PERIODIC TABLE

PERIODIC TABLE

Periodic Table – an arrangement of elements according to similarities in their properties

There are 92 naturally occurring elements.

Demitri Mendeleev – drew the first periodic table; Russian chemist arranged the first periodic table of elements in 1871. Arranged by atomic mass

* The periodic table contains chemical symbol, atomic number, & average atomic mass, physical state of each element, group numbers, and electron configuration.

• Moseley: Later arranged the periodic table by atomic number. (Which is the one we use today.)

MODERN TABLE

• Periods – horizontal rows (7 total)

• Groups – vertical columns (similar physical & chemical properties)

• Metals – high electrical conductivity, luster, ductile, & malleable (Group 1 & 2A)

- Alkali Metals – Group 1A

- Alkaline Earth Metals – Group 2A

Transition Metals & Inner Transition Metals – make up Group B (1B – 8B)

Nonmetals – poor conductors, non lustrous

- Halogens – 7A

- Noble Gases – 8A

Metalloids – elements that border the stair step line

Group # = the outermost electrons

PERIODIC TRENDS

•The elements on the periodic table are arranged

periodically so that trends can be recognized…

TREND OF IONS

Ions are atoms that have lost of gained electrons.

You can determine the charge of an ion by what group it is in.

1A = +1 5A = -3

2A = +2 6A = -2

3A = +3 7A = -1

4A = +/- 4

TREND OF ELECTRONEGATIVITY

• This refers to the ability of an atom to attract a bonded pair of electrons of another

atom to it.

• Increases across the period ( left – right)

• Decreases down the group ( top – bottom)

TREND OF ELECTRON AFFINITY

• Measure of the tendency for atoms to gain electrons.

• Increases across the period; this is caused by the filling of the valence shells

• Decreases down the group; this is due to the electron entering an orbital far

away from the nucleus

TREND OF IONIZATION ENERGY

• The exact quantity of energy that it takes to remove the outermost

electron from the atom.

• Factors affecting Ionization Energy:

- effective nuclear charge

- distance from the nucleus

• Ionization energy increases across the period ( left – right) due to increased nuclear

charge

• Ionization energy decreases down the group ( top – bottom)

TREND OF ATOMIC RADIUS

• Atomic size is determined by how much space the electron takes up. It also depends on

how far its valence electrons are from the nucleus.

• The atom will be large if the electron is far from the nucleus

- size increases down a group (top – bottom)

• The atom will be small if the electron is close to the nucleus

- size decreases across the period ( left – right)

This is due to an increase in effective nuclear charge pulling them closer… the energy

level stays the same

EFFECTIVE NUCLEAR CHARGE

•

TREND OF METALLIC/NON-METALLIC PROPERTIES

• Metallic properties: elements will form cations as they lose electrons (+ve

charge)

• Non-Metallic properties: elements form anions as they gain electrons (-ve

charge)

TREND OF REACTIVITY

How likely/vigorously an atom is to react with other substances

• Metals:

- Period: decreases from left to right

- Group: increases down the group

The farther left and down you go the easier it is for electrons to be taken away.

(Higher Reactivity)

TREND OF REACTIVITY

• Non-Metals

- Period: increases from left to right

- Group: decreases down the group

The farther right and up you go the higher electronegativity – vigorous exchange of

electrons

CLASSIFICATION OF ELEMENTS

Elements can be classified into 4 groups based on electrons.

1. Noble gases: outermost s & p sublevels are filled. Belong to group 18(8A).

(Also called inert gases.)

2. Representative elements (Group A): outermost s or p sublevel is partially

filled

3. Transition metals: metallic elements in which the outermost s sublevel and near d

sublevel contain electrons. (Group B elements)

4. Inner transition metals: metallic elements in which the outermost s sublevel and

nearby f sublevel generally contain electrons. (Lanthanide & Actinide series)

LIGHT AND ATOMIC SPECTRA

• Light consists of electromagnetic waves.

• Light travels at a speed of 3.0 x 10 8 m/s.

• Amplitude: wave height from origin to crest.

• Wavelength (λ): distance between crests or troughs.

• Frequency (ν): number of wavelengths to pass a given point per unit of

time.

(units = hertz Hz, 1/s)

• c = speed of light (3.00 x 10 8 m/s)

• λ= wavelength

• ν= frequency

c=λν

Example: Calculate the wavelength of the yellow light

emitted by a sodium lamp if the frequency of the

radiation is 5.10 x 10 14 Hz (5.10 x 10 14 s-1).

c = 3.00 x 108 m/s

Frequency (ν) = 5.10 x 1014 s-1

wavelength (λ) = ??? m

• Frequency & wavelength are inversely related.

• Frequency & energy are directly related.

• Electromagnetic spectrum: series of waves at different wavelengths (radio waves, radar, microwaves, infrared, visible light, ultraviolet, x-rays, gamma rays, cosmic rays)

ATOMIC EMISSION SPECTRA

• are unique spectra of light emitted by an element when electricity is run through it or

when it is viewed through a prism. Because they are unique, they can act as an element s

fingerprint

COLORED & WHITE LIGHT

• Colored light – one particular wavelength of light or a combination of more than one

wavelength

• White Light – contains all the wavelengths of visible light (blue, red & green)

• What happens for you to see colors?

PLANCK’S CONSTANT (H)– 6.63 X 10 -34 J X S

• E = energy

• h = Planck’s constant

• ν = frequency

E = h x ν

Example: Calculate the energy (J) of a quantum of radiant energy (the

energy of a photon) with a frequency of 5.00 x 1015 s-1.

ν = 5.00 x 1015 s-1

h = 6.63 x 10 -34 J x s

Energy(E) = ??? J