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Periodic Law

5.1 History of the Periodic Table

Objectives:

1. Explain the roles of Mendeleev and Moseley in the development of the periodic table. 2. Describe the modern periodic table 3. Explain how the periodic law can be used to predict the physical and chemical properties of elements. 4. Describe how the elements belonging to a group of the periodic table are interrelated in terms of atomic number.

By 1860, more than 60 elements were known but no way to accurately measure atomic mass or number of atoms of an element in a particular compound.

In September 1860, a group of chemists assembled at the First International Congress of Chemists in Karlsruhe, Germany, to settle the issue of atomic mass as well as some other matters that were making communication difficult.

At the conference, Stanislao Cannizzaro presented a convincing method for accurately measuring the relative mass of atoms. This agreed upon standard enabled chemists to agree on standard values for atomic mass and initiated a search for relationships between atomic mass and other properties of elements.

Russian Chemist Demitri Mendeleev included the new atomic mass values in his textbook on chemistry.

Mendeleev tried to organize the known elements by grouping them by chemical and physical properties. Mendeleev was the first to recognize that when elements were arranged in order of atomic mass, certain similarities in their chemical properties appeared in regular intervials, or PERIODS. (Periodic referring to regularly repeating.)

Mendeleev’s chart – first published in 1869. He predicted the existence of three elements which had been discovered by 1886. This success convinced most chemists to accept the periodic law.

Moseley and the Periodic Law

Moseley: In 1911 Henry Moseley was working with Ernest Rutherford. Moseley noticed that spectra of elements fit into a previously unrecognized pattern. This pattern was governed by atomic number rather than atomic mass. This was consistent with Mendeleev’s ordering of the elements and grouping by properties. Page 2

Mendeleev’s Periodic Law: The physical and chemical properties of the elements are periodic functions of their atomic numbers.

Modern Periodic Table

2 charts

The periodic table is an arrangement of the elements in order of their atomic numbers so that elements with similar properties fall in the same column or group.

Noble Gases

The addition of the Noble Gases was a significant addition to the periodic table.

In 1894 Argon: John William Strutt (Lord Rayleigh) and Sir William Ramsay discovered Argon. They found it in air but it had been previously escaped discovery because it is totally unreactive.

In 1868, Helium was first discovered as a component of the sun based on emission spectra. In 1895 Ramsay showed that He existed on earth as well.

In 1898 Ramsay also discovered Kr and Xe.

Rn was discovered in 1900 by German scientist Friedrich Ernst Dorn.

Lanthanides

Lanthanides are 14 elements with atomic numbers from 58 (Cerium) to 71 (Lutetium). These elements are all very much alike and that made the discovery of each element very tedious.

Actinides

Actinides are the elements from 90 (thorium) to 103 (lawrencium.)

Only through Np are found in nature. The remainder are man made.

Periodicity

Periodicity is a direct result of the way that electrons are arranged around the nucleus.

See the chart.

Make the point that the difference in atomic number is because of filling s, p, d, and f orbitals. Page 3

Electron Configuration and the Periodic Table

Objectives: 1. Describe the relationship between electrons in sub-levels and the length of each period of the periodic table. 2. Locate and name the four blocks of the periodic table. Explain the reasons for these names. 3. Discuss the relationship between group configurations and group numbers. 4. Describe locations in the periodic table and the general properties of the alkali metals, the alkaline-earth metals, the halogens, and the noble gases.

Periods and Blocks of the Periodic Table

Periodic table can be split into four blocks: s-block p-block d-block f-block

Period is the row that a given element is in.

Group is the column that a given element is in. Elements within a group have similar chemical properties.

The period of an elements can be determined from the electron configuration:

For example, arsenic, As, has a configuration [Ar]3d104s24p3. The 4 in the 4p indicates that As is in the 4th period. s-Block Elements: Group 1 and Group 2

Group 1 are the alkali metals: Li, Na, K, Rb, Cs, Fr

Silvery, soft, react with water to produce hydrogen gas.

Group 2 are the alkaline earth metals: Be, Mg, Ca, Sr, Ba, and Ra.

Harder, denser and stronger than alkali metals. Not as reactive but still reactive.

Exceptions: Hydrogen and Helium

Hydrogen should be in Group 1 but has properties very different than alkali metals. We put it there because it does react chemically like the alkali metals. Page 4

He is put in Group 18 because of its inert nature. But unlike all the other Group 18 elements, He has no p electrons and has a closed 1s2 shell.

Example: Give the group, period and block in which the element with the electron configuration of [Xe]6s2 Answer: Group 2, Period 6, s-block d-Block Elements: These are the elements in the middle of the periodic table and are often called the transition elements.

Generally filling the d orbitals. However, there are exceptions (like Cr and Cu)

Ni [Ar]3d84s2 Pd [Kr]4d105s0 Pt [Xe] 4f145d94s1

However, in all cases s electrons + d electrons = group number

Metals, good conductors, not as reactive as alkaline-earth metals. Some extremely unreactive like palladium, platinum, gold found in nature as the elements.

Example: Identify the period, block, group for: [Kr]4d55s1 period 5, d-block, group 6 p-Block Elements: These are the elements from group 13 through group 18 (6 groups, 6 p electrons, no coincidence)

Includes all non-metals except hydrogen and helium and all six metalloids: boron, silicon, germanium, arsenic, antimony, and tellurium.

