Chemistry

Chapter 8

?In 1864, English chemists john newlands noticed that when the known elements were arranged in order of atomic mass, every eighth element had

similar properties. Newlands referred to this peculiar relationship as the

law of octaves. Howe ver, this “law” turned out to be inadequate for

elements beyond calcium, and newland’s work was not accepted by the

scientific community.

?Representative elements are the elements in groups 1A through 7A, all of which have incompletely filled s or p subshells of the highest principal

quantum number. With the exception of helium, the noble gases (the

group 8A elements) all have a completely filled p subshell. The transition

metals are the elements in groups 1B and 3B through 8B, which have

incompletely filled d subshells or readily produce cations with

incompletely filled d subshells (these metals are sometimes referred to as

the d-blok transition elements). The group 2B elements are Zn, Cd, and

Hg, which are neither representative elements nor transition metals. The

lanthanides and actinides are sometimes called f-block transition elements

because they have incompletely filled f subshells

?The outer electrons of an atom, which are those involved in chemical bonding are often called the valence electrons. Having the same number of

valence electrons accounts for similarities in chemical behavior among the

elements within each of these groups.

?Ions, or atoms and ions, that have the same number of electrons and hence the same ground-state electron configuration are said to be isoelectronic.

?Atomic radius of a metal is one-half the distance between the two-nuclei in two adjacent atoms. For elements that exist as diatomic molecules, the

atomic radius is one-half the distance between the nuclei of the two atoms

in a particular molecule.

?When looking at a periodic table:

o The elements are increasing as in atomic radius as you go from

right to left, and from up to down. ****

?Ionic radius is the radius of a cation or an anion. Ionic radius affects the physical and chemical properties of an ionic compound.

?When a neutral atom is converted to an ion, we expect a change in size. If the atoms forms an anion, its size increases, because the nuclear charge

remains the same but the repulsion resulting from the additional electron

enlarges the domain of the electron cloud. On the other hand, a cation is

smaller than the neutral atom, because removing one or more electrons

reduces electron-electron repulsion but the nuclear charge remains the

same, so the electron cloud shrinks.

?Focusing on isoelectronic cations, we see that the radii of tripostive ions (that is, ions that bear three positive charges) are smaller than those of

dipositive ions (that is, ions that bear two positive charges) which in turn

are smaller than unipositve ions (that is, ions that bear one positive charge).

?Ionization energy – is the minimum energy required to remove an electron from a gaseous atom in its ground state. The magnitude of ionization

energy is a measure of the effort required to force an atom to give up an

electron, or of how “tightly” the electron is held in the atom., the higher

the ionization energy the more difficult it is to remove the electron.

?For a many-electron atom, the amount of energy required to remove the first electron, from the atom in its ground state:

o Energy + X(g) -> X+(g) + e-

o Is called the first ionization energy (I1). In this equation X

represent a gaseous atom of any element and e- is an electron.

Unlike an atom in the condensed liquid and solid phases, an atom

is the gaseous phase is virtually uninfluenced by its neighbors.

o Energy + X+(g) -> X2+(g) + e- Second ionization

o Energy + X2+(g) -> X3+(g) + e- Third Ionization

?When a electron is removed from a neutral atom, the repulsion among the remains electrons decreases. Because the nuclear charge remains constant,

more energy is needed to remove another electron from the positively

charged ions. Thus for the same element ionization energies always

increase in this order:

o I1

?Another property that greatly influences the chemical behavior of atoms is their ability to accept one or more electrons. This ability is called electron

affinity, which is the negative of the energy change that occurs when an

electron is accepted by an atom of an element in the gaseous state

o X(g) + e- -> X-(g) deltaH = -XXXkJ

?If delta h has a positive value (ie. 390 kj/mol) means that

the process is exothermic

?If delta h has a negative value, that means that the process

is endothermic

?Another trend in chemical behavior of the representative elements is the diagonal relationship. Diagonal relationship refers to similarities that exist

between pairs of elements in different groups and period of the periodic

table. Specifically the first three members of the second period (Li, Be and

B) exhibit many similarities to the elements located diagonally below

them in the periodic table.

If you would like to further understand this chapter, I suggested reading the summary. Or if you would like to learn more about the individual group elements, then I suggest reading the last few pages of this chapter.

Chapter 9

?Lewis dot symbol – consists of the symbol of an element and one dot for each valence electron in an atom of the element.

?Covalent bond – a bond in which two electrons are shared by two atoms.

Covalent compounds are compounds that contain only covalent bonds.

?Lone pairs – pairs of valence electrons that are not involved in covalent bond formation (ie. F2)

?Lewis structures is a representation of covalent bonding in which shared electron pairs are shown either as lines or as pairs of dots between two atoms, and lone pairs are shown as pairs of dots on individual atoms. Only valence electrons are shown in a Lewis structure.

?Octet rule – an atom other than hydrogen tends to from bonds until it is surrounded by eight valence electrons. In other words, a covalent b ond forms when there are not enough electrons for each individual atom to

have a complete octet. By sharing electrons in a covalent bond, the

individual atoms can complete their octets. The requirement for hydrogen is that it attains the electron configuration of helium, or a total of two

electrons.

o The octet rule works mainly for elements in the second period of the periodic table.

