1. Periodic Trends

2. Periodic Trends
In the Periodic Table, chemical and physical properties appear “periodically”
So, what is a trend?
A trend is the general direction in which something tends to happen.

3. Periodic Trends
The elements on the Periodic Table of Elements show many trends in their physical and chemical properties.
Across the rows (periods or series)
Down the columns (groups or families)

4. Atomic Radius
½ the distance between the nuclei of 2 like atoms in a diatomic molecule
The atoms of the 8 main groups are shown here.

5. Atomic Radius - Groups
Atomic radius increases as you move down a group
Why?
More electrons in more Principal Energy Levels
More Electron shells
Atomic size increases

6. Atomic Radius - Periods
Atomic radius decreases as you move across a period
Why?
(-) electrons increase, but so do (+) protons !!!
Increased (+) nuclear charge pulls the (-) electrons closer to the nucleus
Atomic size decreases

7. Atomic Radius - Periods
The size trend in periods is less pronounced than in groups because of the electron shielding effect.

8. Shielding Effect
Reduction in effective nuclear charge on an electron that is caused by the repulsive forces of other electrons between it and the nucleus
In an atom with one electron, that electron experiences the full charge of the positive nucleus. However, in an atom with many electrons, the outer electrons are simultaneously attracted to the positive nucleus and repelled by the negatively charged electrons.

9. Atomic Radius - Graph

10. Atomic Radius

11. Quick Check Hold up 1 or 2 fingers, Which atom is bigger? K or 2. Li
1. Al or Cl

12. Ionic Radius What are ions? Ions are charged atoms, either + or -
Cations are positive ions
Cations form when atoms lose electrons
Anions are negative ions
Anions form when atoms gain electrons

13. The Octet Rule – Atoms will tend to gain, lose or share electrons so that they have a set of 8 Valence electrons.
Cation
Anion

14. What kind of Ion will it make?
Calcium
Oxygen

15. Ionic Radius

16. Ionic Radius – Cations Group
Cations are smaller than their parent atoms.
Why?
By losing their valence electrons, they lose their entire valence shell
Cations are formed by the metals on the left side of the Periodic Table

17. Ionic Radius – Cations Groups
Ionic size increases as you move down a group for the same reason atomic size increases
Number of principal energy levels increases

18. Ionic Radius – Anions Groups
Anions are larger than their parent ions
Why?
When extra (-) electrons are added, extra (+) protons are NOT added to the nucleus
Effective nuclear attraction is less for the increases number of electrons

19. Ionic Radius – Anions Group
Ionic size increases as you move down a group for the same reason atomic size increases
Number of principal energy levels increases
Anions are formed by nonmetals on the right side of the Periodic Table

20. Ionic Radius - Periods Just like their parent atoms…
Cations get smaller as you move from left to right
Anions get smaller as you move from left to right
Increased (+) nuclear charge pulls the (-) electrons closer to the nucleus

21. Ionic Radius

22. Quick Check
Which one is bigger?
1. Li+ or Na+
1. Cl- or Ar

23. Ionization Energy Energy is needed to remove an electron from an atom
The energy needed to overcome the attraction of the nuclear charge and remove an electron (from a gaseous atom) is called the Ionization Energy

24. 1st Ionization Energy
The energy needed to remove the 1st electron from an atom is the 1st Ionization Energy

25. Factors Affecting Ionization Energy
Atomic Radius
Smaller atoms hang on to valence electrons more tightly, and so have higher ionization energy

26. Factors Affecting Ionization Energy
Charge
The higher the positive charge becomes, the harder it is to pull away additional electrons
Second ionization energy is always higher than the first

27. Factors Affecting Ionization Energy
Orbital Type
It's easier to remove electrons from p orbitals than from s orbitals, which are “deeper”

28. Factors Affecting Ionization Energy
Electron Pairing
Within a subshell, paired electrons are easier to remove than unpaired ones
Reason: repulsion between electrons in the same orbital is higher than repulsion between electrons in different orbitals

29. Factors Affecting Ionization Energy
Electron Pairing – Example
On the basis of gross periodic trends, one might expect O to have a higher ionization energy than N. However, the ionization energy of N is 1402 kJ/mol and the ionization energy of O is only 1314 kJ/mol.
Taking away an electron from O is much easier, because the O contains a paired electron in its valence shell which is repelled by its partner.

30. 1st Ionization Energy

31. 2nd, 3rd, 4th Ionization Energies, etc.
Subsequent electrons require more energy to remove than the first electron
How much more energy is needed depends on what energy levels and orbitals the electrons are in

32. 2nd, 3rd, 4th Ionization Energies, etc.
Source:

33. Electronegativity
Ability of an atom to attract electrons toward itself in a chemical bond

34. Electronegativity
The difference between the electronegativities of two atoms will determine what kind of bond they form
Linus Pauling used an element's ionization energy and electron affinity to predict how it will behave in a bond.
The more energy it takes to pull off the outer electron of an atom, the less likely it is to allow another atom to take those electrons. The more energy the atom releases when it gains an electron, the more likely it is to take electrons from another atom in bonding. These two energies were used to compute a numerical score.

35. Electronegativity - Periods
Electronegativity increases going left to right across the periodic table.
Fluorine's high nuclear charge coupled with its small size make it hold onto bonding electrons more tightly than any other element. Lithium has a lower nuclear charge and is actually larger than fluorine. Its valence electron is not tightly held and it tends to surrender it in chemical bonds. Li Be B C N O F
1.0
1.5
2.0
2.5
3.0
3.5
4.0

36. Electronegativity - Groups
Electronegativity decreases going down a group
The bonding electrons are increasingly distant from the attraction of the nucleus H Li Na K Rb Cs Fr

37. Electronegativity

38. Electronegativity

39. Image Sources http://www.Chem4kids.com http://images.encarta.msn.com

40. Credits
PowerPoint: Adela J. Dziekanowski