1. HISTORY OF THE ATOM Subatomic particles and the periodic table

2. Summary of subatomic particles 2

3. Atoms vs. Elements  Blocks analogy 3

4. Elements  Unique number of protons  Atomic number  Arrangement on periodic table  Symbol and name One or two letters First capital, second lowercase 4

5. Chemical Symbols 5 Some symbols are not obvious. Fe for Iron or Cu for copper, based on older names

6. Nucleus  Isotopes  Same element (same number of protons)  Different number of neutrons  Different masses  Chemically identical  Mass number= protons + neutrons 6

7. Chemical Symbols 7

8. Practice Writing Symbols 8  What is boron’s atomic symbol if it has 6 neutrons?  What is leads atomic symbol if it has 124 neutrons?

9. Subatomic Particles Practice 9 Atomic symbol Atomic number Protons NeutronsElectrons Atomic mass 1124 3137 3989 2935 43100

10. Ions  When compounds are formed, electrons are shared or given away and taken  Atom that has lost or gained electrons  Protons remain the same, it’s the electrons that are affecting the charge  Reminder: electrons are negatively charged  Cations: lost electrons, become positive  Anions: gain electrons become negative 10

11. Ions and Compounds 11  Ions behave much differently than the neutral atom.  e.g., The metal sodium, made of neutral Na atoms, is highly reactive and quite unstable. However, the sodium cations, Na +, found in table salt are very nonreactive and stable.  Since materials like table salt are neutral, there must be equal amounts of charge from cations and anions in them.

12. Atomic Structures of Ions 12  Nonmetals form anions.  For each negative charge, the ion has 1 more electron than the neutral atom.  F = 9 p + and 9 e −, F ─ = 9 p + and 10 e −  P = 15 p + and 15 e −, P 3 ─ = 15 p + and 18 e −

13. Atomic Structures of Ions 13  Metals form cations.  For each positive charge, the ion has 1 less electron than the neutral atom.  Na atom = 11 p + and 11 e −, Na + ion = 11 p + and 10 e −  Ca atom = 20 p + and 20 e −, Ca 2+ ion = 20 p + and 18 e −

14. 14 Practice—Complete the Table Atomic NumberProtonsElectrons Ion Charge Ion Symbol 1618 122+ 36 1−1−

15. 15

16. Periodic Law 16  Mendeleev  Ordered elements by atomic number  Repeating pattern of properties  Periodic Law—When the elements are arranged in order of increasing atomic mass, certain sets of properties recur periodically.  put elements with similar properties in the same column

17. Mendeleev’s Periodic Table 17

18. Mendeleev and Periodic Law 18  used pattern to predict properties of undiscovered elements  Where atomic mass order did not fit other properties, he reordered by other properties.  Te & I

19. Periodic Table 19

20. Metals 20  Solids, except Hg  Shiny  Conductive  Heat  Electricity  Malleable  Ductile  Lose electrons, form cations  ~75 % of elements  Lower left on table

21. Nonmetals 21  Solid, liquid and gas  Poor conductors  Heat  Electricity  Solids are brittle  Gain electrons  Anions  Upper right of periodic table, except H

22. Metalloids 22  Characteristics of both metals and nonmetals  Semiconductors

23. Modern Periodic Table 23  Elements with similar chemical and physical properties are in the same column.  Columns are called Groups or Families.  designated by a number and letter at top  Rows are called Periods.  Each period shows the pattern of properties repeated in the next period.

24. Modern Periodic Table 24  Main Group = Representative Elements = “ A ” groups  Transition Elements = “ B ” groups  all metals  Bottom Rows = Inner Transition Elements = Rare Earth Elements  metals  really belong in Periods 6 & 7

25. Modern Periodic Table 25

26. Periodic Table Groups 26

27. Important Groups—Hydrogen 27  nonmetal  colorless, diatomic gas  very low melting point and density  reacts with nonmetals to form molecular compounds  HCl is acidic gas  H 2 O is a liquid

28. Important Groups—Alkali Metals 28  Group 1A = Alkali Metals  hydrogen usually placed here, though it doesn ’ t really belong  soft, low melting points, low density  flame tests  Li = red, Na = yellow, K = violet  very reactive, never find uncombined in nature

29. Important Groups—Alkaline Earth Metals 29  Group 2A = Alkali Earth Metals  harder, higher melting, and denser than alkali metals  Mg alloys used as structural materials  flame tests  Ca = red, Sr = red, Ba = yellow-green  reactive, but less than corresponding alkali metal

30. Important Groups—Halogens 30  Group 7A = Halogens  nonmetals  F 2 and Cl 2 gases; Br 2 liquid; I 2 solid  all diatomic  very reactive

31. Important Groups—Noble Gases 31  Group 8A = Noble Gases  all gases at room temperature  very low melting and boiling points  very unreactive, practically inert  very hard to remove electron from or give an electron to

32. Ion Charge and the Periodic Table 32  The charge on an ion can often be determined from an element ’ s position on the Periodic Table.  Metals always form positively charged cations.  For many main group metals, the charge = the group number.  Nonmetals form negatively charged anions.

33. Ionic Charges and Periodic Table 33

34. Practice—What is the charge on each of the following ions? 34  potassium cation  sulfide anion  calcium cation  bromide anion  aluminum cation