1. Exploring the Periodic Table
Modern Chemistry; Holt, Rinehart, & Winston

2. Chapter 5 – Section 1 history of the periodic table
In the late 1800s, scientists had identified over 60 elements. Certain characteristic physical and chemical properties were associated with each element. The physical property called atomic mass provided chemists with a convenient way to organize the elements. At the same time, it was recognized that there were certain elements that had similar chemical properties. Mendeleev arranged the elements in rows according to atomic weight and kept elements with similar chemical properties in the same columns. Today elements are ordered according to atomic number rather than atomic mass.

3. Learning Targets
I can explain the roles of Mendeleev and Moseley in the development of the periodic table.
I can describe the modern periodic table.
I can explain how the periodic law can be used to predict the physical and chemical properties of elements.
I can describe how the elements belonging to a group of the periodic. table are interrelated in terms of atomic number.

4. Stanislao Cannizzaro (1826-1910)
Italian chemist
Determined a method for accurately measuring the relative masses of atoms
His method allowed chemists to search for a relationship between atomic mass and other properties of elements

5. Dmitri Mendeleev (1834-1907) Russian chemist
Credited as being the creator of the first version of the periodic table of elements
Arranged his periodic table according to atomic mass so that elements with similar properties were in the same group
Some elements could not be arranged according to atomic mass in order to keep the elements arranged according to properties
Predicted the properties of elements that had not yet been discovered using his periodic table

6. Mendeleev’s Periodic Table
“I began to look about and write down the elements with their atomic weights and typical properties, analogous elements and like atomic weights on separate cards, and this soon convinced me that the properties of elements are in periodic dependence upon their atomic weights.” --Mendeleev, Principles of Chemistry, 1905, Vol. II

7. Henry Moseley (1887-1915) English chemist Worked with Rutherford
Proved Mendeleev’s arrangement of the periodic table to be correct – only, the periodic table was arranged according to atomic number, not atomic mass

8. The Periodic Law
States that when elements are arranged in order of increasing atomic number, their physical and chemical properties show a periodic pattern

9. Chapter 5 – Section 2 electron configuration and the periodic table
The modern periodic table has 112 squares, which represent a unique element. The distinctive shape of the periodic table comes in part from the periodic law. Elements in the same column have similar properties. These columns are referred to as groups or families of elements. The horizontal rows of the periodic table are called periods. The elements in the periodic table are also grouped as metals, nonmetals, and semimetals. Metals make up most of the periodic table and are located in the center and at the left of the table. With the exception of hydrogen, nonmetals are on the right side, and semimetals are located between the metals and nonmetals. The periodic table can also be viewed in terms of orbital blocks. These orbital blocks refer to the orbitals (s, p, d, and f ) which contain the elements’ incompleted sublevels of electrons.

10. Learning Targets
I can describe the relationship between electrons in sublevels and the length of each period of the periodic table
I can locate and name the four blocks of the periodic table and explain the reasons for these names
I can discuss the relationship between group configurations and group numbers
I can describe the locations in the periodic table and the general properties of the alkali metals, the alkaline-earth metals, the halogens, the transition metals, the noble gases, the actinides, the lanthanides, the metals, the nonmetals, the metalloids, and the main group elements

11. Periodic Law Demonstrated in Groups
Why do elements in groups have similar physical and chemical properties?
They have the same number of valence electrons in their outer energy levels.
Generally, the configurations of the outermost electron shells of elements within the same group are the same.

12. METALS METALLOIDS NONMETALS
In the periodic table below, indicate the location of the groups, periods, alkali metals, alkaline earth metals, halogens, noble gases, lanthanides, actinides, transition metals, inner transition metals, main group elements, metals, nonmetals and metalloids.
METALS
METALLOIDS
NONMETALS

13. In the periodic table below, indicate the location of the groups, periods, alkali metals, alkaline earth metals, halogens, noble gases, lanthanides, actinides, transition metals, inner transition metals, main group elements, metals, nonmetals and metalloids.
ALKALI METALS

14. ALKALINE-EARTH METALS
In the periodic table below, indicate the location of the groups, periods, alkali metals, alkaline earth metals, halogens, noble gases, lanthanides, actinides, transition metals, inner transition metals, main group elements, metals, nonmetals and metalloids.
ALKALINE-EARTH METALS

