The Periodic Table

1. The Periodic Table • A few elements, including copper, silver, and gold, have been known for thousands of years • There were only 13 elements identified by the year 1700. • Chemists suspected that other elements existed. • As chemists began to use scientific methods to search for elements, the rate of discovery increased.

2. The Periodic Table • Early chemists attempted to organize the known elements • Some used the properties of the elements • Dobereiner was a German chemist who published his classification of the elements • He organized the elements into triads

3. The Periodic Table • A triad is a set of three elements with similar properties. • The elements shown here formed one triad. Chlorine, bromine, and iodine may look different, but they have very similar chemical properties. • Dobereiner noted a pattern in his triads. One element in each triad tended to have properties with values that fell midway between those of the other two elements.

4. The Periodic Table • In 1869, a Russian chemist and teacher, Dmitri Mendeleev, published a table of the elements. • The organization he chose was a periodic table. • Elements in a periodic table are arranged into groups based on a set of repeating properties. • This concept is known as periodicity.

5. The Periodic Table • Mendeleev left spaces in his table • He predicted that elements would be discovered to fill those spaces, and he predicted what their properties would be based on their location in the table.

6. The Periodic Table • Mendeleev’s table was so successful because chemists were able to make predictions from it. • However his table had one error • It was arranged by increasing atomic mass.

7. The Periodic Table • Henry Moseley modified Mendeleev’s table • He arranged the elements by increasing atomic number. • This led to the Periodic Law • The physical and chemical properties of elements are periodic functions of their atomic numbers.

8. The Periodic Table • There are seven horizontal rows or periods in the table • Each period corresponds to the energy level • There are eighteen vertical columns or groups in the table

9. The Periodic Table • Although Mendeleev’s table only had 60 elements, today’s table has 118 elements • Many elements have been discovered since his original work • The noble gases - due to their unreactivity • The lanthanides and actinides series – many are man-made

10. The Periodic Table • Periodicity can be observed in the periodic table • Each group has similar properties • The electron configuration tells an element’s position in the periodic table

11. The Periodic Table • The elements can be grouped into three broad classes based on their general properties. Three classes of elements: • Metals • Nonmetals • Metalloids

12. The Periodic Table • The majority of elements are metals • Metals have three key properties • Shiny or luster • Flexible (malleable – hammer into a sheet and ductile – drawn into a wire) • Good conductor of energy (electricity and heat)

13. The Periodic Table • Although there are fewer nonmetals, they are more abundant on Earth • Nonmetals have three key properties • Dull • Brittle • Poor conductor of energy

14. The Periodic Table • Metalloids have properties of metals and nonmetals • Their properties can be changed by conditions • They are found along the stair step of the periodic table • Silicon is the most famous metalloid • It is responsible for computer chips

15. The Periodic Table • Most elements are solids • There are two elements that are a liquid at room temperature • Bromine • Mercury • There are 11 elements that are gases at room temperature

16. The Periodic Table • s & p block elements are called main group (representative) elements • s block • Group 1 – Alkali metals • Most reactive metals • So reactive, not found in nature as elements • Group 2 – Alkaline Earth metals • Less reactive than group 1 metals

17. The Periodic Table • p block • Group 13 – 16 Mixed group • Group 17 – Halogens • Most reactive nonmetals • Group 18 – Noble Gases • Completely unreactive nonmetals

18. The Periodic Table • d block • Groups 3 – 12 – Transition metals • Less reactive than groups 1 & 2 • f block • Bottom two rows • Known as the lanthanides and actinides (inner transition metals)

19. The Periodic Table • There are several trends in the periodic table • Atomic radii • Valence electrons • Ionic radii • Ionization energy • Electronegativity

20. The Periodic Table • Atomic radii is one half the distance between the nuclei of identical atoms that are bonded together

21. The Periodic Table • Atomic Radius Trend: • Down a group – atomic radii increases • This happens because of the increased number of energy levels • The energy levels shield the electrons from the attraction of protons in the nucleus • Across a period – atomic radii decreases • This happens because as more electrons are added to the same energy level, those electrons are pulled closer due to the increased number of protons in the nucleus • Largest atomic radii – francium • Smallest atomic radii – fluorine

22. The Periodic Table • Valence electrons are the electrons found in the outermost energy level • These are the electrons available to be gained, lost, or shared • All atoms want 8 valence electrons or a full outer energy level • Noble gas electron configuration • Valence electrons determine the chemical properties of the atom

23. The Periodic Table • Valence electrons can be represented using Lewis Dot Diagrams

24. The Periodic Table • Atoms are neutral because there are equal numbers of both protons and electrons • Sometimes atoms can gain or lose electrons to form ions • An ion is an atom or group of atoms that has a positive or negative charge • Losing electrons results in a positive ion called a cation • Gaining electrons results in a negative ion called an anion

25. The Periodic Table • Metals (left side of the table) form cations • Cations are smaller than their atom counterparts because they are losing an electron (and sometimes an energy level) • More positive charges have a greater pull on less negative charges

26. The Periodic Table • Nonmetals (right side of the table) form anions • Anions are larger than their atom counterparts because they are gaining an electron • Less positive charges cannot pull in the greater number of negative charges

27. The Periodic Table • Ionization energy is the energy required to remove an electron from an atom • a low IE means it is easier to remove the electron • Atoms can lose an electron, to form an ion • They do this to achieve noble gas electron configuration (or 8 valence electrons) • When an atom easily loses electrons, it is said to be active • Metals tend to lose electrons

28. The Periodic Table • Ionization Energy Trend: • Down a group – ionization energy decreases • As the valence electrons are farther from the nucleus, the atom gives them up with less energy • Across a period – ionization energy increase • As the number of valence electrons increases in the same energy level, the atom is more resistant to giving up an electron (more energy) • Greatest IE – fluorine • Least IE - francium

30. The Periodic Table • Electronegativity is the measure of the ability of an atom in a chemical compound to attract electrons • All values are based on fluorine • Fluorine is most electronegative atom with a value of 4.0 • The trend decreases in either direction from fluorine