1. Periodic Table and Periodic Trends

2. Activity You and partner pull out all the writing utensils you have
Create a method of organizing your writing utensils without talking to other groups

3. I. Development of Modern Periodic Table
A. History

4. Lavoisier
Antoine Lavoisier was first to compile list of known elements in 1790’s

5. Mendeleev
Russian chemist, Mendeleev, organized a table by arranging elements in order of increasing atomic mass

6. Mendeleev is credited with the first periodic table
Mendeleev’s table predicted the existence and properties of undiscovered elements

7. Moseley
English chemist, Moseley, discovered that each element contained a unique number of protons
Moseley arranged elements in a table in order of increasing atomic number

8. Periodic Law:
There is a periodic repetition of chemical and physical properties of the elements when arranged by increasing atomic number

9. B. Modern Periodic Table
118 ?

10. Periodic Table of the Elements: organization of elements
Periodic Table of the Elements: organization of elements. Each square shows the name of an element, its chemical symbol, atomic number, and average atomic mass

12. Periodic Table
elements arranged by increasing atomic number

13. Groups vertical columns
also called ______________ or just families (we’ll call them groups)
chemical families

14. Main Group
Groups 1,2 through are called main group elements (A groups )

15. Transition Metals
Groups 3-12 are called transition elements (B groups )

16. contains elements with similar chemical properties
EX: Li, Na, K

17. Periods horizontal rows
physical and chemical properties change somewhat regularly across a period
elements close to each other in the same period are more similar than those further apart
EX: K, Ca, Sc

18. The Staircase

19. The two sides of the periodic table can be divided into metals and non-metals by the ____________ line
staircase

20. Non-metals are found to the right of the staircase
all elements to the left of the staircase are considered metals (except hydrogen)
elements that border the staircase are called metalloids
Ex. Si, Ge

21. A look into Metals
What state of matter are most of them in at room temperature?
Do any metals look familiar to you?
Solids (except for Mercury – it’s a liquid)
Silver, Gold, Platinum, Lead, Tin

22. Metals at the left of the staircase and bottom two rows
conduct heat and electricity easily
most are solid at room temperature (Hg is a liquid at RT)

23. Metals
exhibit malleability (can be hammered or rolled into thin sheets)
high tensile strength (ability to resist breaking when pulled)

24. Physical Properties of Metals
range from soft (sodium) to extremely hard (platinum)
Physical properties of metals include: malleable, ductile, lustrous, and conductivity of heat and electricity
Malleable: can be hammered into thin sheets
Ductile: can be pulled into wires
Lustrous: shiny appearance

25. Metals – Metalloids – Non-Metals

26. But some are liquids and solids too Oxygen, Hydrogen, Nitrogen
A look into Non-Metals
What state of matter are most of them in at room temperature?
Do any non-metals look familiar to you?
Gases
But some are liquids and solids too
Oxygen, Hydrogen, Nitrogen

27. Non-Metals Can be solids, liquids, or gases
toward right of periodic table
most are gases at RT

28. Non-Metals examples of gases: nitrogen, oxygen, hydrogen
examples of solids: carbon, phosphorus

29. Non-Metals solids are typically brittle
poor conductors of heat and electricity

30. Pure carbon

31. Metalloids
along the stair step line that separates metals and non-metals
have some characteristics of metals and some of non-metals
all are solid at room temperature

32. Metalloids less malleable than metals not as brittle as nonmetals
semiconductors of electricity
used in electronics

33. II. Electrons and the Periodic Table

34. A. Valence electrons
Valence electrons are the ­electrons in the outer-most energy level in an atom
atoms in same group have similar chemical properties because they have the same number of valence electrons

35. Energy Level Diagrams Turn to set II of your Study Packet
# e- = # p+ in an atom
Fills center levels first
Electrons are dots
8 e- max
8 e- max
8 e- max
2 e- max
Nucleus
Turn to set II of your Study Packet

36. Valence Electrons

37. Valence Electrons
the period (row) can indicate the energy level of the element’s valence electrons
the Roman numeral for the main group (A group) elements indicates the number of valence electrons for that group (exception: Helium – 2 valence electrons only)

38. Valence electrons

40. B. Electron blocks
s-block consists of groups 1(IA) and 2 (IIA) and the s orbitals are being filled
p-block consists of groups 13 (IIIA) through 18 (VIIIA) and the p orbitals are being filled
the d-block consists of groups 3 (IIIB) through 12 (IIB) and the d orbitals are being filled

41. Electron Blocks
the f-block includes the Lanthanide series and the Actinide series and the f orbitals are being filled

