unit) indicates what elements are present and the simple whole number ratio of those elements.back
Ionic compound is composed of positive and negative ions that are combined so that the numbers of positive and negative charges are equal. Back
Resonance refers to bonding in molecules or ions that cannot be correctly represented by a single Lewis structure. Back
A triple covalent bond is a covalent bond produced by the sharing of three pairs of electrons between two atoms. Back
Multiple bonds are double or triple covalent bonds. Back
A double covalent bond is a covalent bond produced by the sharing of two pairs of electrons between two atoms. Back
A single covalent bond is a covalent bond produced by the sharing of one pair of electrons between two atoms. Back
Structural formula indicates the kind, number, arrangement, and bonds but not the unshared pairs of the atoms in a molecule. Back
An unshared pair of electrons is a pair of electrons that is not involved in bonding and that belongs exclusively to one atom. Back
Electron dot notation is an electron configuration notation in which only the valence electrons of an atom of a particular element are shown, indicated by dots placed around the element's symbol. Back
Chemical compounds tend to form so that each atom, by gaining, losing, or sharing electrons, has an octet of electrons in its highest occupied energy level. Back
Bond energy is the energy required to break a chemical bond and form neutral isolated atoms. Back
The distance between two bonded atoms at their minimum potential energy is the bond length. Back
A diatomic molecule is a molecule containing only two atoms.
The diatomic molecules are fluorine, chlorine,bromine, iodine, hydrogen, oxygen, and nitrogen.
A molecular formula shows the types and numbers of atoms combined in a single molecule of a molecular compound. Back
A chemical formula indicates the relative numbers of atoms of each kind in a chemical compound by using atomic symbols and numerical subscripts. Back
A molecular compound is a chemical compound whose simplest units are molecules. Back
A molecule is a neutral group of atoms that are held together by covalent bonds. Back
Polar means an uneven distribution of charge. Back
A polar covalent bond is a covalent bond in which the bonded atoms have an unequal attraction for the shared electrons. Back
A nonpolar covalent bond is a covalent bond in which the bonding electrons are shared equally by the bonded atoms, resulting in a balanced distribution of electrical charge. Back
Covalent bonding results from the sharing of electron pairs between two atoms. Back
Ionic bonding is a chemical bond that results from the electrical attraction between large numbers of cations and anions. Back
Chemical Bond is a mutual electrical attraction between the nuclei and valence electrons of different atoms that binds the atoms together. Back