?Atoms can form different types of covalent bonds. In a single bond – two atoms are held together by one electron pair. Many compounds are held together by multiple bonds, that is, bonds formed when two atoms shre two or more pairs of electrons. If two atoms share two pairs of electrons, the covalent bond is called a double bond.

? A triple bond arises when two atoms share three pairs of electrons, (N2) ?Bond length – is defined as the distance between the nuclei of two covalently bonded atoms in a molecule.

?The bond HF is called a polar covalent bond, or simply a polar bond, because the electrons spend more time in the vicinity of one atom than the other. The HF bond and other polar bonds can be though of as being

intermediate between a (nonpolar) covalent bond, in which the sharing of electrons is exactly equal, and an ionic bond, in which the transfer of the electron(s) is nearly complete.

? A property that helps us distinguish a nonpolar covalent bond from a polar covalent bond is electronegativity, the ability of an atom to attract toward itself the electrons in a chemical bond. Elements with high

electronegativity have a greater tendency to attract electrons than do

elements with low electronegativity.

o Electronegativity is related to electron affinity and ionization energy.

o Electronegativity is a relative concept, mea ign tha t an element’s ectronegativity can be measured only in relation the

electronegativity of other elements.

o Linus Pauling devised a method for calculating relative

electronegativities of most elements.

?There is no sharp distinction between a polar bond and an ionic bond, but the following rule is helpful in distinguishing between them. An ionic

bond forms when the electronegativity difference between the two

bonding atoms is 2.0 more. This rule applies to most but not all ionic

compounds. Sometimes chemists use the quantity percent ionic character

to describe the nature of a bond. A purely ionic bond would have 100

percent ionic character, although no such bond is known, whereas a

nonpolar or purely covalent bond has 0 percent ionic character.

?Electronegativity and electron affinity are related but different concepts.

Both indicate the tendency of an atom to attract electrons. However,

electron affinity refers to an isolated atom’s attraction for an additional

electron, whereas electronegativity signifies the ability of an atom in a

chemical bond (with another atom) to attract the shared electron.

Furthermore, the electron affinity is an experimentally measurable

quantity, whereas electronegativity is an estimated number that cannot be measured.

?An atom’s formal charge is the electrical charge difference between the valence electrons in an isolated atom and the number of electrons assigned to an atom in a lewis structure.

?To assign the number of electrons on an atom in a lewis structure, we proceed as:

o All the ato m’s nonbonding electrons are assigned to the atom

o We break the bond(s) between the atom and other atom(s) and assign half of the bonding electrons to the atom

?When you write formal charges, these rules are helpful:

o For molecules, the sum of the formal charges must add up to zero because they are electrically neutral species.

o For cations, the sum of the formal charges must equal the positive charge

o For anions, the sum of the formal charges must equal the negative charge

?Keep in mind, that formal charges do not represent actual charge separation within the molecule.

?Resonance structure – one of two or more lewis structures for a single molecule that cannot be represented accurately by only one lewis structure.

The double-headed arrow indicates that the structures shown are

resonance structures.

?The term resonance itself means the use of two or more lewis structures to represent a particular molecule.

?Exceptions to the octet rule:

o The incomplete octet:

?In some compounds the number of electrons surround the

central atom in a stable molecule is fewer than eight.

?Elements in group 3A, particularly boron and aluminum,

also tend to form compounds in which they are surrounded

by fewer than eight electrons.

? A resonance structure with a double bond between

B and F can be drawn that satisfies the octet rule for

B.

?The B-N bond is different from the covalent bonds

discussed so far in the sense that both electrons are

contributed by the N atom. A covalent bond in which one

of the atoms donated both electrons is called a coordinate

covalent bond. Although the properties of a coordinate

covalent bond do not differ from those of a normal covalent

bond (because all electrons are alike no matter what their

source), the distinction is useful for keeping tack of valence

electrons and assigning formal charges)

o Odd-Electron Molecules

?Some molecules contain an odd number of electrons.

Among them are nitric oxide (NO) and nitrogen dioxide

(NO2)

?Because we need an even number of electrons for complete

pairing (to reach eight) the octet rule clearly cannot be

satisfied for all the atoms in any molecule that has an odd

number of electrons

o The expanded octet:

?In a number of compounds there are more than eight

valence electrons around an atom. These expanded octets

are needed only for atoms of elements in and beyond the

third period of the periodic table.

? A measure of the stability of a molecule is its bond energy, which is the enthalpy change required to break a particular bond in 1 mole of gaseous

molecules. (bond energies in solids and liquids are affected by

neighboring molecules.)

?In many cases, it is possible to predict the approximate enthalpy of

reaction by using the average bond energies. Because energy is always

required to break chemical bonds and chemical bond formation is always

accompanied by a release of energy, we can estimate the enthalpy of a

reaction by counting the total number of bonds broken and formed in the

reaction and recording all the corresponding energy changes. The enthalpy

of reaction in the gas phase is given by:

o deltaH o = sigma(BE(reactants)) – sigma(BE(products))

o where be stands for average bond energy and sigma is the

summation sign

To further understand Bond energies, and Lewis dot structures and resonance I suggest taking a deeper look into the textbook.