42. S S 1 2 P 2 3 3 d 4 3 4 5 5 4 6 6 5 7 6 f 4 5
Energy levels

43. III. Properties of Elements

44. A. Hydrogen Group 1/IA and has 1 valence electrons
Only non-metal in group 1
has 3 naturally-occurring isotopes

45. A. Hydrogen
can act like a nonmetal and lose an e- or act like a metal and gain an e-
very reactive

46. B. Alkali Metals link Group 1 or IA 1 valence electrons
forms ions with a charge (cations)
Most reactive metals
only found combined with other elements in nature

47. C. Alkaline Earth Metals
Group IIA or 2 and all have 2 valence electrons
forms ions with a 2+ charge (cations)
less reactive than alkali metals, but still pretty reactive

48. D. Halogens Group VIIA or 17 and have 7 valence e-,
most reactive nonmetals
only found combined with other elements in nature
forms ions with 1- charge (anions)

49. Halogens are commonly referred to as “Halides” (write that down)

50. E. Noble Gases (Inert Gases)
Group VIIIA or 8
8 valence electrons, except He that has 2 valence e-
rarely reactive
most are gases at RT (room temperature)

51. Noble Gases
Examples: helium, neon

52. F. Transition Metals
Groups 3 through 12
fills the d block

53. G. Inner Transition Metals
Also referred to as rare earth metals
The two rows on the bottom of the periodic table
Include Lanthanide and Actinide Series
Fills the f block

54. Keep these 3 factors in mind when considering periodic trends:
IV. Periodic Trends
Keep these 3 factors in mind when considering periodic trends:

55. 1. Nuclear charge
Whenever a proton is added to the nucleus, it creates a stronger “nuclear magnet” pulling in the electrons even more.

56. Electrons added to the same energy level (period) will be pulled in tighter toward the nucleus_.
Ex:

57. 2. Nuclear Shielding
When an energy level is added to the atom (each new period you are adding “layers” between the nucleus and the valence electrons.

58. As energy levels are added, the atom becomes larger and you dilute the pull of the nucleus for the valence electrons because not only are there more “layers” , but the valence electrons are also now farther from the nucleus.

59. (It is easier to remove a valence electron as energy levels are added
Ex.

60. 3. Octet Rule
Atoms will lose, gain or share electrons so they can achieve the electron configuration of the closest noble gas.
As elements get closer to the noble gases on the periodic table (further right ), the greater the attraction for electrons.

61. Noble gases DO NOT attract electrons
Elements on the left side of the periodic table want to lose electrons, so they will not have a great attraction for electrons. EX

62. A. Atomic Radius
Atomic radius is defined as one-half the distance between nuclei of two like atoms in a diatomic molecule
Ex.

63. measured in picometers (pm), 10-12 m or Angstroms (A), 10-10 m
atomic radius indicates relative size of the atom

64. Just the main groups

65. 1. Group trends.
Atomic radius generally increase as you move down a group. This is mainly due to succeeding energy levels being filled.

66. 2. Period trends
Atomic radius generally decrease as you move across a period from left to right. This is mainly due to increasing nuclear charge

67. Atomic Radius Smallest atomic radius = Helium
Largest atomic radius = Francium
Decreases
Increases
increases

68. B. Ionic Radius

69. B Ionic Radius
Ionic radius is the measurement of an ion in a crystal lattice
The units of measurements is picometers (pm) or Angstrom (Å)

70. 1. Group trends for ionic radii
Ionic radius increase as you move down a group. This is because of the added “layers” of electrons.
2. Period Trends for ionic radii
Ionic radius decrease as you move left to right across a period. This is mainly due to the nuclear charge.

71. Ionic Radius
Decreases
Increases
increases

72. C. Ionization Energy
Ionization energy is defined as the amount of energy required to remove an electron from an atom.

74. 1. Group Trends
Ionization energy decrease as you move down a group. This is mainly due to the energy levels.

75. 2. Period trends
Ionization energy increases as you move left to right across a period. This is mainly due to octet

76. Ionization Energy Smallest Ionization NRG =Francium
Largest Ionization NRG = Helium
Increases
Increases
decreases

77. D. Electronegativity
Electronegativity is defined as the tendency for atoms of the element to attract electrons when they are chemically combined with atoms of another element.

78. 1. Group trends
Electronegativity decreases as you move down a group. This is because the distance from the nucleus is increasing

79. 2. Period Trends
Electronegativity increases as you move left to right across a period. This is mainly due to the octet rule.
Note: Noble gases have no electronegativity because they don’t attract electrons at all. Again, think of the octet rule.

80. Electronegativity Most Electronegative = F
Least Electronegative = Fr (+ all noble gases, of course)
Increases
Increases
Decreases

81. increases increases decreases decreases increases increases decreases
Electronegativity
Atomic Radius
Ionizing Energy
Nuclear Charge
Ionic Radius
Shielding
increases
increases
decreases
decreases
increases
increases
decreases
Atomic Radius
Ionic Radius
Ionizing Energy
Electronegativity
Nuclear Charge
Shielding
decreases
increasing
increasing
increasing
decreases