1. The Periodic Table

4. History of the Periodic Table
In 1869, Dmitri Mendeleev arranged the elements according to atomic mass in an attempt to classify them.
He left spaces for elements that he predicted existed, yet hadn’t been discovered yet. This set him apart from others, and he was later proven correct.

6. Henry Moseley
In 1913, Henry Moseley developed the concept of atomic numbers. He determined the frequencies of X-rays emitted as different elements were bombarded with high energy electrons.
Each element produces X-rays of unique frequency, and the frequency increased as atomic mass increased.

7. Atomic Number
Moseley arranged these frequencies in order by assigning a unique whole number, which he called the atomic number.
He correctly identified the atomic number as the number of protons in the nucleus of the atom.
This clarified some problems that Mendeleev had; Ar and K, Te and I.

8. Mendeleev and the Periodic Law
When elements are arranged in order of increasing atomic number, their physical and chemical properties show a pattern or trend.

10. Effective Nuclear Charge
Electrons are (-) charged and nuclei have a (+) charge.
Many properties of the atoms depend on their electron configurations and how strongly the valence electrons are attracted to the nucleus.

11. Effective Nuclear Charge, continued
A valence electron in an atom is attracted to the nucleus, and is repelled by the other electrons in the atom.
The inner electrons (core) are effective in shielding or blocking the nuclear attraction. (Electron shielding effect.)

12. Effective Nuclear Charge = +1.

13. Effective Nuclear Charge, continued
Zeff = Z - S
Z is the number of protons in the nucleus
S is the screening constant
The value of S is usually close to the number of core electrons in the atom.
The effective nuclear charge increases as we move across a period.

14. Periodic Table Arrangement
Periods: Horizontal Rows
Indicates the energy level where the valence electrons are located.
Example: An element in Period 4 has its valence electrons in the 4th energy level.

15. Periodic Table Arrangement
Groups/Families: Vertical columns
Elements in the same column have similar valence electron configurations and therefore similar properties.
Example: All elements in Group I have 1 valence electron.

16. Hydrogen is NOT a part of Group 1. It is a non-metal and a gas.

17. Groups 1 and 2 Group 1 – Alkali Metals Group 2 – Alkali Earth Metals
1 valence electron
VERY reactive
Group 2 – Alkali Earth Metals
2 valence electrons
Reactive, but not as much as Group 1

18. Both Group 1 and Group 2 Reactivity increases as you go down a group
Form cations
Found as compounds only , in nature
Low ionization energy and low electronegativity
Large radii
Radii of positive ions are smaller than the radii of atoms.

19. Radius of Sodium atom versus its ionNa Na
2e-
8e-
8e-
1e-
Sodium ion with 10 electrons
Sodium atom with 11 electrons

20. Groups 3 – 11 Transition Elements
Metallic characteristics
Have multiple (+) ions
Colored compounds

22. Group 17 – Halogens 7 valence electrons
These are the most active of the non-metals!
Gains electrons easily and forms (-) ions
F2 is the most reactive and I2 is the least reactive
F2 is the most electronegative and has the most ionization energy

24. Group 18 – Noble Gases 8 valence electrons
Inactive gases not found in compounds in nature

26. Inner Transition Elements (Bottom Row)
Lanthanides – Lanthanoid series – ones that follow the element Lanthanum (#57)
Rare earth elements that are less than 0.01% in nature
Actinides – Actinoid series – ones that follow Actinum (#89)
All are radioactive and unstable
Those after U are not found in nature
Elements above 92 are synthetic

27. Broad Categories of the Periodic Table

28. Metals The elements on the left side of the table. Malleable Ductile
Lustrous
Conducts heat and electricity
Low ionization energies and low electronegativities
Form cations because they lose electrons
Mostly solids
Reactivity increases as you go down a group.

