PERIODIC TABLE: DEVELOPMENT OF THE PERIODIC TABLE In the late 1790s Lavoisier developed a list of the 23 known elements

PERIODIC TABLE: DEVELOPMENT OF THE PERIODIC TABLE

• In the late 1790s Lavoisier developed a list of the 23 known elements.

DEVELOPMENT OF THE PERIODIC TABLE

• During the 1800s many more elements were discovered.

• Scientists used electricity to separate compounds.

• By 1870 there were 90 known elements.


• Chemists were overwhelmed by the huge volume of information.

• A new way of categorizing the information was needed.

• A breakthrough came in 1860 when chemists agreed upon a method of accurately determining the atomic masses of the elements.


• The first attempt was made by John Newlands.

• He noticed the properties of the elements repeated every eighth element.

• He called this the Law of Octaves.



• This was not widely accepted because it did not work for all the elements.

• But, Newlands was correct in that the properties of elements do repeat in a periodic way.


• Next came Dmitri Mendeleev


• Dmitri Mendeleev also arranged the elements by increasing atomic mass.

• He did not limit the length of his rows

• He noticed the elements fell into columns with similar properties.


• Mendeleev also left empty spaces to account for elements that had not yet been discovered.


• There were problems with Mendeleev’s table because new elements were discovered and atomic masses better determined that placed them out of order.

• Henry Moseley had discovered that the atoms of each element contain a unique number of protons in their nuclei.

• He arranged the elements in order of increasing atomic number – this adjustment fixed the periodic table.


• The modern periodic table is based on the periodic law which states that the physical and chemical properties of the elements tend to change with increasing atomic number in a periodic way.

MODERN PERIODIC TABLE

• Each element has its own box containing information about the element.

• The elements are arranged in order of increasing atomic number into a series of columns called groups, or families, and rows called periods.

REGIONS OF THE PERIODIC TABLE

• There are five major regions of the periodic table:

metals nonmetals

metalloids noble gases

hydrogen


• The metals are the largest region of the periodic table. Elements within this region are solid, generally lustrous, ductile, and malleable. They are also excellent conductors of heat and electricity.

REGIONS OF THE PERIODIC TABLE• Nonmetals make up the second largest region of the

table. These elements have a variety of physical states and properties. In general, they are poor conductors of heat and electricity, and are brittle and non-lustrous.


• Metalloids are elements that have properties of both metals and nonmetals. These elements are sandwiched in between metal and nonmetal regions along the “stair step”.

METALLOIDS• Boron (B) – used in fiberglass, glassware, ceramics,

polymers, detergents, insecticides• Silicon (Si) – semiconductors• Germanium (Ge) – semiconductors• Arsenic (As) – wood preservative, insecticide,

semiconductors• Antimony (Sb) – batteries, low friction metals, flame-

proofing, ceramics, paints, glass, pottery• Tellurium (Te) – blasting caps, added to copper and

stainless steel to improve machinability, reduce corrosivity of lead

• Astatine (At) – radioactive tracer in cancer treatment

REGIONS OF THE PERIODIC TABLE• The metalloids must be memorized since the periodic

tables we use are not color-coded.

• The noble gases are gases located in Group 18. These gases do not normally react with anything.

FAMILY CHARACTERISTICSGroup 18 – Noble Gases- Once called the inert gases- No compounds of He, Ne or Ar- Compounds of Xe, Kr and Rn- Have full orbitals in the highest energy level; this is

called an octet.- Their electron configuration is very stable

- Atoms from the other 17 groups either gain or lose electrons to achieve a noble gas electron configuration

Group 1: Alkali Metals• Have metallic properties – soft, shiny, highly reactive,

can be cut with a knife, react with oxygen, good conductors of heat and electricity

• Single electron in the highest energy level

• By losing this electron the alkali metal achieves a noble gas electron configuration and a positive charge of +1.

Group 2: Alkaline Earth Metals• Also have metallic properties – harder, stronger, denser

than Group 1 metals.

