1. TRENDS in the PERIODIC TABLE A trend is a pattern or a repetition of particular properties.

2. The periodic table is arranged in a certain way to keep elements with similar properties close together

3. Demetri Mendeelev Developed the first real periodic table which ours is based upon. He was able to predict the missing elements and their properties once the table was put together.

4. History of the Periodic Table 1871 – Mendeleev arranged the elements according to: 1. Increasing atomic mass 2. Elements w/ similar properties were put in the same row 1913 – Moseley arranged the elements according to: 1. Increasing atomic number 2. Elements w/ similar properties were put in the same column

5. GROUPS vs. PERIODS Groups go up and down. Periods go left and right. Groups share many similarities. Periods show periodically (regularly) changing properties.

6. THERE ARE SEVERAL BASIC TRENDS (or patterns) THAT WE NEED TO RECOGNIZE AND UNDERSTAND.

7. Periodic Trends Periodic Trends – patterns (don’t always hold true) can be seen with our current arrangement of the elements (Moseley) Trends we’ll be looking at: 1.Atomic Radius 2.Ionization Energy 3. Electronegativity (or Electron Affinity)

8. Atomic Radius Atomic Radius – size of an atom (distance from nucleus to outermost e - )

9. Atomic Radius Trend Group Trend – As you go down a column, atomic radius increases As you go down, e - are filled into orbitals that are farther away from the nucleus (attraction not as strong) Periodic Trend – As you go across a period (L to R), atomic radius decreases As you go L to R, e - are put into the same orbital, but more p + and e - total (more attraction = smaller size)

10. SHOWS Atomic Sizes for Groups and Periods

11. Atomic Radius The effect is that the more positive nucleus has a greater pull on the electron cloud. The nucleus is more positive and the electron cloud is more negative. The increased attraction pulls the cloud in, making atoms smaller as we move from left to right across a period.

12. Effective Nuclear Charge What keeps electrons from simply flying off into space? Effective nuclear charge is the pull that an electron “feels” from the nucleus. The closer an electron is to the nucleus, the more pull it feels. As effective nuclear charge increases, the electron cloud is pulled in tighter.

13. Atomic Radius On your help sheet, draw arrows like this:

14. Shielding As more energy levels are added to atoms, the inner layers of electrons shield the outer electrons from the nucleus. The effective nuclear charge on those outer electrons is less, and so the outer electrons are less tightly held.

15. Ionization Energy Ionization Energy – energy needed to remove outermost e -

16. Ionization Energy Group Trend – As you go down a column, ionization energy decreases As you go down, atomic size is increasing (less attraction), so easier to remove an e - Periodic Trend – As you go across a period (L to R), ionization energy increases As you go L to R, atomic size is decreasing (more attraction), so more difficult to remove an e - (also, metals want to lose e -, but nonmetals do not)

17. Ionization Energy This is the second important periodic trend. If an electron is given enough energy (in the form of a photon) to overcome the effective nuclear charge holding the electron in the cloud, it can leave the atom completely. The atom has been “ionized” or charged. The number of protons and electrons is no longer equal.

18. Ionization Energy The energy required to remove an electron from an atom is ionization energy. (measured in kilojoules, kJ) The larger the atom is, the easier its electrons are to remove. Ionization energy and atomic radius are inversely proportional. Ionization energy is always endothermic, that is energy is added to the atom to remove the electron.

19. Ionization Energy (Potential) Draw arrows on your help sheet like this:

20. Electronegativity Electronegativity- tendency of an atom to attract e -

21. Electronegativity Trend Group Trend – As you go down a column, electronegativity decreases As you go down, atomic size is increasing, so less attraction to its own e - and other atom’s e - Periodic Trend – As you go across a period (L to R), electronegativity increases As you go L to R, atomic size is decreasing, so there is more attraction to its own e - and other atom’s e -

22. Electron Affinity What does the word ‘affinity’ mean? Electron affinity is the energy change that occurs when an atom gains an electron (also measured in kJ). Where ionization energy is always endothermic, electron affinity is usually exothermic, but not always.

23. Electron Affinity Electron affinity is exothermic if there is an empty or partially empty orbital for an electron to occupy. If there are no empty spaces, a new orbital or PEL must be created, making the process endothermic. This is true for the alkaline earth metals and the noble gases.

24. Electron Affinity Your help sheet should look like this: ++

25. Reactivity Reactivity – tendency of an atom to react Metals – lose e - when they react, so metals’ reactivity is based on lowest Ionization Energy (bottom/left corner) Low I.E = High Reactivity Nonmetals – gain e - when they react, so nonmetals’ reactivity is based on high electronegativity (upper/right corner) High electronegativity = High reactivity

26. Overall Reactivity This ties all the previous trends together in one package. However, we must treat metals and nonmetals separately. The most reactive metals are the largest since they are the best electron givers. The most reactive nonmetals are the smallest ones, the best electron takers.

27. Overall Reactivity Your help sheet will look like this: 0

28. The Octet Rule The “goal” of most atoms (except H, Li and Be) is to have an octet or group of 8 electrons in their valence energy level. They may accomplish this by either giving electrons away or taking them. Metals generally give electrons, nonmetals take them from other atoms. Atoms that have gained or lost electrons are called ions.

29. Ions When an atom gains an electron, it becomes negatively charged (more electrons than protons ) and is called an anion. In the same way that nonmetal atoms can gain electrons, metal atoms can lose electrons. They become positively charged cations.

30. Ions Here is a simple way to remember which is the cation and which the anion: This is a cat-ion. This is Ann Ion. He’s a “plussy” cat! She’s unhappy and negative. +

31. Cation Formation 11p+ Na atom 1 valence electron Valence e- lost in ion formation Effective nuclear charge on remaining electrons increases. Remaining e- are pulled in closer to the nucleus. Ionic size decreases. Result: a smaller sodium cation, Na +

32. Anion Formation 17p+ Chlorine atom with 7 valence e- One e- is added to the outer shell. Effective nuclear charge is reduced and the e- cloud expands. A chloride ion is produced. It is larger than the original atom.

33. Summary of Periodic Trends