Periodic Table History

1. Periodic Table History • The periodic table as we have it today has not always existed; it developed much in the same way as atomic theory did. In the early 1800’s scientists began looking for ways to classify the elements that had been discovered.

2. A German chemist, Johann Dobereiner, found, that the properties of three metals (Ca, Ba, Sr) were very similar. He also recognized that the atomic mass of Sr was about midway between that of Ca and Ba. He grouped these three elements together and called them a triad. **Two other triads he found included Cl, Br, and I as well as S, Se, and Te.

3. John Newlands saw that if he put the elements in rows of seven that there would be a repeating pattern every eighth element. This is often referred to as the Law of Octaves.

4. In 1860, Stanislao Cannizaro described a method for accurately determining the atomic weights of elements. This gave scientists a set of standardized measurements to use in all countries. Based on Cannizaro’s work other chemists began working on other ways to organize the elements.

5. Dmitri Mendeleev, a Russian chemist in the mid 1800’s, was looking for a way to organize all known elements. By organizing the elements by density, reactivity, and melting point, he came up with a periodic table that put the elements in order of increasing atomic mass. There were gaps (undiscovered elements) in Mendeleev’s table. For each of these gaps, Mendeleev predicted the properties of the missing elements.

6. In the early 1900’s, English scientist, Henry Moseley found a new pattern in Mendeleev’s table. This pattern showed that the nuclei of each element in the table increased by one positive charge (proton). This observation plus the discoveries of new elements to fill in Mendeleev’s gaps led to the modern periodic table being organized by increasing atomic number.

7. From both Mendeleev and Moseley’s work, the periodic law was developed. The periodic law states that the chemical and physical properties of the elements are periodic functions of their atomic number.

8. Horizontal rows are called periods. Vertical columns are called groups or families.

9. Family Names • Group 1 or IA= Alkali Metals (does not include hydrogen) • Group 2 or IIA = Alkaline Earth Metals. • Group 13 or IIIA = Boron family • Group 14 or IVA = Carbon family • Group 15 or VA = Nitrogen family.

10. Family Names continued: • Group 16 or VIA = Calcogens. • Group 17 or VIIA = Halogens. • Group 18 or VIIIA = Noble Gases - once known as the Inert Gases.

11. The General Layout • The middle section of the periodic table contains the Transition Metals (d-block) • The two rows below the main part of the table are the Inner Transition Metals (f-block)

12. General Layout Continued: • The metals are found on the left side of the periodic table. • The non-metals are found on the right side of the table. • The metalloids are found above and below the stair-step line started between boron and aluminum.

13. Where are the phases? • The most solid elements are located to the left and below the metalloid line. • Two elements are liquid at room temperature: mercury and bromine. • Several elements are gases at room temperature: hydrogen, helium, nitrogen, oxygen, fluorine, neon, chlorine, argon, krypton, xenon, and radon.

14. Metals • Shiny, metallic luster • Mostly solid, yet easily deformed – malleable & ductile • Good conductors of heat & electricity • Loosely held electrons Metal Activity Increases

15. Nonmetals • Solids lack luster • Many gases at room temperature • Poor conductor of electricity • Lower melting points • Tightly held valence electrons Nonmetal Activity Increases

16. Periodic Trends

17. Energy Level Trend • Energy levels – Increases down a group, does not change across a period In other words, there are more layers of electrons as move down the table (think about an onion) Remember – energy levels are regions of space in which electrons can move around the atom’s nucleus

18. Atomic Radius Trend • What is atomic radius? • It is one half the distance between 2 adjacent atoms of the same element. • It increases as you go down a group. • Because you are adding electrons to new levels • It decreases as you go from left to right in a period. • Electrons added to the same energy level are pulled closer to the nucleus because more protons are also added. ___________________ has the largest atomic radius. Increases

19. Ionic radius • Cation- a positively charged ion because it has LOST electrons. The radius always decreases with a lose of electrons • Metals tend to form cations! • Anion – a negatively charged ion because it has GAINED electrons. The radius always increases with a gain of electrons. • Nonmetals tend to form anions! • Label your table with the correct Charges!

20. Electronegativity Trend • What is electronegativity? • The measure of the ability of an atom in a chemical compound to attract electrons • Does NOT include the Noble Gases!! • Decreases as you move down a group. • Because electrons are farther from the nucleus in large atoms • Increases from left to right across a period. • Because electrons are closer to the nucleus in smaller atoms • ___________________ has the highest electronegativity. Increases

21. Ionization Energy Trend • What is Ionization Energy? • The amount of energy necessary to remove an electron from a gaseous atom. • Decreases as you move down a group. • Because electrons are easier to remove since they are farther from the nucleus • Increases as you move from left to right across a period. • Because electrons are more difficult to remove since they are closer to the nucleus • __________________ has the highest ionization energy. Increases