• Less reactive than Group 1 metals.

• Need to lose 2 electrons to achieve a noble gas electron configuration. Results in a positive charge of +2.

Groups 3 – 12: Transition Elements• NOT A FAMILY – just the name of a section of the

periodic table.

• These are metals, but they are not as reactive as Group 1 and Group 2 metals.

• Electron configurations are unusual and sometimes do not come out as predicted by the diagonal rule.

• Will lose electrons from both the s-orbitals and the d-orbitals to form ions with various positive charges.

Bottom Rows: Inner Transition Elements• NOT A FAMILY – just the name of a section of the

periodic table.

• Each row has its own name: Lanthanides and Actinides

• The lanthanides are reactive, shiny metals.

• The actinides are usually radioactive due to an unstable arrangement of their protons and neutrons.

Groups 13 – 18: Main Block Elements• NOT A FAMILY – just the name of a section of the

periodic table. These are also known as the representative elements.

• They represent a wide variety of chemical and physical properties.

• The metals will lose electrons and form positive ions while the nonmetals will gain electrons to form negative ions.

• The metalloids can both lose and gain electrons.

Group 17: Halogens• The halogens combine easily with metals, especially the

alkali metals, to form compounds known as salts.

• The halogens are the most reactive nonmetals.

• Their electron configuration is one electron short of being a noble gas electron configuration.

• Will gain an electron to form an ion with a charge of -1.

Hydrogen

• Most common element in the universe.

• It has only one electron and it reacts very rapidly with other elements – behaving as a metal.

• It can gain an electron and attain a charge of -1.

EXAMPLES• For each of the following elements, list the region of the

periodic table from which it comes and list the name of its family (if applicable):

• Ba• Metal, alkaline earth• Sb• Metalloid• I• Nonmetal, halogen• Rn• Noble gas, noble gas

EXAMPLES• For each of the given elements, list two other elements

with similar chemical properties.• Fe• Ru, Hs, Os• Se• S, O• Br• I, Cl, F• Rb• K, Na

CLASSIFICATION OF THE ELEMENTSOrganizing by Electron Configuration

• Elements within the same group have the same electron configuration in their outermost energy level.

• These electrons are called valence electrons.

• Atoms in the same group have similar chemical properties because they have the same number and arrangement of valence electrons.

Organizing by Electron Configuration• The energy level of an element’s valence electron

indicates the period, or row, on the periodic table where it can be found.

The s-, p-, d-, and f- Block Elements• The s-block consists of Groups 1 & 2 along with

hydrogen and helium. In this block, the valence electrons occupy only the “s” orbitals.

• The p-block elements consist of Groups 13 – 18. As one progresses from left to right, one more electron is added to the “p” orbital until it is filled with six electrons.

• The noble gases exhibit great stability because both their “s” and “p” orbitals are filled.

The s-, p-, d-, and f- Block Elements• The d-block contains the Transition Elements and it is

the largest of the electron blocks.

• These elements are characterized by a filled outermost “s” orbital and either partially or completely filled “d” orbitals.

• The f-block elements are found in both the lanthanide and actinide series.

EXAMPLE• Strontium has an abbreviated, noble gas, electron

configuration of [Kr] 5s2. Without using the periodic table, determine:

• Group• 2• Period• 5• Block• s

EXAMPLE• Write the abbreviated, noble gas, electron configuration

of the element in Group 12, Period 4.

• Zn: [Ar] 3d104s2

PERIODIC TRENDS

• The periodic law states that the physical and chemical properties of the elements are periodic functions of their atomic numbers.

• Because the elements are arranged side by side in order of increasing atomic number, one can view certain important vertical and horizontal trends.

Atomic Radius• Recall from Rutherford’s experiment that the nucleus

was found to occupy only a small fraction of the atom’s entire volume.

• It is the electron cloud surrounding the nucleus that determines the boundaries of the atom.

Atomic Radius• Because the electrons travel in this “cloud region”, it is

difficult to measure the size of an atom.

