Evolution and Properties of the Periodic Table

1. Chapter 5The Periodic Table

2. History • Russian scientist Dmitri Mendeleev taught chemistry in terms of properties. • Mid 1800 – only atomic masse of elements were known. • Wrote down the elements in order of increasing mass. • Found a pattern of repeating properties.

3. Mendeleev’s Table • Grouped elements in columns by similar properties in order of increasing atomic mass. • Problems - felt that the properties were more important than the mass, so switched order. • Found some gaps - Must be undiscovered elements. • Predicted their properties before they were found.

4. The modern table • Elements are still grouped by properties - similar properties are in the same column. • Order is in increasing atomic number. • Added a column of elements Mendeleev didn’t know about – Noble Gases. • The noble gases weren’t found because they didn’t react with anything.

5. Horizontal rows are called periods • There are 7 periods

6. Vertical columns are called groups or families. • Elements are placed in columns by similar properties.

7. 1A 8A • The elements in the A groups are called the representative elements or main group elements 2A 3A 4A 5A 6A 7A

8. These are called the inner transition elements and they belong here The group B are called the transition elements

9. Group 1A are the alkali metals - most reactive • Group 2A are the alkaline earth metals

10. Group 7A is called the Halogens • Group 8A are the noble gases

11. S- block s1 • Alkali metals all end in s1 • Alkaline earth metals all end in s2 • really have to include He but it fits better later. • Helium has the properties of the noble gases. s2

12. Transition Metals -d block s1 d5 s1 d10 d1 d2 d3 d5 d6 d7 d8 d10

13. The P-block p1 p2 p6 p3 p4 p5

14. f6 f13 f1 f2 f3 f4 f5 f7 f8 f10 f12 f14 f11 f9 F - block • inner transition elements

15. Atomic Size • First problem - where do you start measuring. • The electron cloud doesn’t have a definite edge. • They get around this by measuring more than 1 atom at a time.

16. Atomic Size } • Atomic Radius = half the distance between two nuclei of a diatomic molecule. Radius

17. Trends in Atomic Size • Influenced by two factors. • Energy Level - Higher energy level is further away from nucleus. • Charge on nucleus - More charge (more protons) pulls electrons in closer.

18. Group trends H • As we go down a group • Each atom has another energy level, • So the atoms get bigger. Li Na K Rb

19. Periodic Trends • As you go across a period the radius gets smaller. • Same energy level. • More nuclear charge. • Outermost electrons are pulled closer to nucleus. Na Mg Al Si P S Cl Ar

20. Ionization Energy • The amount of energy required to completely remove an electron from a gaseous atom. • Removing one electron makes a +1 ion. • The energy required is called the first ionization energy.

21. Ionization Energy • The second ionization energy is the energy required to remove the second electron. • Always greater than first IE. • The third IE is the energy required to remove a third electron. • Greater than 1st or 2nd IE.

22. What determines IE • The greater the nuclear charge the greater IE. • The smaller the atom the higher the IE. • Filled and half filled orbitals have lower energy, so achieving them is easier, lower IE. • Shielding

23. Shielding • The electron on the outside energy level has to look through all the other energy levels to see the nucleus. • A second electron has the same shielding.

24. Group trends • As you go down a group first IE decreases because the electron is further away from the nucleus. • More shielding.

25. Periodic trends • All the atoms in the same period have the same energy level. • Same shielding. • Increasing nuclear charge • So IE generally increases from left to right. • Exceptions at full and 1/2 fill orbitals.

26. What Causes the Trends • Full Energy Levels are very low energy. • Noble Gases have full orbitals. • Atoms behave in ways to achieve noble gas configuration.

27. Electron Affinity • The energy change associated with adding an electron to a gaseous atom. • Easiest to add to group 7A - gets them to full energy level. • Increase from left to right atoms become smaller, with greater nuclear charge. • Decrease as we go down a group.

