1. UNIT 6: PERIODIC TABLE

2. Aim # 1 How has the Periodic Table evolved overtime?
John Newlands, an English scientist listed all the known elements by their atomic masses.
Elements with the least
amount of mass were first
followed by elements with
greater masses
1864

3. 1869
Dimitri Mendeleev, Russian scientist organized elements according to their atomic mass and properties
Developed a chart with columns and rows
Columns called family or group

4. Families contained elements with similar properties
Each row is called period
Left to right across a
period elements increase
in atomic mass
Mendeleev believed there were more elements that would be added and left places in his chart empty

5. MODERN PERIODIC TABLE Lists elements according to their atomic number
Still used columns called families which consist of elements with similar properties
Still uses rows called periods . Elements increase by 1 moving from left to right.
Periodic table lists 118 element.
May need to be revised as new elements are discovered.

6. Periodic Table Geography

7. The horizontal rows of the periodic table are called PERIODS.

8. The elements in any group of the periodic table have similar physical and chemical properties!
The vertical columns of the periodic table are called GROUPS, or FAMILIES.

9. Periodic Law
When elements are arranged in order of increasing atomic number, there is a periodic pattern in their physical and chemical properties.

10. AIM# 2: What are the properties and locations of metals and nonmetals?
Ductile: able to be pulled into a wire
Malleable: able to be hammered into shapes
Brittle: will shatter into pieces if struck
Luster: shiny
Dull: not shiny

11. METALS Location: to the left of the staircase
State(s) the metals exist as at STP: all solids except Hg – liquid
Properties:
Malleable - low ionization energy
Ductile - low electronegativity
Has luster - lose electrons to form positive
Good conductors ions with smaller radii
High MP, BP
MELPS HELP!!!!!!!!!!!

12. NONMETALS Location: to the right of the staircase
State(s) the metals exist as at STP: all solids except H, He, N, O, F, Cl, Group 18- Gases; Br - liquid
Properties:
Brittle - high ionization energy
Dull - high electronegativity
Low MP, BP gain electrons to form neg
Poor conductors ions with larger radii

13. PROPERTIES OF TRANSITION METALS
Transition metals are generally located in groups 3 to 12. They form colored ions. When mixed with water they produce colored aqueous solutions.
Multiple oxidation states

14. METALLOIDS
Metalloids are located on the staircase between metals and nonmetals. They display some metallic properties and some nonmetallic properties. There are six metalloids
List them below and make sure to put their name and symbol
→ B, Si, Ge, As, Sb, Te

15. NOBLE GASES
Noble gases are located in group 18. They have a full valence shell and are therefore very stable. They are not very reactive/ inert

16. AIM# 3: Are there patterns (trends) within the Periodic Table
AIM# 3: Are there patterns (trends) within the Periodic Table? – IONIZATION ENERGY
Amount of energy needed to remove the most loosely bound electron from a neutral gaseous atom. (Listed in Table S)

17. PERIOD TREND
Ex) List the ionization energies for the following elements in period 2
What do you notice is happening to the ionization energies when looking across the period from left to right? increasing
Values from left to right across a period increase
Element Li B F
Ionization Energy

18. GROUP TREND
Ex) List the ionization energies for the following elements in group 2
What do you notice is happening to the ionization energies when looking down a group from top to bottom? decreasing
In a group ionization energy decrease.
Element
Ionization Energy Be Sr Ba

19. ATOMIC RADIUS
Atomic radius is distance the distance between nuclei of identical atoms bonded together. (Basically tells the size of a neutral atom. Listed in Table S)

20. PERIOD TREND
Ex) List the atomic radius for the following elements in period 2
What do you notice is happening to the atomic radius when looking across the period from left to right? decreasing
Values from left to right across a period decrease
Element Li B F
Atomic Radius

21. GROUP TREND
Ex) List the atomic radius for the following elements in group 2
What do you notice is happening to the atomic radius when looking down a group from top to bottom? increasing
In a group atomic radius increase.
Element
Atomic Radius Be Sr Ba

22. ELECTRONEGATIVITY
The electronegativity value of an atom is a measure of its attraction for electrons when bonded to another atom. Electronegativity values are listed in Table S.

23. PERIOD TREND
Ex) List the electronegativity values for the following elements in period 2
What do you notice is happening to the electronegativity when looking across the period from left to right? increasing
Values from left to right across a period increase
Element Li B F
Electronegativity

24. GROUP TREND
Ex) List the electronegativity values for the following elements in group 2
What do you notice is happening to the electronegativity when looking down a group from top to bottom? decreasing
In a group atomic radius decrease.
Element
Atomic Radius Be Sr Ra

25. VALENCE ELECTRONS
Valence electrons are electrons in the outer energy level of an atom

26. PERIOD TREND
Using your Periodic Table as a reference, list the # of valence electrons for period 2 elements in the order below)
What is happening to the number of valence electrons going from left to right across a period? Increases
Going from left to right across a period the number of valence electrons increase.
Element Li Be B C N O F Ne
Valence e-

27. GROUP TREND
Using your Periodic Table as a reference, list the # of valence electrons for group 2 elements in the order below)
What is happening to the number of valence electrons as you move down a group from top to bottom? Stay the same
From top to bottom in a group the number of valence electrons remains the same.
Element
Atomic Radius Be Mg Ca Sr Ba Ra

28. PRINCIPLE ENERGY LEVELS
The orbitals in an atom form a series of energy levels in which electrons may be found. Each electron in an atom has its own distinct amount of energy that corresponds to the energy level that it occupies. Electrons can gain or lose energy and move to different levels.