15. In the periodic table below, indicate the location of the groups, periods, alkali metals, alkaline earth metals, halogens, noble gases, lanthanides, actinides, transition metals, inner transition metals, main group elements, metals, nonmetals and metalloids.
HALOGENS

16. In the periodic table below, indicate the location of the groups, periods, alkali metals, alkaline earth metals, halogens, noble gases, lanthanides, actinides, transition metals, inner transition metals, main group elements, metals, nonmetals and metalloids.
NOBLE GASES

17. In the periodic table below, indicate the location of the groups, periods, alkali metals, alkaline earth metals, halogens, noble gases, lanthanides, actinides, transition metals, inner transition metals, main group elements, metals, nonmetals and metalloids.
TRANSITION METALS

18. INNER TRANSITION (Rare Earth) METALS
In the periodic table below, indicate the location of the groups, periods, alkali metals, alkaline earth metals, halogens, noble gases, lanthanides, actinides, transition metals, inner transition metals, main group elements, metals, nonmetals and metalloids.
INNER TRANSITION (Rare Earth) METALS

19. In the periodic table below, indicate the location of the groups, periods, alkali metals, alkaline earth metals, halogens, noble gases, lanthanides, actinides, transition metals, inner transition metals, main group elements, metals, nonmetals and metalloids.
LANTHANIDES

20. In the periodic table below, indicate the location of the groups, periods, alkali metals, alkaline earth metals, halogens, noble gases, lanthanides, actinides, transition metals, inner transition metals, main group elements, metals, nonmetals and metalloids.
ACTINIDES

21. In the periodic table below, indicate the location of the groups, periods, alkali metals, alkaline earth metals, halogens, noble gases, lanthanides, actinides, transition metals, inner transition metals, main group elements, metals, nonmetals and metalloids.
PERIODS

22. In the periodic table below, indicate the location of the groups, periods, alkali metals, alkaline earth metals, halogens, noble gases, lanthanides, actinides, transition metals, inner transition metals, main group elements, metals, nonmetals and metalloids.
GROUPS

23. In the periodic table below, indicate the location of the groups, periods, alkali metals, alkaline earth metals, halogens, noble gases, lanthanides, actinides, transition metals, inner transition metals, main group elements, metals, nonmetals and metalloids.
MAIN GROUP ELEMENTS

24. Let’s Compare! Metals Nonmetals Metalloids
Good conductors of heat and electricity
Malleable
Ductile
Luster
Typically solids at room temperature
Solids, liquids and gases at room temperature
Solids are brittle and dull
Poor conductors of heat and electricity
Have properties of both metals and nonmetals
Mostly brittle solids
Intermediate conductors of electricity- AKA semiconductors

25. Properties of Alkali Metals
Extremely reactive
Readily react with water and air
Silvery in appearance
Soft enough to cut with a knife
Lower densities than other metals
Lower melting points than other metals

26. Properties of Alkaline-Earth Metals
Harder & stronger than alkali metals
Higher densities & melting points than alkali metals
Less reactive than alkali metals

27. Properties of Halogens
Most reactive nonmetals
React readily with most metals to form salts
Most electronegative elements

28. Properties of Noble Gases
Least reactive elements because their highest occupied energy levels are completely filled with an octet of electrons (except He, which only requires 2 electrons to be filled).

29. Properties of Transition Metals
High densities
High melting points
Good conductors of heat & electricity
High luster
Less reactive than alkali and alkaline-earth metals

30. Properties of p Block Metals
Harder and more dense than the s block metals
Softer and less dense than the d block metals.

31. Properties of Lanthanides
Soft, silvery metals
Similar reactivity to alkaline-earth metals

32. Properties of Actinides
All radioactive
The first 4 have been found naturally on Earth

33. Did you know?
Oxygen, carbon, hydrogen and nitrogen make up 96% of the human body mass
Calcium and phosphorous make up 3%
Sodium, potassium, chloride and magnesium make up 0.7%
Iron, cobalt, copper, zinc, selenium, cyanide and fluorine are found in trace amounts
All of these elements are essential for life!

34. Chapter 5 – Section 3 electron configurations and periodic properties
Many of the properties of the elements change in predictable ways as you move across a period or move down a group of the periodic table. The predictable changes in these properties are called periodic trends. There are periodic trends for properties such as atomic radius, ionic size, ionization energy, electron affinity, and electronegativity. Knowledge of these trends helps develop a better understanding of the periodic table and of the patterns of behavior of the elements.

