1. The Periodic Table

2. Organizing the Elements Early chemists used the properties of elements to sort them into groups. As early as 1829 there was a published classification system for the known elements - they were grouped into “triads” with similar properties

3. The First Periodic Table In 1869 Dmitri Mendeleev published a table of the known elements. There were 60 known elements at the time. Mendeleev wrote the properties of the elements on cards, which allowed him to move them around until he found a pattern that worked. He organized the elements into groups based on a set of repeating properties, and arranged the elements in order of increasing atomic mass. (They did not know about protons back then!)

4. Mendeleev’s Table

5. The Current Periodic Table Mendeleev wasn’t too far off, but he had to break the pattern of increasing atomic mass to keep elements with similar properties in the same columns. In 1913 Henry Mosley put the elements in rows by increasing ATOMIC NUMBER!!

6. Names & Symbols Sources of the element names: – Most common source for name is some property of the element – Element’s place of discovery – Mineral in which element was found – To honor person or place – Symbols derived from English/Latin name Element symbols have only the first letter capitalized

7. Positions of Key Elements/Groups The horizontal rows are called periods numbered from 1 to 7. The vertical columns are called groups are labeled from 1 to 18.

8. Three Classes of Elements Metals – high luster good conductors most are ductile/malleable Nonmetals – Non-lusterous poor conductors most are gases at room temp. Metalloids – along the staircase; properties of both

9. Families Periods Columns of elements are called groups or families. Elements in each family have similar but not identical properties. For example, lithium (Li), sodium (Na), potassium (K), and other members of family IA are all soft, white, shiny metals. All elements in a family have the same number of valence electrons. Each horizontal row of elements is called a period. The elements in a period are not alike in properties. In fact, the properties change greatly across even given row. The first element in a period is always an extremely active solid. The last element in a period, is always an inactive gas.

10. Common Groups Representatives – Group A (the “tall” columns) These tend to be the most reactive group of elements, with very predictable properties. We also number the representative columns as columns 1A through 8A. Active Metals Group 1 (1A) – Alkali Metals Group 2 (2A) – Alkaline Earth Metals Nonmetals Group 17(7A) – Halogens Group 18(8A) – Noble Gases

11. Transitions – Group B These are metals with less predictable properties than the representative group Inner Transition Elements (Actinide and Lanthinides) Similar to the transition metals, because this group actually fits within the transition metals

13. “Long” view of the periodic table

14. Hydrogen Hydrogen belongs to a family of its own. Although it looks like it is in the alkali metal group, it is NOT a metal! Hydrogen is a reactive gas. Hydrogen was involved in the explosion of the Hindenberg. Hydrogen is promising as an alternative fuel source for automobiles

15. Helium Although its also a “lighter than air” gas like hydrogen, helium has very different properties It is a noble gas, with a filled outer electron shell. Unlike hydrogen, it is very UNreactive.

16. Alkali Metals 1 st column on the periodic table (Group 1) not including hydrogen. Very reactive metals, always combined with something else in nature (like in salt). In pure (elemental) form they are soft enough to cut with a butter knife

17. Alkaline Earth Metals Second column on the periodic table. (Group 2) Reactive metals that are always combined with nonmetals in nature. Several of these elements are important mineral nutrients (such as Mg and Ca)

18. Transition Metals Elements in groups 3- 12 Less reactive, harder metals Includes metals used in jewelry and construction. Metals used “as metal.”

19. Halogens Elements in group 17 Very reactive, volatile, diatomic, nonmetals Always found combined with other element in nature. Used as disinfectants and to strengthen teeth.

20. The Noble Gases Elements in this group are all gases at room temperature. Called “noble” because they do not readily interact with other elements. They are inert. They do not react with other elements because their outermost energy level (valence shell) is already full.

21. Using the Periodic Table to Determine Electron Configuration The period number of an element signifies the highest energy level (first quantum number) an electron in that element occupies The representative group number of an element (1A – 8A) signifies the number of valence electrons in the highest energy level. Valence electrons are the electrons in the outermost energy level of an atom.

22. Periodic Table Electron configuration An element’s position is related to its atomic # and its electron configuration

24. What is the outer electron configuration of: Phosphorus Calcium Bromine

25. Periodic Trends A trend is the tendency of a set of data to move in a certain direction You can predict the properties of an element based on its position in the periodic table and the known periodic trends For example, atomic radius (size of each atom) increases as you go down the periodic table, but decreases as you go across it. How do we explain these periodic trends?

26. The Shielding Effect The shielding effect helps explains periodic trends. In atoms with many electrons, each electron feels: the pull of the nucleus, as well as the repulsion from the other electrons. Repulsion from other electrons in lower energy levels blocks or “shields” the outer electrons from the full attractive force of the nucleus.

27. 1. Atomic Radius Atomic radius (size) decreases across a period – the increased number of protons in the nucleus pull the electrons in the outer energy level closer. Atomic radius increases as you go down a group – As you progress to the next period, the electrons are in a higher energy level, “shielded” from the pull of the nucleus by the lower energy levels

28. 2. Ionization Energy 1 st ionization energy – energy needed to remove the first electron from an atom (while in a gaseous state) Increases across row/period (the increased number of protons in the nucleus increases the pull on the electrons, so it’s harder to pull an electron off) Decreases down a column/group due to the shielding effect (lower energy levels blocking the pull of the nucleus) Metals have low ionization energies Nonmetals have high ionization energies

29. 3. Electronegativity Tendency for atoms to attract electrons when combined w/ another element Follows the same trend as ionization energy: increases across a row, decreases down a group

30. Summary of Periodic Trends

31. Valence Electrons Valence electrons are the electrons in the highest occupied energy level of an atom. These are the electrons that determine the reactivity of an element. The number of valence electrons in the atom of an element is equal to its representative group number. Examples: oxygen (group 6A = 6 valence e - ) calcium (group 2A = 2 valence e - )

32. What does it mean to be reactive? We described elements according to their reactivity. Elements that are reactive bond easily with other elements to make compounds. Some elements are only found in nature bonded with other elements.

33. What makes an element reactive? An incomplete valence electron level. All atoms (except the smallest ones) want to have 8 electrons in their very outermost energy level. This is the most stable configuration. Atoms bond until this level is completely full. Atoms with few valence electrons lose them during bonding. Atoms with 5, 6, or 7 valence electrons gain electrons during bonding.

34. Electron Dot Diagram Valence Electrons are presented as dots on the diagrams called “Lewis Dot Structures”. Each single electron is placed before a pair, just like in orbital diagrams.

35. Electron Dot Diagram

36. Octet Rule Noble gases (except He) have the most stable, “perfect” configuration with 8 valence electrons in outer level. As a result, they are inert (unreactive in chemical reactions). OCTET RULE: Atoms of all other elements tend to achieve the electron configuration of Noble gases by either losing or gaining or sharing electrons to have a complete octet of electrons in outer level. 1s 2 2s 2 2p 6

39. Metals: Metals achieve octet by forming cations: losing their valence electrons and leaving a compete octet on a previous (lower) energy level. Example: 11 Na 1s 2 2s 2 2p 6 3s 1 11 Na + 1s 2 2s 2 2p 6

40. Nonmetals: Nonmetals achieve an octet by forming anions: gaining electrons to complete an octet in the outermost energy level Example: 16 S 1s 2 2s 2 2p 6 3s 2 3p 4 16 S 2- 1s 2 2s 2 2p 6 3s 2 3p 6