29. Non-Metals The elements on the right side of the table
Brittle as solids, with no luster
Poor conductors of heat and electricity
High ionization energies and high electronegativities
Reactivities decrease as you go down a group
Form anions because they gain electrons
They can be gases, liquids or solids

31. Semi-Metals (Metalloids)
Some properties are metallic, some are non-metallic and some are in-between
They are found adjacent to the darkened line

32. Other information, continued
Mono-atomic elements – He and all the Noble gases in Group 18
7 Diatomic Elements - BrINClHOF
Allotropes – Two different structures of the same element
Example – Carbon
Graphite diamond buckyballs

34. Buckyball

35. Electronegativity This is the ability to attract electrons
Based on Fluorine, which is 4.0
F has the greatest electronegativity
(Foxy Fluorine steals other elements electrons!)

37. Ionization Energy
The first ionization energy (I1) is the ENERGY needed to remove the first electron in a neutral atom. The second ionization energy (I2) is the energy needed to remove the second electron and so on.
Measured in kJ/mol
The greater the ionization energy the more difficult it is to remove an electron.

38. Trends in the Periodic Table

39. Trends of Atomic Radii (across a period)
The trend increases from LEFT to RIGHT
The radius decreases because the pulling power of the nucleus increases (due to an increase in the effective nuclear charge.)
Atomic # 13
Atomic # 14
Atomic # 11
Atomic # 15 Na Al Si P
Radius = 132
Radius = 128
Radius = 190
Radius = 143

40. Trend in Atomic Radii

41. Trends in Atomic Radii (down a group)
As you go down (↓) a group, the atomic radius increases.
Li 155
Na 190
K 235
Rb 240
Because There is an increase in principal quantum numbers (n).

42. Radii of Neutral Atom versus Its IonNa
2e-
8e-
8e-
1e-
1 e- removed
Sodium ion with 10 electrons (Na +1)
Radius = 120 pm
Sodium atom with 11 electrons (Na 0)
Radius = 190 pm

43. Radii of Neutral Atom versus Its Ion
1 e- added
Fluorine ion with 10 electrons (F -1)
Radius = 95 pm
Fluorine atom with 9 electrons (F 0)
Radius = 57 pm

45. The difference between the non-bonding atomic radius and the bonding atomic radius.
The bonding atomic radius is shorter than the non-bonding radius.

47. Atomic Radii and Bond Lengths
Knowing atomic radii allows us to estimate the bond lengths between different atoms in molecules.
Ex: The Cl ̶ Cl bond length in Cl2 is 1.99 Å, so a radius of 0.99Å is assigned to Cl. In CCl4, the measured length of the C ̶ Cl is 1.77Å, close to the sum of the atomic radii of C and Cl (0.77Å Å).

48. Isoelectronic Series
This is a group of ions all containing the same number of electrons.
The series: O-2, F-, Na+, Mg2+, and Al3+ all have 10 electrons.
Nuclear charges increase as you move through a series, and the atomic radius of the ion decreases.

49. Trends of Electronegativity and Ionization Energy
(across a period)
Electronegativity INCREASES
Ionization Energy INCREASES
The increased effective nuclear charge makes it harder to remove electrons.
WHY?

50. Trends of Electronegativity and Ionization Energy
(down a group)
WHY?
Ionization Energy DECREASES
Electronegativity DECREASES
Electrons are held less tightly because energy levels increase and there is more electron shielding.

53. Electron Affinity
Most atoms gain electrons to form anions. The energy change that occurs when an electron is added to an atom is called electron affinity because it measures the attraction (or affinity) of the atom for the added electron.
For most atoms, energy is released when electrons are added.
Cl(g) + e- → Cl-(g) ΔE = -349 kJ/mol

54. Trends of Metallic Characteristics
Down a group –
INCREASE
in metallic character
Across a period –
DECREASE
In metallic character
(There is a decreased attraction for electrons.)
(There is an increased attraction for electrons.)

56. Metallic Character
Metals have low ionization energies and therefore tend to form positive ions easily.
As a result, metals are oxidized (lose electrons – LEO) when they undergo chemical reactions.
Most metal oxides are basic. Those that dissolve in water react to form metal hydroxides.
CaO (s) + HOH →Ca(OH)2 (aq)