• The atomic radius is determined in two primary ways. One is to measure the distance between centers of like atoms in a diatomic molecule. The other is to measure the bond lengths of atoms in compounds.

Atomic Radius• Looking at the periodic table one can notice things that

affect the size of an atom.

• As one moves down through a group a new principal energy level is added. Each new level is physically farther from the nucleus.

• As each energy level is added, the energy levels and electrons closer to the nucleus shield electrons in the new energy level from the pull of the nucleus. This shielding effect allows the outer electrons to drift farther away from the nucleus.

Atomic Radius• When crossing a period, row, from left to right each atom

gains both a proton and an electron.

• But, no new principal energy level is added so the new electrons enter the same energy level as all the others.

• There is no shielding.

• The additional protons in the nucleus provide more pull on the electron cloud and bring it closer to the nucleus.

Atomic Radius• THE TREND:

Atomic radius increases from top to bottom in a group and decreases from left to right across a period.

Atomic Radius

Ionic Radius• Atoms can gain or lose electrons to form charged

particles.

• These particles are called ions.

• When atoms lose an electron they form a positively charged ion and become smaller in size.

• Positive ions are smaller than their parent atom.

Ionic Radius• When atoms gain an electron they form a negatively

charged ion and become larger in size.

• Negative ions are larger than their parent atom.

Ionization Energy• To form a positive ion, an electron must be removed from

a neutral atom.

• The electron is removed from the outermost energy level of the atom, known as the valence shell.

• The electrons in that “shell” are known as valence electrons.

• This requires energy to overcome the attraction between the positive nucleus and the negatively charged electron.

Ionization Energy• This energy is known as the ionization energy.

• Ionization energy is defined as the energy required to remove an electron from a gaseous atom.

• The energy required to remove the first electron from an atom is called the first ionization energy.

• The process is:

• ATOM + IONIZATION ENERGY → ION+ + e-

Ionization Energy• As one moves down a group (column) it becomes easier

to remove electrons because they are farther from the nucleus and the shielding effect increases.

• As one moves left to right across a period (row) it becomes more difficult to remove electrons because there is no increase in the shielding effect, the positive charge in the nucleus increases, and the electrons are closer to the nucleus.

• Trend:• Ionization energy decreases from top to bottom down a

group and increases from left to right across a period.

Ionization Energy

Ionization Energy

Electronegativity• The electronegativity of an element indicates the relative

ability of an atom to attract electrons to itself when the atom is involved in a chemical bond.

• These values cannot be measured. In fact, they were developed by a chemist, Linus Pauling, as a way to explain chemical bonding in molecules. They are based on fluorine being the element with the strongest attraction for electrons, so it is a relative scale.

Electronegativity• The noble gases are ignored since they are basically

inert.

• Fluorine is the most electronegative element and is assigned a value of 3.98 while cesium and francium are the least electronegative elements.

• In a chemical bond, the atom with the greater electronegativity more strongly attracts the electrons in that bond.

Electronegativity• Trend:• Electronegativity decreases from top to bottom in a

group (column) and increase from left to right across a period (row).

Electronegativity

Example:• In each of the following sets of elements, which element

would be expected to have the highest ionization energy?

• a. Cs, K, Li

Li• b. Ba, Sr, Ca

Ca• c. I, Br, Cl

Cl• d. Mg, Si, S

S

Example:• Arrange the following sets of elements in order of

increasing atomic size.• a. Sn, Xe, Rb, Sr

Xe < Sn < Sr < Rb

• b. Rn, He, Xe, Kr

He < Kr < Xe < Rn

• c. Pb, Ba, Cs, At

At < Pb < Ba < Cs

Example:• In each of the following sets of elements, indicate which

element has the smallest electronegativity value.

• a. Na, K, Rb

Rb• b. S, Na, Si

Na• c. P, N, As

As• d. O, N, F

N

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PERIODIC TABLE: DEVELOPMENT OF THE PERIODIC TABLE In the late 1790s Lavoisier developed a list of the 23 known elements