The Periodic Table

chapter 6

Developing the Periodic Table In the early 1800s, scientists began

to find ways to classify the elements.

German chemist Dobereiner grouped elements based on similar properties.

English chemist Newlands, arranged elements based on atomic mass.

Developing the Periodic Table Russian chemist, Dmitri Mendeleev

organized elements into a table based on atomic mass and similar properties.

Mendeleev stated that the properties of elements are a periodic function of their atomic masses.

Mendeleev’s Periodic Table

Mendeleev’s Prediction Mendeleev’s table had several

missing elements. When these elements were discovered, they were almost exactly as Mendeleev predicted.

The following is an example of the element we know as Germanium.

Germanium is located below silicon. Mendeleev predicted its properties based on this location in his table.

Ekasilicon (Es) Germanium (Ge)

1. Atomic mass: 72 1. Atomic mass: 72.61

2. High melting pt. 2. Melting pt: 945° C

3. Density: 5.5g/cm3 3. Density: 5.323g/cm3

4. Dark gray metal 4. Gray metal

5. Will obtain from K2EsF6

5. Obtain from K2GeF6

6. Will form EsO2 6. Forms oxide (GeO2)

Modern Periodic Law Henry Moseley

revised Mendeleev’s periodic law.

He used atomic number to organize elements.

Atomic number is the basis for our current periodic law.

Periodic Table

Periodic Table Review: Rows on the

periodic table are called PERIODS

Columns on the periodic table are called GROUPS or FAMILIES

Periodic Table Review

There are 7 periods and 18 groups. Electron arrangements are repeated

in periods. Elements with similar e-

configurations are placed in the same group.

Elements in groups are also listed in order of their increasing principal quantum numbers.

Electron Configuration

Sublevel / e- capacity s 2 p 6 d 10 f 14

S - block Contains elements in Group 1, Group

2, and He from Group 18. Electrons are added to the s –

orbitals. EX: H = 1s1

He = 1s2

Li = 1s22s1

Be = 1s22s2

P - block Contains elements in Group 13,

Group 14, Group 15, Group 16, Group 17, and the remaining elements from Group 18 (except He)

Electrons are added to the p – orbitals.

Ex: B = 1s22s22p1

C = 1s22s22p2

N = 1s22s22p3

D – block Contains elements from the center

of the periodic table. These elements are called transition

metals. Electrons are added to the d –

orbitals of the transitions metals as well as La and Ac of the inner transition elements (rare earth).

F - block Contains elements from the inner

transition metals (rare earth elements)

Electrons are added to the f – orbitals.

Ex: Ce Lu Th Lr

Octet Rule Atoms with full outer levels are stable

(less reactive) For elements (except He) this stable

configuration would have eight e-.(two in the outer s sublevels and six in the outer p sublevels)

These outer eight e- (valence electrons) are called an octet.

Octet Rule Eight electrons in an outer level

render an atom unreactive. This is referred to as the Octet Rule. When atoms react with one another,

they do so to obtain a stable config. Some atoms gain or lose e- (ions)

and some share e- (molecules).

Organizing Information on the Periodic Table

Use a pen to label the following:Group 1 Alkali metalsGroup 2 Alkaline earth metalsGroup 16 ChalcogensGroup 17 HalogensGroup 18 Noble gasesSc – Uub Transition metalsLa – Lu LanthanoidsAc – Lr Actinoids

Organizing Information on the Periodic Table

Draw a stair step dark line starting between B and Al.

Label the right side: metals Label the left side: nonmetals Write METALLOID along stair step line. Label the valence e- (outer electrons). Use colored pencils to shade each group

or category a different color.

Basic Properties of Metals, Nonmetals, and Metalloids

Metals:1. Dense and shiny (luster).2. Conduct heat and electricity well.3. Have high melting points.4.Malleable and ductile.

Nonmetals:1. Generally gases or brittle solids.2. If solid, dull surface.3. Good insulators.

Metalloids:1. Properties of both metals and nonmetals.

EX: Silicon, for example, possesses a metallic luster, yet it is an inefficient conductor and is brittle.

Properties of Alkali Metals Group 1 metals Soft silver metals. Less dense than other

metals and lower melting points.

Very reactive due to large size and one loosely held valence electron.

Too reactive to be found free in nature.

