CH 101 - Chapter 5 Notes

Ionic substances

extended networks of large numbers of ions and not individual molecules

What makes the ionic bond non-directional?

The charge on the ions is spherical

covalent bonds

Bonds that are directional that form between nonmetals

Bond length

The separation between the nuclei at the minimum energy

What type of nuclei can get closer to atoms than the other?

Smaller nuclei can get closer to other atoms than can larger nuclei

Bond energy

The amount by which the energy of the two atoms is reduced by forming the bond

χ

electronegativity, which measures an atom's ability to attract the bonding electrons

What kind of orbitals do atoms with high electronegativities have?

They have unfilled, low-energy orbitals and strong affinities for bonding electrons

What happens if the electronegativities of the bound atoms differ?

The bonding electrons are not shared equally as the more electronegative atom becomes electron rich and the less electronegative atom becomes electron poor.

Partial charge δ (delta)

The charge that is produced in a covalent bond and is only a part of what it would be if the bond were ionic

Bond dipole

the polarity of the bond

How does the strength of a bond dipole (the polarity of the bond) increase?

It increases as the electronegativity difference (Δχ) between the bound atoms increases.

When is the bond between two atoms is purely covalent?

If their electronegativities are identical; but it becomes increasingly polar as the difference in their electronegativities increases

percent ionic character

The more polar a bond is, the higher is its percent ionic character.

Why is no bond 100% ionic?

Every atom exerts some pull on the bonding electrons; i.e., no atom has zero electronegativity

Nonmetal-nonmetal bonds are covalent, while

metal-nonmetal bonds are ionic.

Binary compounds

compounds that consist of only two elements; they are named by giving the name of the less electronegative atom followed by the name of the more electronegative atom with its ending changed to -ide

Which element is written first in a formula?

the less electronegative element

What does the placement of H mean in a formula?

If H is at the beginning of a formula, the compound is acidic. If it's at the end, it's not acidic.

Lewis structures

representation of the molecule that shows the distribution of the valence electrons of its constituent atoms

Why are the orbitals in a Lewis structure are already drawn with the positions required for bonding?

Lewis symbols represent an atom that is about to bond, not an isolated atom

In ionic compounds, nonmetals gain electrons from a metal to attain filled valence shells. What about covalent compounds?

In covalent compounds, nonmetals share electrons with other nonmetals to attain filled valence shells.

octet of valence electrons

A closed valence shell for a nonmetal consists of eight electrons (filled s and p sublevels)

the octet rule

The tendency of nonmetals to obtain eight valence electrons, with Hydrogen as the only exception (it only wants 2 electrons)

How are Lewis structures of molecules obtained?

By giving each atom (except hydrogen) an octet of valence electrons

nonbonding electrons or lone pairs

electrons not involved in bonding

ER

electrons required to give each atom an octet

VE

valence electrons

SP

shared pairs/one-half the number of shared
electrons

SP

½[ER - VE]

What does each line in a bond represent?

two electrons

bond order

number of shared pairs the bond contains

single bond

bond order of 1 because it contains a single pair of shared electrons

double bond

a covalent bond in which two pairs of electrons are shared between two atoms

triple bond

three pairs of shared electrons

The bond strength of a specific bond increases and its bond length ________

decreases as its bond order increases

DETERMINING LEWIS STRUCTURES

1. Determine the number of electrons required (ER) to give each isolated non-hydrogen atom eight electrons and each hydrogen two electrons. ER is eight times the number of non-hydrogen
atoms (NA) plus two times the number of hydrogen atoms (NH).
ER = 8NA + 2NH
2. Determine the number of valence electrons in the molecule (VE), which is the number of electrons that must be shown in the final structure.
3. The difference, ER - VE, divided by 2 is the number of shared pairs (SP).
SP = ½(ER - VE)
4. Draw the molecule with the bonded atoms connected by single bonds. Then satisfy the number of shared pairs required (SP) by adding double or triple bonds as necessary. Finally, add lone pairs of electrons to complete each atom’s octet, but remember that hydrogen atoms have only two electrons.

The following should also be considered when constructing a Lewis structure:

• The first atom in the formula is the least electronegative and is generally the central atom.
• Hydrogen atoms have only one bond.
• Multiple bonds should not be drawn to terminal halogen atoms.* More than one single bond can be drawn to a halogen when it is the central atom, as in the perchlorate ion (ClO41-).

Resonance structures

Lewis structures that differ only in the placement of electrons

How is charge assigned to atoms?

By assigning each of the valence electrons in a molecule to one of the atoms

formal charge

The resulting charge on the atom if the bonds are assumed to be covalent

oxidation state

The resulting charge on the atom if the bonds are assumed to be ionic

assigned charge

the number of valence electrons in the free (neutral) atom minus the number of valence electrons assigned to the atom in the molecule

NB

nonbonding electrons

BE

bonding electrons

The formal charge of atom A (FCA) is determined to be

FCA = VE - (NB + ½ BE)

The sum of the formal charges must equal what?

the charge on the species (zero for a molecule)

Coulomb’s law

energy is required to separate positive and negative charge

When can formal charge be used to help determine the preferred structure of a molecule?

When more than one Lewis structure can be drawn for a molecule

What is the Lewis structure that is preferred?

The one in which the formal charges of all atoms are closest to zero. If non-zero formal charge must be assigned, negative formal charge should reside on the more electronegative atoms.

Why should placing more than one bonding pair to a halogen be avoided?

Because it places positive formal charges on these electronegative atoms

OXA

VE - (NB + ∑j ajBE)

CHAPTER SUMMARY AND OBJECTIVES

Covalent bonds result when bonding electrons are shared. The energy of the bound atoms
is less than that of the separated atoms because the bonding electrons experience the positive charge of both nuclei. The energy required to break the bond is called the bond energy. When the electronegativities of the atoms are different, the bonding electrons are not shared equally to produce a bond dipole. The polarity of a bond increases as the electronegativity difference between the two atoms increases. The less electronegative atom is written first in the formula and maintains its name, while the more electronegative element is written last, and its ending is changed to -ide. The number of each atom present in the formula is given by a prefix.
Lewis structures show the valence electrons in a compound as dots if they are in lone pairs or lines if they are bonding pairs. The number of shared pairs is determined as SP = 1/2 (VE - ER). The number of shared pairs in a bond is called the bond order. Resonance structures are Lewis structures that differ only in the placement of the electron pairs. Electron counting can be done with formal charge or oxidation state. Bonding electrons are assigned equally to the bound atoms in the formal charge, but they are
assigned to the more electronegative element in the oxidation state. Thus, formal charge is a better description of charge distribution in purely covalent compounds, while oxidation state gives a better picture in very polar or ionic compounds.