Chapters 11.1-11.6

11.1 Periodic Trends in Atomic Properties

   The properties of traits of substances can

be physical or chemical. Chemical properties

are the ability of a substance to form new

substances by reacting with other substances

or decomposition. Physical properties are non

composition alterations such as color, taste, odor,

density, melting and boiling point.

The first chemical property trend we see when

looking at the periodic table is that of the metal,

nonmetal categories. On either side of the stairstep

are metalloids. Metals tend to lose electrons and

nonmetals tend to gain electrons.

*Hydrogen is neither a nonmetal nor a metal.

The atoms radius is another property within the

periodic table. The radii of the atoms decrease

down each group due to the additional energy

levels. Now if we move across going to the right,

staying in the same energy level, the atomic radii

decreases. This is due to a the additional protons

(positive charge) pulling the electrons closer to

the nucleus and decreasing it's size.

The increase in pull of the protons and electrons

to the nucleus increase the amount of energy

required to remove an electron. The amount of

energy required to remove an electron is known

as ionization energy. Ionization energy decreases

as atomic radii increases. As we move across the

period the ionization increase as atomic radii

decreases.

11.2 Lewis Structure of Atoms

   Lewis structure or diagram "represents valance

electrons and covalent bonds in an atom or

molecular species". Named after the chemist

Gilbert N. Lewis who suggested using dots for the

valance electrons and the chemical symbol of

the element.

Exp:

Electrons of the valance of nitrogen will fill

the p orbit one by one, then begin to pair.

Exp:

11.3 Ionic Bond/Electron Transfer

   The most stable position of an element is that

of the configuration of it's nearest noble gas. That

is, completion of the valance or an octet. An octet

is a structure of eight electrons in its outer shell.

We will look at K, and Br to see how this occurs.

Now Bromine has an octet at the energy level

4 which is isoelectronic with Krypton,

The K⁺ ion and the Br ^- ion are strongly attracted

to each other due to opposite charges and the

formation of KBr is the result.

11.4 Predicting Formulas of Ionic Compounds

   Ionic bonds form when strong electrostatic

forces hold ions in a fixed position. In section

11.3 we gave you an example of an ionic

compound forming when a metal donates an

electron and a nonmetal accepts it. Applying

this concept to other representative elements

leads to this chemical principle: " In almost all

stable compounds of representative elements, each

atom attains a noble gas electron configuration".

That is to say that it becomes isoelectronic with

the nearest noble gas. Even if the ratio of ions is

not 1:1 the overall outcome of the compound is

electrically neutral.

Exp:

Calcium has two electrons to donate and it will

combine with the elements in group 7A (halogens)

in a 1:2 ratio. Similarly any element in group 2A

will combine with chlorine in a 1:2 ratio.

11.5 Covalent Bond/Sharing Electrons

   Another bonding concept was introduced by

G.N. Lewis in 1916 in which a covalent bond

occurs between pairs of shared electrons.

Molecular compounds consist of covalent bonds,

whose ultimate structural unit is a molecule.

* Ionic bonds are not molecules.

The simplest covalent bond and the simplest

molecule appear in hydrogen, H₂.

Equal sharing is common among identical

noble gases. Exp; O₂, N₂, F₂, Br₂, and I₂.

The easiest way to indicate a covalent bond

is to draw a line between the two atoms.

11.6 Electronegativity

   What happens when two different kinds of

atoms share electrons is dependant on

electronegativity. Electronegativity is a scale

of the ability of one atom in a covalent bond to

attract the shared electrons more strongly than

the other atom. In the case of identical atoms

in a covalent bond the attraction is mutual and

the bond is termed nonpolar. However, when

two different atoms share electrons, one atoms

electronegativity can be greater and the term

is polar. In other words the distribution of the

bonding electron charge is unsymmetrical.

When trying to determine the greater

electronegativity it is helpful to know that

nonmetals are greater than metals and that

fluorine has the greatest of the nonmetals.

Electonegativity decreases from top to

bottom and increases from left to right.

Exp:

11.7 Lewis Structure of Compounds

   We have already touched on accomplishing

Lewis structures in the last three sections. Now

lets add steps to follow for more complex

diagrams.

Step 1 Find the total number of valance electrons
of all the atoms. If ions, add for each negative
charge or subtract for each positive charge.

Step 2 Write down possible skeletal arrangement,
usually the single atom in the center. Connect
covalent bonds with a dash or two dots.

Step 3 Subtract the number of electrons in the
bonds you made from the total amount of valance
electrons of all atoms.

Step 4 Place electrons left around each atom until
each of those atoms has a electron configuration
of a noble gas(octet).

Step 5 If there is not enough electrons to give each
atom 8 electrons each, you may need to change
single bonds to double bonds(count as four electrons).

Exp: PBr₃

Step 1 Phosphorus has five electrons and bromine

has seven each for a total of 26 electrons available.

Step 2

Step 3 Do the subtraction 26-6+20 electrons left

Step 4 Filling valance of leftover electrons

Step 5 is not need because all atoms have an octet.

* Lets try one more but this time more complex.

Exp: HCO₃^-

Step 1 Hydrogen has one electron, carbon has four,

six from each oxygen, and one for the negative charge.
                 1+4+3(6)+1=24electrons available

Step 2

Step 3 Do the subtraction 24-8=16 available electrons

Step 4

Step 5 Use square brackets to show this as an ion

and look for a double bond placement

Now we could have formed a double to the carbon

with either of the oxygen atoms. There are two

possible Lewis structures for HCO₃^-.

The formation of a structure by moving lone pairs

from one or more identical outer atoms is called

resonance.