Elements in Group 17 are called the halogens (fluorine, chlorine, bromine, iodine, astatine)

React vigorously with most metals. F and Cl colored gases, Br orange liquid, I a dark purple solid.

Metaloids: semiconducting, brittle solids, some properties of metals and some properties of non metals.

Metals of the p-block are generally harder and denser than the s-block alkaline-earth metals but softer and less dense than the d block metals.

Table on Relationship among group numbers, blocks, and electron configurations

Example: Write the electron configuration for the Group 14 element in the second period. Name the element and tell whether it is a metal or non-metal. 18-14 = 4 therefore there must be 2 p electrons. Thus for second period: 1s22s22p2 This is carbon a non-metal. Page 5 f-Block Elements: Lanthanides and Actinides These are the elements are shiny metals similar to alkaline-earth metals, Group 2.

These are filling f orbitals.

Between Lanthanum and Hafnium in the sixth period And Thorium to Rutherfordium (only first four Th through Np are found in nature)

Example: Name the block and group, name the element, metal/non-metal/metalloid, high or low reactivity: [Xe]4f145d96s1 Period 6, d-block, group 10, Pt, low reactivity.

Video segment on Blocks

5-3 Electron Configuration and Periodic Properties

Objectives: 1. Define atomic and ionic radii, ionization energy, electron affinity, and electronegativity. 2. Compare the periodic trends of atomic radii, ionization energy, and electronegativity, and state the reasons for these variations. 3. Define valence electrons, and state how many are present in atoms of each main- group element. 4. Compare the atomic radii, ionization energies, and electronegativities of the d- block elements with those of the main-group elements.

Atomic Radii

Atomic radius is defined as ½ the distance between the nuclei of identical atoms when they are bonded together.

Atomic Radii generally decrease from left to right because of the larger positive charge in the nucleus.

Atomic Radii generally increase in size going down group in the periodic table.

Specific example: Gallium – size smaller than aluminum because it has 10 more protons in the nucleus than expected because gallium has a d orbital and aluminum does not.

Ionization Energy

An electron can be removed from an atom if enough energy is supplied:

A + energy  A+ + e-

This forms an ion: an atom or a group of atoms which have either a positive or negative charge. Page 6

A process which forms an ion is called ionization.

The energy required to remove one electron from a neutral atom is called the first ionization energy, IE1 and is given in Kilojoules/mole

In general the first ionization energy increases across a period and decreases down a group.

Show chart from ATE of first ionization energy vs. Atomic number.

Energy required to remove additional electrons are referred to as: Second ionization energy (IE2) And Third ionization energy (IE3)

Removal of each successive electron is progressively harder due to the stronger attraction of the positive nucleus.

If you reach a noble gas shell configuration, then the next electron is extremely hard to remove.

Example: Consider two main-group elements A and B. Element A has a first ionization energy of 419 kJ/mole. Element B has a first ionization energy of 1000 kJ/mole. For each element decide if it is more likely to be in the s-block or the p-block. Which element is more likely to for a positive ion? A in s-block, B in p-block, A more likely to form a positive ion.

Electron Affinity

The energy change when a neutral atom acquires an electron is called an atom’s electron affinity

A + e-  A- + Energy

By convention, when energy is released it is represented by a negative number.

Group 17 elements acquire an electron most readily.

In general, electrons add with greater difficulty going down a group. Competition between more positive nucleus (higher affinity) and larger nuclei (lower affinity). Larger usually predominates.

Adding a second electron is always positive in energy (requires energy to happen).

Ionic Radii Page 7

A positive ion is called a cation (attracted to the cathode, - electrode)

A negative ion is called an anion (attracted to the anode, + electrode)

Cations are smaller and anions are larger than the corresponding atomic radii, follows the same trends as atomic radii.

Show video clip on cations and anions.

Valence Electrons

Electrons available to be gained, lossed, or shared in the formation of chemical compounds are referred to as valence electrons.

Example: the electron lost from Na to form Na+ is a valence electron. Page 8

Electronegativity

Linus Pauling invented a scale of numerical values reflecting the tendency of an atom to attract electrons.

Electronegativity is a measure of the ability of an atom in a chemical compound to attract electrons.

In general electronegativities tend to increase across a period.

Electronegativities tend to either decrease down a group or remain about the same.

Video clip on Electronegativity

Example: Among the elements, gallium, bromine, and calcium, which has the highest electronegativity? Explain why in terms of periodic trends.

The elements are all in the 4th period. Bromine has the highest atomic number and is farthest to the right in the period. Therefore, it should have the highest electronegativity because electronegativity increases across periods.

Periodic Properties of d- and f-block Elements

Atomic radii of d-block elements generally decrease across periods. The change is usually small because the (n-1)d electrons shield the outer electrons.

Ionization energies for the d-block and the f-block generally increase across the periods.

The order in which electrons are removed from the d-block and the f-block elements is exactly the reverse of the order given in the electron-configuration notation.

For example, Fe:

Fe(0) [Ar] 3d64s2

Fe(1) [Ar] 3d64s1

Fe(2) [Ar] 3d6

Fe(3) [Ar] 3d5