The Periodic Table Unit 5

1. The Periodic Table Unit 5 Objectives: Explain the history of the periodic table Show the organization of the periodic table according to physical and chemical properties Explain what is contained in an element cell on the periodic table Classify elements into groups Explain elemental groupings Show and explain periodic trends

2. Dimitri Mendeleev (1869) • Arranged the 63 known elements into a Periodic Table • Based on atomic mass • Provided for variance from strict atomic weight order • Left space for new elements • Predicted three yet-to-be-discovered elements including silicon(Si)and boron (B) • Did not include any of the Noble Gases

3. Moseley (1887-1915) • Studied under Rutherford • Developed the application of X-ray spectra to study atomic structure • Discovered positioning of elements in the periodic table by atomic numbers • Published results that wavelengths of X-ray emissions coincided with ordering by atomic number • Determined that atomic weight was not the significant player in the periodic law • Found properties of the elements varied periodically with atomic number

4. Cells on the Periodic Table Contains information about an element’s chemical properties • Symbol – tells the name of the element (also see table S) • Atomic Mass – shows the average mass • Atomic Number – gives the number of protons • Electron Configuration – states the relative location of the electrons • Oxidation Number – Shows the possible ways an element can lose or gain electrons

5. The Periodic Cell on the Reference Tables

6. Periods and Periodicity Periods are in horizontal rows • Shows the shells of the valence electrons • Valence shells are the same for that row. Ex: Period 4 – valence electrons are in the 4th shell where n=4 Ex: Period 5 – valence electrons are in the 5th shell where n=5

7. Groups Groups are located in vertical columns and are numbered 1-18 • Shows elements with similar physical and chemical properties • Shows the same number of valence electrons Ex: Group 2 – are all alkaline earth metals and have only 2 electrons in it’s valence shell

8. Classifications • Some Group Classifications • Group 1 – alkali metals • Group 2 – alkaline earth metals • Group 3-16 – transition elements • Group 17 – halogens • Group 18 – noble gases or inert gases

9. Periodic Law The properties of the elements are a periodic function of their atomic numbers When arranged by atomic number, elements show similar properties that repeat periodically

10. Metals, Nonmetals and Metalloids Elements are also classified as metals, nonmetals or metalloids. This diagram shows the area each are located.

11. Metals • Take up more than 2/3 of all the elements • Have unique physical and chemical properties

12. Physical Properties of Metals Some Physical Properties of Metals: • Luster (shininess) • Good conductors of heat and electricity • High density (heavy for their size) • High melting point • Ductile (most metals can be drawn out into thin wires) • Malleable (most metals can be hammered into thin sheets)

13. Chemical Properties of Metals Some Chemical Properties of Metals: • Easily loses electrons • Forms cations (+ charge) • Corrodes easily. Corrosion is a gradual wearing away. Example: silver (Ag) tarnishing and iron (Fe) rusting.

14. Nonmetals Some physical Properties of Nonmetals: • No luster (dull appearance) • Poor conductor of heat and electricity • Brittle (breaks easily) • Not ductile • Not malleable • Low density • Low melting point Some chemical Properties of Nonmetals: • Tend to gain electrons • Nonmetals form anions

15. Metalloids Metalloids are elements that have properties of both metals and nonmetals. Some Physical Properties of Metalloids: • Solids at room temperature • Can be shiny or dull • Ductile • Malleable • Conduct heat and electricity better than nonmetals but not as well as metals

16. Noble Gases • located in group 18 • have no desire to bond with other elements except under extreme conditions • Often called inert gases

17. Periodic Trends • Reactivity • Atomic Radius • Ionic Size • Ionization Energy • Electron Affinity • Electronegativity

18. Reactivity Across a period (row): • Elements on the left lose electrons better than the element to the right of it. • It requires less energy to lose its valence electrons. Down a Column • Elements on the bottom lose electrons better than the element on the top. • Electrons are in shells farther away from the nucleus.

19. Atomic Radius Across a period: • Radii of the atoms decrease left to right • Electrons are pulled closer to the nucleus due to the positive charge Down a Column: • Atomic radii increase from top to bottom • Electrons are farther away from the nuclear charge • There are more electron shells.

20. Ionic Size When an element loses electrons: • Radius gets smaller • More positive charges pulling on the negative electrons When an element gains electrons: • the radius gets larger • More negative charges pulling away from the positive nucleus

21. Ionic Radius

22. Ionization Energy The energy necessary to remove an electron from a neutral atom. • From left to right the ionization energy increases. • From top to bottom the ionization energy decreases • Ionization energy increases for each electron removed

23. Electron Affinity The energy given off when a neutral atom in the gas phase gains an electron to form a negatively charged ion.

24. Electronegativity • An atom’s ability to attract a bonding pair of electrons. • Fluorine (F) is the most electronegative element at 4.0. • Cesium (Cs) and francium(Fr) are the least electronegative at 0.7. (Table S)

25. The Octet Rule Question: How many electrons fit in the valence shell? Answer: • Comes from the electron configuration • Wants a total of eight electrons in the outer shell (valence shell)