29. PERIOD TREND
Using your Periodic Table as a reference, list the # of principle energy levels for period 2 elements in the order below)
What is happening to the number of principle energy levels going from left to right across a period? nothing
 Going from left to right across a period the number of principle energy levels remains the same.
Element Li Be B C N O F Ne
Energy level

30. GROUP TREND
Using your Periodic Table as a reference, list the # of principle energy levels for group 2 elements in the order below)
What is happening to the number of principle energy levels as you move down a group from top to bottom? Increasing
From top to bottom in a group the number of principle energy levels increases.
Element
PEL Be Mg Ca Sr Ba Ra

31. METALLIC CHARACTER
Most known elements are metals. Most active metals are located in groups 1 and 2. In any group on the Periodic Table the metallic properties of the elements increase from top to bottom.
Most active metals are in group 1 and 2
Fr is the most reactive

32. IONIC RADIUS
Atoms gain or lose electrons to become charged particles called ions. Metals tend to lose their valence electrons in chemical reactions to become positive ions. Nonmetals tend to gain electrons in chemical reactions and become negative ions. As these atoms gain or lose electrons, they complete an octet of valence electrons. Ionic radius is the distance from the nucleus to the outer principle energy level of the ion.

33. IONIC RADIUS
When a metal loses its valence electrons and becomes positively charged it loses an energy level, its radius decreases. When a nonmetal gains an electron and becomes negatively charged its radius increases.
**The ionic radius of metals is smaller than the radius of the original atom. The radii of nonmetallic ions are larger than those of the original atoms** .

34. Aim # 4 What are the effects of atomic structure on Periodic Trends?
Magnet Demonstration
Atomic Radius
Electronegativity
Ionization energy

35. Effect of atomic structure on Periodic Trends
Remember that all of the elements of the periodic table are made up of the simplest building blocks of matter called atoms
The properties of the elements follow general patterns called trends.

36. Effect of atomic structure on Periodic Trends
When we observe the properties across a period or down a group, these trends can be explained by changes in atomic structure from one element to the next.

37. Effect of atomic structure on Periodic Trends
The periodic trends of atomic radius, electronegativity, and ionization energy can all be explained by changes in nuclear charge, and/ or the number of P.E.L.

38. Effect of atomic structure on Periodic Trends
[Think about magnets and attraction! Nuclear charge represents the strength of a magnet's attractive force and the number of principle energy levels represents the distance between the magnets.]

39. ATOMIC RADIUS

40. ATOMIC RADIUS: Period Trend
As we look at the elements across period the nuclear charge in each successive atom increases. However, the number of principle energy levels remains the same
Therefore, with each successive element the second shell is pulled closer to the nucleus. This causes the atomic radius to decrease

41. ATOMIC RADIUS: Period Trend
CONCLUSION: Increased nuclear charge (# of protons) across the period causes the atomic radius to decrease

42. ATOMIC RADIUS: Group Trend
As we go down group 1 the number of principle energy levels increases. Therefore, each successive atom down a group increases in size. ( atomic radius)

43. ATOMIC RADIUS: Group Trend
CONCLUSION: Increased number of P.E.L. causes the atomic radius to increase down a group

44. ELECTRONEGATIVITY
The measure of the relative attraction of an atom for a foreign electron.

45. ELECTRONEGATIVITY

46. ELECTRONEGATIVITY: Period Trend
As we go across a period from left to right the nuclear charge increases.
However, the number of principle energy levels remains the same.
Therefore, elements to the right of the periodic table have an increased attraction for foreign electrons.( electronegativity)

47. ELECTRONEGATIVITY: Period Trend
CONCLUSION: Increased nuclear charge across a period causes an increase in electronegativity

48. ELECTRONEGATIVITY: Group Trend
As we go down a group the number of P.E.L. increases
Therefore, there is an increased shielding effect and a weaker attraction by the nucleus for foreign electrons
Therefore, each successive atom down a group decreases in electronegativity

49. ELECTRONEGATIVITY: Group Trend
CONCLUSION: Increased number of P.E.L. Causes an increased shielding effect, which decreased electronegativity

50. IONIZATION ENERGY
The amount of energy required to remove the outermost electron from an atom.

51. IONIZATION ENERGY

52. IONIZATION ENERGY: Period Trend
As we go across a period from left to right the nuclear charge increases.
However, the number of principle energy levels remains the same.

53. IONIZATION ENERGY: Period Trend
Therefore, elements to the right of the periodic table hold onto their own valence electrons more strongly.
This means it requires a larger amount of ionization energy exerted by a foreign atom to remove

54. IONIZATION ENERGY: Period Trend
CONCLUSION: An increase in nuclear charge across a period causes the ionization energy to increase

55. IONIZATION ENERGY: Group Trend
As we go down a group the number of P.E.L. increases
Therefore, there is an increased shielding effect and a weaker attraction for its own valence electrons
Therefore, it requires less ionization energy to remove a valence electron from these atoms as we go down a group

56. IONIZATION ENERGY: GROUP Trend
CONCLUSION: An increased shielding effect due to an increasing number of P.E.L. causes a decrease in ionization energy down a group