Properties of Alkaline Earth Metals

Group 2 Metals Shiny silvery-white

metals Have 2 valence

electrons Not as reactive as alkali

metals but very reactive All found in the Earth’s

crust in mineral form Too reactive to be found

in free element form

Properties of Halogens Group 17 nonmetals All diatomic gases at

room temperature EX: F2, Br2

Too reactive to be found as free elements in nature

Most important group to be used in industry

Properties of Chalcogens

Group 16 nonmetals Diverse group that

includes nonmetals, metalloids, and metals

Properties of Noble Gases Group 18

nonmetals Complete octet of

valence electrons Largely

unreactive Monotomic gases

Periodic Trends

Using the Periodic Table to Predict Properties of Elements

The basis of the periodic table is the atomic structures of the elements.

Position on the table and properties of these elements arise from the e- configurations of the atoms.

Properties such as density, atomic radius, oxidation numbers, ionization energy, and e- affinity can be predicted.

Atomic Radius As principal quantum number increases,

the size of the electron cloud increases. Size of atoms increase moving down Per. Table.

Atoms in the same period have the same quantum number; however, positive charge on the nucleus increases by one proton for each element in a period. This pulls the e- cloud in tighter, decreasing atomic radius.

Predicting Atomic Radius General rule: atomic size increases

as you move diagonally from top right corner to bottom left corner.

When graphed, atomic radii demonstrates a periodic trend

Radii of ions: Ions are atoms that have gained or lost e- from the outer orbitals.

Cations: (+) Become smaller

1. Positive charged nucleus attracting fewer e-.2. Reduced the number of energy levels.EX:

Anions: (-) Become larger1. Positive charged

nucleus attracting more e- expands e- cloud.

Trends in Oxidation Numbers Our knowledge of e- configurations

and the stability of noble gases allows us to predict oxidation numbers for elements.

Oxidation numbers represent the charge an ion obtains after losing or gaining valence electrons.

1+

2+ Tend to have more than one oxidation number

3+2+ or 4+

3- 2- 1-0

3+

3+ or 4+

Two hydrogen atoms are walking down the road. One said, “I think I lost an electron!”.

“Really”, the other replied, “ Are you sure?”.

“Yes, I’m positive”.

Ionization Energy The energy required to remove an e-

from an atom. The larger the atom, the less energy is

required because the e- are farther from the positive center.

Remove the most loosely held e- is first ionization energy.

Measured in kilojoules per molekJ/mol

Ionization energy increases diagonally from bottom left corner to top right corner.

Classification based on First Ionization Energy

METAL1. Low 1st ionization energy.2. Located on left side of Periodic Table.3. Form positive ions.

NONMETAL

1. High 1st ionization energy.2. Located on the right side of Periodic Table.3. Form negative ions.

Multiple Ionization Energies

Additional e- can be lost from an atom and the ionization energies can be measured.

IONIZATION ENERGIES (kilojoules per mole)

Element 1st 2nd 3rd 4th 5th

H 1312.0

He 2372.3 5220

Li 520.2 7300 11750

Be 899.5 1760 14850 20900

B 800.6 2420 3660 25020 32660

Electronegativity

Electronegativity is the ability of an atom to capture an electron.

It increases from bottom left to top right corners.

Review

Review Based on our trends:

The most reactive metal element would be

FranciumThe most reactive nonmetal element would be

Fluorine

Electron Affinity e- affinity is a measure of an atom’s

attraction for an e-. Metals have low e- affinities. Nonmetals have high e- affinities. Chemical reactions occur between

atoms with high e- affinity and those with low e- affinity.

EX: Al + Br Al2Br3

(low) (high) (more stable)

In Summary Periodic table is a chart of elements

in which the elements are arranged based on their e- configurations which dictates their properties.

Moving down a group in the periodic table, atomic radii becomes larger because more energy levels are needed for more e-.

In Summary As the size becomes larger, the e-

are located farther away from the positive center.

This decreases the affinity of that atom to hold on to these outer e-, thus decreasing e- affinity.

Ionization energy is low because it is easy for the atom to lose these outer e-.

In Summary Moving across a period in the periodic

table, atomic radii becomes smaller because the energy levels of periods are the same but the positive centers of atoms increase. This pulls the e- cloud closer to the nucleus, making the atom smaller.

Ionization energy and e- affinity increases for these smaller atoms.

THE END