Chapters 10.1-10.6

10.1 Brief History

   Chemists strive to describe the atom based on

several theories and discoveries. In section 5.2,

Dalton's atomic theory and the modifications

of his theory that still hold true today. Michael

Faraday discovered certain substances(ions)

conduct electrical current when dissolved in

H₂O. Up until now we have looked at the atoms

as individuals. But how does the structural

behavior of a atom relate to the periodic table?

Within the periodic table elements have similar

behaviors.

10.2 Electromagnetic Spectrum

   Paving the way for the modern model of the

atom is the study of energy and light. The

transfer of energy through space can occur

within what is known as the electromagnetic

spectrum.

Scientists describe light in one of two ways, having

properties of a wave-like nature

and as photons (tiny packets of energy). For this

course you need only basic terminology of wave

mechanics in order to understand Bohr's contribution

to the modern concept of the atom.

1. Wavelength(lambda, ) "is the distance between

consecutive peaks(or troughs in a wave).

2. Frequency(v, expressed in hertz) how many waves

pass a particular point per second.

3. Velocity(c, speed of light) how fast a wave moves

through space. c= 3.00x10⁸ m/s.

10.3 The Bohr Atom

    Neil's Bohr(1912-1913) using the earlier

works of Max Planck's energy quanta, developed

a relationship between quanta and his studies of

the line spectra of hydrogen.

This relation introduced quantified electron energy

levels which appear in modern theory as principle

energy levels or principle quantum number, n.

10.4 Energy Levels of Electrons

    Definition of the quantization of energy, within

atoms, states that electrons may be at one of the

energy levels but not between two such levels.

Description of electron energy, uses four quantum

numbers.

1) Principle Energy level identified by the principle

quantum number, n= 1, < n=2,< n=3...ect.

2) The sublevel, (for each principle energy level) are

designated by s,p,d,f.

3) The orbitals, each orbital has a specific number of

them per sublevel, one for s, three for p, five for d,

and seven for f.

4) The allowed amount of electrons in a orbital. The

Pauli exclusion principle limits the amount of

electrons possible within an orbital (that is 2). An

orbital may be occupied by one electron occupied

by two electrons, or it may be unoccupied.

10.5 Atomic Structure first 18 Elements

    Every element has a certain energy level in it's

stable position. Within that energy level is sublevels

ie: s< p< d< f ect.... Each sublevel has specific

amounts of orbitals associated with it( s, has 1,

p, has 3, d, has 5, and f has seven). These orbitals

only hold a maximum of two electron per orbital.

Schematic Table

* Please follow along with a copy of the periodic table.

Starting with H and moving across to He we have our

first energy level. Hydrogen only has one electron

and this electron is in the sublevel s. The s

sublevel is capable of a maximum of two electrons

and thus this sublevel is filled at He .

Moving on to Li and across , is the next energy

level n=2. The sublevels get a little tricky so you

may want to refer to the table we placed in 10.4.

Here we will show the electrons in boxes and

also write out the electron configuration.

B, C, N, O, F, and Ne have a p sublevel and this

sublevel will be full at Neon. *If your concerned

that our illustration looks as if the s sublevel has

more than one orbital and more than two

electrons, that is incorrect. Remember that there

is one s sublevel and two electrons max. per

principle energy level. We also include these

preceding energy levels in our electron

configuration unless otherwise abbreviated.

Exp:

This element has 5 electrons total and three valance

electrons from the principle energy level two.

However when we write out an electron configuration

we include all of the previous elements configuration

which are in fact a part of entity. *Remember

Borons p sublevel has three orbitals with six

electrons capacity even though it is not full they

are still there.

Try to work out the electron configuration for

the first eighteen elements using the a ordinary

periodic table. Then if you get confused look

at the one provided in section 10.4. When you

have finished move down this page to check

your answers.

Electronic Configuration Answers

10.6 Electron Structures/ Periodic Table

    Valance Electrons, the ones in the highest

principle energy level of an atom, participate

in bonding atoms to form compounds. The more

stable(subshells orbitals filled) an atom, the

more energy must put put in to remove an

electron ie: Ionization. Lets break down the

periodic table into identifiable parts and relate

them to electronic structures. Moving down

the horizontal rows we have periods and

as mentioned earlier increasing principle

energy levels. However, elements with similar

chemical properties are found vertically on

the table. These are called groups or families

and are numbered 1-18. The groups are also

separated by letters A and B. The A groups

are called the transitional elements and the B

groups are transitional elements. You will also

notice a corresponding number associated with

the A and B lettering. * This is the number of

electrons in the atoms valance.

Alkali Metals /Group 1A

Reactivity increases as you move down
Lost valance electrons/isoelectronic with noble gas
Good conductors of heat and electricity
Density increases as atomic number(e-) increase
Boiling and melting point decrease down

Alkaline Earths/Group 2A

Readily gives up the 2 valance electrons
Lost valance electrons creates noble gas configuration
Reactivity increases as you go down table

Halogens/Group7A(salt formers)

Only the first four elements F, Cl, Br, I
Reactivity decreases as you move down
Gains electrons/noble gas configuration
Density, melting, boiling point increase w/atomic

Noble Gas/Group 8A

Very unreactive
Filled valance
Resist gaining or losing electrons

Groups 1 and 2 tend to lose their valance

electron(s) thus leaving the neutral element

with a net charge of 1⁺

Exp: Na, Z(atomic number) =11

        Removing an electron leaves

        Na⁺ (monoatomic cation)

Groups to the right of the stairstep, especially

in 16 and 17 tend to gain electrons leaving them

with 1^- charge.

Exp: F, Z(atomic number) =9

        gaining an electron

        F^- (monoatomic anion).