28. Ionic Size • Cations form by losing electrons. • Cations are smaller than the atom they come from. • Metals form cations. • Cations of representative elements have noble gas configuration.

29. Ionic size • Anions form by gaining electrons. • Anions are bigger than the atom they come from. • Nonmetals form anions. • Anions of representative elements have noble gas configuration.

30. Group trends • Adding energy level • Ions get bigger as you go down. Li+1 Na+1 K+1 Rb+1 Cs+1

31. Periodic Trends • Across the period nuclear charge increases so they get smaller. • Energy level changes between anions and cations. N-3 O-2 F-1 B+3 Li+1 C+4 Be+2

32. Electronegativity

33. Electronegativity • The tendency for an atom to attract electrons to itself when it is chemically combined with another element. • How fair it shares - large electronegativity means it pulls the electron toward it. • Atoms with large negative electron affinity have larger electronegativity.

34. Group Trend • The further down a group the farther the electron is away from nucleus and the more electrons an atom has which increases shielding. • More willing to share. • Low electronegativity.

35. Periodic Trend • Metals (left side) - let their electrons go easily • Low electronegativity • Nonmetals (right side) - they want more electrons. • High electronegativity.

36. Ionization energy, electronegativity Electron affinity INCREASE

37. Atomic size increases, shielding constant Ionic size increases

38. Chemical and Physical Properties

39. Descriptive Chemistry • The study of elements and the compounds they form.

40. Hydrogen • Lightest Element • Most Abundant in the Entire Universe • Henry Cavendish – First to systematically collect and study. • Hydrogen – “water-former” because it produces water when combusted in air.

41. Physical Colorless Odorless Tasteless Least dense High speed Diffuse quickly Little attraction Chemical Diatomic If split – highly flammable Acts like 1A and 7A NH3 – ammonia HCl – Hydrochloric acid Metallic hydrides Hydrogen

42. Commercial Sources Byproduct of fuel processes Electrolysis Passing steam over heated coke (impure carbon obtained from coal) Uses Mostly ammonia Rocket fuel Hydrogenated oil

43. Alkali Metals • Most reactive metals. • Most isolated by Sir Humphrey Davy in 1807. • Form cations with +1 charge by losing one electron; 1 valence electron; s-block • Sodium – • most abundant alkali metal • 6th most common element in the earth’s crust. • Never found as pure metals, only combined in nature.

44. Physical Good conductors Bright luster Low densities Soft Easily cut Chemical Very reactive React violently with water to produce bases Stored in oil to prevent reaction with oxygen or water in air Alkali Metals

45. Uses Sodium vapor street lights Na/K pump Baking soda Soap Rayon Paper Salt substitutes Lithium batteries Lithium makes lubricants water resistant and extreme temperature resistant Medication

46. Alkaline Earth Metals • Mostly metal-oxygen compounds which are mostly geologic minerals that dissolve slightly in water. • More dense, harder, and have higher melting points than alkali metals. • Form +2 cations by losing 2 electrons; 2 valence electrons • s-block element

47. Alkaline Earth Metals-Elements • Beryllium – important in minerals; emerald • Magnesium – 8th most common in earth’s crust; important component in chlorophyll • Calcium – 5th most abundant in earth’s crust; major component of bones and teeth • Radium – naturally radioactive and luminesces

48. Physical Dull exterior with shiny interior Harder metal Low density Malleable Good conductors Hard water Chemical Reactivity increases as atom gets larger Uses X-ray tubes Lightweight alloys Flares Plaster Medicine Fireworks Alkaline Earth Metals

49. Transition Metals • d-block elements • Form varying ions both positive and negative. • All have 2 valence electrons

50. Physical Typical metals High density Shiny luster Good conductors Malleable Ductile Solids at room temperature Ex. Mercury – liq. Chemical Nonreactive can resist corrosion Copper and iron corrode quickly Iron, copper, and nickel produce magnetic fields Uses Construction Medical Transition Metals