1. The Periodic Table

2. Organizing Principle Chemists used properties of elements to sort them into groups

3. Mendeleev Properties of elements differ in orderly way - elements can be grouped Dmitri Mendeleev (Russian Chemist) – discovered periodicity

4. Mendeleev’s Periodic Table Increase in atomic mass from left to right - this is how Mendeleev arranged the elements Columns of elements with similar properties Mendeleev’s chart had gaps - undiscovered elements

5. Modern Periodic Table Henry Mosely removed irregularities in Mendeleev’s Periodic views - placed elements according to increasing atomic numbers (in modern periodic table) - caused elements to be grouped easier according to similarities

6. Periodic Law Periodic Law – when elements arranged in order of increasing atomic number, there is periodic repetition of their physical and chemical properties Periodic Table – page 162-163

7. Metals, Nonmetals, Metalloids Three classes of elements are: - metals - nonmetals - metalloids

8. Metals Metals – one class of elements that are good conductors of heat and electric current - tend to be ductile (draw into wire), malleable (pound into thin sheets), and shiny - make up about 80% of elements - Example: Copper (Cu) Page 158 shows metals (yellow section)

9. Nonmetals Nonmetals – element that tends to be a poor conductor of heat and electric current - generally have properties opposite to those of metals - most are gases at room temperature - Example: Oxygen (O) Page 158 (blue section)

10. Metalloid Metalloid – an element that tends to have properties that are similar to both metals and nonmetals Example: Silicon (Si) Page 158 (green section)

11. Metals, Nonmetals, Metalloids Copper Oxygen Silicon

12. Squares in the Periodic Table Periodic Table displays symbols and names of the elements - also displays information about structure of atoms

13. Group 1 (1 A) – Alkali Metals Alkali Metals – soft, highly reactive metals - shiny surfaces dull as react with Oxygen - stored under oil or kerosene - produce alkaline solutions - metallic properties Examples: Lithium, Sodium, Potassium

14. Alkali Metals

15. Group 2 (2A) – Alkaline Earth Metals Alkaline-Earth Metals – harder, denser, stronger and have higher melting points than alkali metals - less reactive than Group 1 - more difficult to lose 2 electrons than 1 Examples: Magnesium and Calcium

16. Group 2 – Alkaline-Earth Metals

17. Halogens Halogens – Group 17 (7 A) elements - combine easy with metals to form salts - derived from Latin = salt-former - most reactive nonmetals - Example: Chlorine

18. Chlorine

19. Electron Configurations in Groups Elements can be sorted into the following based on their electron configurations: - noble gases - representative elements - transition metals - inner transition elements

20. Noble Gases Noble Gases – Group 18 (8A) elements - sometimes called inert gases – rarely take part in a reaction - full s and p orbitals in highest principle energy levels - electron configurations very stable – resistant to change - Examples: Helium, Neon, Argon

21. Noble Gases

22. Representative Elements Representative Elements – display a wide range of physical and chemical properties - some metals, some non-metals, and some metalloids - most are solids at room temp - some gases at room temp - one liquid (bromine) at room temp - page 165

23. Transition Elements Transition Metal – one of Group B elements in highest occupied s sublevel and nearby d sublevel generally containing electrons (page 162-163) Inner Transition Metal – element in lanthanide or actinide series - highest occupied s sublevel and nearby f sublevel generally contain electrons

24. Groups 3-12 (1B-8B)

25. Lanthanides and Actinides Lanthanides – shiny, metallic elements - atomic numbers 58-71 - fill 4f orbitals - some used to create phosphor dots in television tubes

26. Lanthanides

27. Lanthanides and Actinides Actinides – metallic elements - atomic numbers 90 – 103 - have unstable arrangements of protons and neutrons - all have radioactive forms - Example: Uranium

28. Actinides

29. Hydrogen Hydrogen in chemical family by itself Most common use in ammonia

30. Blocks of Elements Periodic table can be broken down into blocks (page 166) – based on electron configurations and positions of elements s-block = Groups 1A, 2A and Hydrogen p-block = Groups 3A – 8A d-block = 1B – 8B f-block = lanthanides and actinides

31. Atomic Radius Increases within a family Atomic Radius – ½ of the distance between nuclei of two atoms of the same element when atoms are joined - usually measured in picometers In general, atomic size increases from top to bottom within a group - decreases from left to right across a period

32. Atomic Radius

33. Shielding Effect Shielding effect – reduction of the attractive force between a nucleus and its outer electrons - due to blocking effect of inner electrons - allows outer electrons to be farther away from nucleus

34. Shielding Effect

35. Atomic Size Atomic radii generally decrease as you move across a period from left to right - each atom gains one electron and one proton - as proton numbers increase = positive charge increases creating a greater pull on electrons - decrease in atomic radius

36. Atomic Size

37. Ions Ion – atom or group of atoms that has gained or lost one or more electrons to acquire a net electric charge Positive and negative ions form when electrons are transferred between atoms Cation – ion with positive charge Anion – ion with negative charge

38. Ionization Energy Ionization Energy – amount of energy needed to remove electron from a specific atom or ion in ground state in the gas phase - generally decreases as move down a group of elements - increases across a period from left to right Takes different amounts of ionization energy for every electron taken off

39. Ionization Energy

40. Electron Affinity Electron Affinity – energy change that accompanies addition of electron to an atom in gas phase - become more negative from left to right - top to bottom within a group, electron affinity tends to become less negative

41. Electron Affinity

42. Trends in Ionic Size Cations are always smaller than the atoms from which they form Anions are always larger than the atoms from which they form

43. Trends in Ionic Size

44. Electronegativity Electronegativity – ability of an atom to attract electrons to itself when it is combined with another atom - decrease going down a group - increase going left to right across a period

45. Electronegativity

46. How are elements created? Of 110 known elements, only elements up to number 92 occur naturally - naturally occurring elements come from forged matter and energy deep in space - remaining elements are synthetic (man made)

47. Elements

48. Natural Elements Hydrogen atoms and Helium atoms fuse to form different elements (mostly in stars) - when star dies and explodes – releases elements into space

49. Synthetic Elements First man made element – created in 1919 Nitrogen + Helium = Oxygen + Hydrogen Used alpha particles = helium nuclei - targeted toward Nitrogen to form Oxygen Elements heavier than Uranium are synthetic Today, scientists bombard nuclei with protons, neutrons, alpha, and beta particles to create elements