# 1 Trends on the Periodic Table Chapter 7 Written by JoAnne L. Swanson University of Central Florida.

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1 Trends on the Periodic Table Chapter 7 Written by JoAnne L. Swanson University of Central Florida

2 Effective Nuclear Charge Z eff = Z – S = (The protons in the nucleus) – (the number of electrons between the nucleus and electron of interest) Z eff allows us to estimate trends in properties ionization energy atomic and ionic size electron affinity

3 Electrons in an atom are attracted to the positive pull of the nucleus while at the same time they are __________________ ________________________________________________-. Outer electrons are also ________________________________ ________________________________________________ These attractions and repulsions are taken into consideration when interpreting trends on the periodic table.

4 IONIZATION ENERGY Ionization energy is the energy required to remove the most loosely held electron from an atom in its gaseous state. Example to demonstrate: Al(g)  Al +1 (g) + e - 1 st E i = 580 KJ / mol Al +1 (g)  Al +2 (g) + e - 2 nd E i = 1815 KJ / mol Al +2 (g)  Al +3 (g) + e - 3 rd E i = 2740 KJ / mol Al +3 (g)  Al +4 (g) + e - 4 th E i = 11600 KJ / mol

5 The outermost electron is always ____________________. This is because it is ______________________________________ ______________________. In this case the 3p electron is removed. The second ionization energy is considerably higher because now the electron (the 3s) is being removed from ___________________. ______________________________________________. The fourth ionization energy is the highest for Al +3 because this would be removing a “____” electron. That is, _______________ ______________________________________________________ ______________________________________________________.

6 In the graph of ionization energies, there is a general trend of _________________________________________________. There are deviations from this trend due to shielding, Z eff, and repulsions. Generally, ionization energy increases going to the right of a period because the electrons are in the same PEL, and are not shielded from the nuclear charge. That is, the positive protons are increasing across the period along with the negative electrons, in the same energy level. There is a stronger positive to negative attraction as both protons and electrons increase. Therefore it would take more energy to remove an electron.

7 The ionization energy _________________________. Here the electron is in a higher energy level, farther from the nucleus, and thus held less tightly. Deviations from the trend – Be has a higher ionization energy than B the 2p electron in B is shielded from the nuclear charge by the 2s electrons and is thus easier to remove. N has a higher ionization energy than O N has 3 unpaired electrons. O has a paired electron in the 2p x orbital and thus an increase in repulsion. The electron is easier to remove due to the repulsion.

8 ELECTRON AFFINITY Electron affinity (E a ) is _______________________________ ____________________________________. The more negative the number, the more ____________________. Therefore, atoms with a more negative electron affinity tend to _______________. Repulsions and nuclear charge affect electron affinity – N - (g) is unstable and so has a positive E a (endothermic). Here an electron would be added to an already occupied ‘p’ orbital in which the repulsion is greater than the Z eff (the positive attraction of the nucleus).

9 O(g) + e -  O - (g) is exothermic. Here the difference in Z eff (it has more protons therefore more attraction to electrons), overcomes the repulsion of adding a 2 nd electron to an orbital. Oddly enough, O 2- is not stable and a 2 nd electron cannot be added to O -. The reason we see so many compounds of O 2- has to do with the O 2- being stabilized by the large attraction of the positive charge of the ion it is bonding to.

10 So the basic trend for Electron affinity with some deviations is – a more negative E a going to the right in a period. Going down a group, E a becomes more positive with some deviations. The electrons are being added at increasing distances from the nucleus and also to increasingly larger orbitals. Z eff and repulsion play a smaller role. It is helpful to look at the basic trends, deviations aside, as depending on metallic or non metallic character. Metallic character increases down and left. Non metallic character increases up and right. Metals tend to lose electrons therefore, have low ionization energies and more positive electron affinities. Non metals tend to gain electrons therefore have high ionization energies and more negative, exothermic, electron affinities.

11 ATOMIC AND IONIC SIZE Atomic size ______________going from left to right in a period. The Z eff _______________ as p + and e - are added. Shielding decreases in the same period. Atomic size ___________ going down a group. Electrons are being added to a larger orbital, farther away from the nucleus. There is less effective nuclear charge as the distance from the nucleus increases, therefore less attraction between the electron and the nucleus.

12 Ionic Size – Cations are ____________than their respective neutral atoms due to less repulsions and the emptying of the largest, outer orbital. Anions on the other hand are larger than their neutral atoms due to repulsions between the added electrons while no additional protons are added. This will cause the electrons to spread out more.