35. Learning Targets I can define the term periodic trend.
I can define atomic radius, ionic radius, ionization energy, electron affinity and electronegativity.
I can describe the general trends on the periodic table for atomic radius, ionic radius, electron affinity, ionization energy and electronegativity.
I can apply the trends on the periodic table to answer questions regarding size, electron affinity, ionization energy and electronegativity.

36. Distance between nuclei
Atomic Radii
Atomic radius – one-half the distance between the nuclei of identical atoms that are bonded together
Atomic Radius
Distance between nuclei

37. Period Trends Decreases across a period
Higher effective nuclear charge – more protons being added to the nucleus at the same time electrons are being added to the same energy level

38. Why?
Protons are added to the nucleus moving across a period from left to right
This increases the charge of the nucleus (effective nuclear charge – Zeff)
As Zeff increases, the electrons are pulled closer to the nucleus

39. Period Trends +
+ +
+ + +
+ +

40. Group Trends Increase down a group
Although more protons are being added to the nucleus going down a period, energy levels are also being added. The inner energy levels create a shielding effect from the attractive nuclear forces.

41. Why?
The addition of shells increases the electrons’ distance from the nucleus and the size of the atom
Electron-electron repulsion “plumps” up the atom
Zeff decreases the further the electrons are from the nucleus
n=3
n=2
n=1

42. Variations in Atomic Radii

43. Atomic Radii Trends
DECREASES
DECREASES

44. Ionization Energy
The energy required to remove one electron from a neutral atom of an element creating an ion
A + Energy  A+ + e-

45. Period Trends Increase across a period Why?
Zeff increases across the period

46. Group Trends Decrease down the group Why?
Electron shielding causes a decrease in effective nuclear charge
Electron-electron repulsion forces increase
Electrons have higher PE as the energy level increases
Shielding effect – outer electrons are shielded from the nuclear charge attractive forces by the electrons in the inner energy shells

47. Variations in Ionization Energies
Draw the orbital notation for Group 5A and Group 6A.
Can you explain the dips in the chart for these 2 groups?

48. Variations in Ionization Energies
If removing an electron will create an empty or ½ filled subshell, ionization energy will decrease.

49. Successive Ionization Energies
Each successive electron removed from an ion feels an increasingly stronger effective nuclear charge (Zeff) – therefore, successive ionization energies are larger than 1st ionization energies
A large jump in ionization energy occurs when removing an electron from an ion that assumes a noble gas configuration

50. Ionization Energy Trends
INCREASES
INCREASES

51. Electron Affinity
The change in energy that a neutral atom undergoes when an electron is acquired (the ability to attract an e -)
A + e-  A- + energy
[negative energy value (exothermic)]
A + e- + energy  A-
[positive energy value (endothermic)]

52. Period Trends Increase across a period Why?
Zeff increases across the period

53. Group Trends Decrease down the group Why?
Electron shielding causes a decrease in effective nuclear charge
Electron-electron repulsion forces increase
Electrons have higher PE as the energy level increases
Shielding effect – outer electrons are shielded from the nuclear charge attractive forces by the electrons in the inner energy shells

54. Variations in Electron Affinities
Why is there such a large decrease in energy for groups 2A and 5A?

55. Electron Affinity Trends
INCREASES
INCREASES

56. Ionic Radii Cation – positively charged ion
Cations are smaller than their parent atom – why?
Anion – negatively charged ion
Anions are bigger than their parent atom – why?
Removal of an electron creates an unbalanced positive charge increasing Zeff and decreasing the radius of the ion.
Addition of an electron creates an unbalanced negative charge decreasing Zeff and increasing the radius of the ion. +

57. Ionic Radii Trends
DECREASES
DECREASES

58. Valence Electrons
Electrons available to be gained, lost or shared in the formation of a chemical compound
Located in the outer energy level

59. Electronegativity
A measure of the ability of an atom in a chemical compound to attract a bonding pair of electrons
NOTE *Electronegativity is a property of atoms in compounds and thus differs from ionization energy and electron affinity, which are properties of isolated atoms*

60. Trends Increase across a period Decrease down a group
Effective nuclear charge increases
Decrease down a group
Increase in atomic size and increase in electron shielding decreases the effective nuclear charge
Electronegativity depends upon:
The number of protons in the nucleus
The distance from the nucleus
Electron shielding

61. Electronegativity Trends
INCREASES
INCREASES