1. Periodicity Trends in the Periodic Table

2. Electron Dot Diagrams Atoms can be represented by electron dot diagrams. The dots on the dot diagram identify only the outside shell electrons.

3. Steps For Drawing Electron Dot Diagrams Write the atomic symbol for the atom. This will represent the nucleus and the inner energy levels. Use a dot to represent an outer shell electron. One dot is placed in each of the four sides before any pairing occurs. Beginning with the fifth dot, pairing may occur to a maximum of eight dots. This is the octet rule.

6. Unpaired electrons are called bonding electrons because they are involved in bond formation. Paired electrons are called lone pairs and are generally not involved in bond formation, however some complex molecules and ions do use the lone pairs to bond.

7. So far you have learned that the elements are arranged in the periodic table according to their atomic number. You have also learned that there is a correlation between the arrangement of elements and their electron configurations. There is also a correlation between how the elements are arranged and their atomic radii, ionization energy, electron affinity, ionic radii, valence electrons, and electronegativity.

8. Explaining Reactivity Period Patterns: – Each element’s period number is the same as the number of shells it’s electrons occupy – 2 nd row has two shells (two energy levels) Group Patterns: – All elements in the main group have the same number of electrons in their outer shell. These are called valence electrons. – Elements in group 2 have two electrons in their outer shell

9. Valence Electrons Chemical compounds form because electrons are lost, gained, or shared between atoms. The electrons that interact in this manner are those electrons in the highest energy levels. The electrons available to be lost, gained, or shared in the formation of chemical compounds are referred to as valence electrons. Atoms with partially filled outer energy levels are unstable. The close the element is to a noble gas, the more reactive/unstable it is

10. Group Number Number of Valence Electrons 11 22 133 144 155 166 177 188 Valence Electrons in Main-Group Elements (any main energy level)

11. Graphing Activity Complete the graphing activity ‘Trends in the periodic table’ –Do questions

12. Atomic Radii The atomic radius is defined as half the distance between the nuclei of identical atoms bonded together. It is measure in picometers.

14. TREND: Atomic Radius Decrease across a period – due to increasing positive nuclear charge & increased # e- on shell (more attraction) Increase down a group – due to increasing number of energy levels (outer electrons are farther from the nucleus = less attraction)

17. Ionization Energies An electron can be removed from an atom if enough energy is supplied. Li + energy  Li + + e - Ionization energy is the energy required to remove one electron from a neutral atom of an element, measured in kilojoules/mole (kj/mol)

18. 1.IE tends to increase across each period. Atoms are getting smaller, electrons are closer to the nucleus 2.IE tends to decrease down a group Atoms are getting larger, electrons are farther from the nucleus Outer electrons becomes increasingly more shielded from the nucleus by inner electrons TREND: Ionization Energy

19. Also 3. Metals have characteristically low IE 4. Nonmentals have high IE 5. Noble gases have a very high IE

21. Group 1 elements have the lowest ionization energies. They lose electrons most easily. Group 17 elements have high ionization energies because they do not lose electrons easily. These elements would rather gain an electron. Going across a period, from left to right, the ionization energy increases.

22. As you go down a group, the electrons are further and further away from the nucleus, and are therefore easier to remove. Going down a group, the ionization energies decrease.

23. Pop Quiz! 1.Rank these elements by decreasing atomic size: a)Se, Br, Ca b)N, Bi, As 2. Which has the greater ionization energy? a) P or Bi? b) P or Ar?

24. Ion: An atom or group of atoms that has a positive (cation) or negative (anion) charge

25. Ionic Radii - Cation A positive ion is known as a cation. – The formation of a cation by loss of one or more electrons always leads to a decrease in radius. The remaining electrons will be pulled in tighter to the nucleus.

26. Ionic Radii - Anion A negative ion is known as an anion. – The formation of an anion by the addition of one or more electrons always leads to an increase in radius. – This is because the total positive charge of the nucleus remains the same when an electron is added so the electrons are not drawn to the nucleus as strongly as they were before the addition. – The electron cloud spreads out because of the greater repulsion between the increased number of electrons.

27. TREND: Ionic Radius Cations have a smaller ionic radius than corresponding atom Protons outnumber electrons Less shielding of electrons Anions have a larger ionic radius than corresponding atom Electrons outnumber protons Greater electron-electron repulsion Ion size increases down a group

30. Electron Affinity Neutral atoms can also acquire electrons. The energy that occurs when an electron is acquired by a neutral atom is called the electron affinity. Most atoms release energy when they acquire an electron. Cl + e -  Cl 1- + energy

31. Sodium on the other hand would rather give up an electron than take one in. If an atom is forced to take an electron, not much energy is released due to the amount of energy used to force the electron in. Na + e - + energy  Na +

32. TREND: Electron Affinity 1.EA tends to increase across each period. 1.EA tends to decrease down a group 1.Halogens have the highest electron affinities 1.Metals have low electron affinities

35. Electronegativity Valence electrons hold atoms together in a chemical compound. In many compounds the negative charge of the valence electrons is concentrated closer to one atom than the other. – This uneven concentration of charge has a significant effect on the chemical properties of a compound. – It is therefore useful to have a measure of how strongly one atom attracts the electrons of another atom within a compound.

36. Electronegativity is a measure of the ability of an atom in a chemical compound to attract electrons.

37. The ability of an atom to attract electrons when bonded 1.EN tends to increase across each period atomic radius is smaller, bonding pair can get closer to the nucleus and be more attracted 2.EN tends to decrease down a group atomic radius is larger, attracts bonding pairs less strongly 3.Nonmetals have high electronegativity 4.Metals have low electronegativity Electronegativity

41. Pop Quiz! 1.Which element has the greatest electron affinity? a)F, Br, Cl b)I, Rb, Sn 2. Which element is most electronegative? a)Po, O, Se b)O, Li, C