13 Properties within Groups The representative elements (s and p filling elements) exhibit trends that change in a regular way 1. The number and type of valence electrons primarily determine an atoms chemistry. 2. The electron configuration is one of the most valuable pieces of information for being able to predict behavior of elements. 3. The most basic division of elements is that of the metals and nonmetals.

14 Metals with more metallic character (lower left), are the most reactive metals. They have low ionization energies and so lose electrons more easily. The non metals with more nonmetallic character (upper right, excluding inert gases), are the most reactive non metals and tend to gain electrons. They have high ionization energies and more exothermic E a (more negative).

15 METALS - Metals tend to form ionic compounds with non metals, that is,_____________________________________________. These are held together by strong ________________ (+ and – attractions). Metals in water form H 2 gas and base.(Since metals have low E i they form bases that will ionize in water. Na + H 2 O  NaOH + H 2 (g) Metal oxides in water are basic. Na 2 O + H 2 O  2NaOH NiO + H 2 O  Ni(OH) 2 THEREFORE SINCE METAL OXIDES ARE BASIC, WHEN PUT IN ACID THEY FORM SALT AND WATER: Na 2 O + 2 HCl  2NaCl + H 2 O

16 THE Lithium REACTS WITH OXYGEN TO FORM A METAL OXIDE AND ALSO REACTS WITH HYDROGEN TO FORM A METAL HYDRIDE: 4Li + O 2  2Li 2 O 2Li + H 2  2LiH But other alkali metals form peroxides (O 2 2- ) with O 2 : Na + O 2  Na 2 O 2 What is a hydride???

17 The alkaline earth metals get more reactive going down a group: Be + H 2 O(l)  NR Be + H 2 O(g)  NR Mg + H 2 O(l)  NR Mg + H 2 O(g)  Mg(OH) 2 Ca + H 2 O(l)  Ca(OH) 2 Sr + H 2 O(l)  Sr(OH) 2 THEY ALSO FORM METAL OXIDES: Mg + O 2  MgO

18 NON METAL OXIDES ARE _________, THAT IS, WHEN REACTED WITH WATER PRODUCE ______. IT FOLLOWS THAT NON METAL OXIDES WITH BASE SHOULD THEN PRODUCE _____________________. P 4 O 6 (s) + 6 H 2 O  4 H 3 PO 3 P 4 O 6 (s) + 12 KOH(aq)  4 K 3 PO 3 + 6 HOH

19 PHYSICAL CHARACTERISTICS OF METALS: LUSTEROUS AND SOFT HIGH MELTING POINTS GOOD CONDUCTORS OF ELECTRICITY (WHY?) AND HEAT MALLEABLE AND DUCTILE

20 NON METALS - THEY ARE SOLIDS LIQUIDS AND GASES AT ROOM TEMP. SOLIDS ARE BRITTLE AND DULL LOWER MELTING POINTS THAN METALS POOR CONDUCTORS OF ELECTRICITY WHEN NON METALS BOND TO OTHER NONMETALS THEY GENERALLY FORM __________________ AND THEREFORE ______________ SUBSTANCES. IN THESE COVALENT BONDS, ELECTRONS ARE _________ BETWEEN ATOMS BY OVERLAP OF ORBITALS.

21 HALOGENS (X) REACT WITH HYDROGEN TO FORM HX(g) HALOGENS REACT WITH METALS TO FORM MX(S). F AND Cl ARE THE MOST REACTIVE HALOGENS. FLUORINE IS THE MOST REACTIVE HALOGEN AND WILL TAKE ELECTRONS FROM ALMOST ANY SUBSTANCE. F 2 + H 2 O  HF + O 2 EXOTHERMIC F 2 + SiO 2  SiF 4 + O 2 EXOTHERMIC THUS, FLUORINE GAS IS VERY DANGEROUS TO WORK WITH. SULFUR AND OXYGEN HAVE SIMILAR PROPERTIES AND FORM SULFIDES AND OXIDES WITH OTHER NONMETALS. OXYGEN BEING MORE REACTIVE DOES SO MORE EXOTHERMICALLY.

22 SUMMARY OF CHAPTER: 1.Z EFF, REPULSIONS, AND SHEILDING 2.IONIZATION, ELECTRON AFFINITY, ATOMIC AND IONIC SIZE TRENDS AND DEVIATIONS FROM TRENDS 3.SIMILARITIES IN GROUPS METALS AND THEIR REACTIONS WITH WATER, ACID, OXYGEN, HYDROGEN AND OTHER NON METALS NON METALS AND THEIR REACTIONS WITH OXYGEN AND WATER BASE, AND OTHER NON METALS 4. CHARACTERISTICS OF METALS AND NON METALS

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