Chapter 5 Notes

1. Periodic Law Chapter 5 Notes

2. During the 1800’s, chemists had discovered more than halfof all of the elements and were forced with the task of rememberingall of their properties and what happened when they were combined with other elements into compounds. There was no table of elementsdeveloped to indicate what these elements behaved like. Section 1: History of The periodic Table

3. Also, the method for determiningthe atomic masses of atoms was not standardized. Therefore, several different chemists were reportingdifferent atomic masses for the same element. This caused it to be very difficult for any scientist to replicate any other scientist’s work. In September of 1860, a group of chemists met to discuss these problems. Section 1: History of The periodic Table

4. Italian chemist StanislaoCannizzaro presented a method that allowed the masses to be determined very accurately. This paved the way for standardvalues for atomic mass and initiated a search for relationships between atomic mass and other properties of the elements. Section 1: History of The periodic Table

5. After Russian chemist Dmitri Mendeleevheard about the new atomic masses discussed at the meeting in Germany, he decided to include them in a chemistry textbook that he was working on. He had hoped to organize the elements according to their properties. Mendeleev and Chemical Periodicity

6. He placed the names of each element on a card along with the atomic masses and any other properties that the element had. Then he arranged the cards according to various propertiesand looked for any trends. Mendeleev and Chemical Periodicity

7. Mendeleev noticed that when the elements were arranged in order of increasing atomic mass, certain similaritiesin their chemical properties appeared at regular intervals. A repeating pattern is referred to as periodic. Mendeleev and Chemical Periodicity

8. Mendeleev then created a table in which elements with similar propertieswere grouped together. The first periodic table of the elements had been made. In Mendeleev’s first table, he placed iodine after tellurium even though it was thought tellurium had a larger atomic mass. He did this because iodine’s properties fit better with the other elements in the group. Mendeleev and Chemical Periodicity

9. He was correctin this assumption. He also left several emptyspaces and predicted the existence of other elements based on his table. The properties of these elements are very close to what Mendeleev predicted they would be. Mendeleev and Chemical Periodicity

10. Most chemists were persuaded to accept Mendeleev’s periodic table. However, two questions still remained. Why could most of the elements be arranged in the order of increasing atomic mass? What was the reason for chemical periodicity? Mendeleev and Chemical Periodicity

11. The first question was not answered until more than 40years after Mendeleev’s periodic table was published. In 1911, English scientist Henry Moseley examined the spectra of 38 different metals. He discovered a previously unrecognizedpattern. The atoms actually fit better when they were arranged according to an increasing number of protons. Mosley and the Periodic Law

12. Moseley’s work led to the modern definition of atomic number and the recognition that atomic number, not atomic mass, should be the basis for the organizationof the periodic table. Mosley and the Periodic Law

13. This is why Mendeleev noticed that iodine fit better in that group than tellurium. Iodine had a smalleratomic mass, but a largeratomic number and therefore should be after tellurium. Mendeleev ‘s principle of chemical periodicity is stated in the periodic law. Mosley and the Periodic Law

14. Periodic Law– The physical and chemical properties of the elements are periodic functions of their atomic numbers. Mosley and the Periodic Law

15. The periodic table has undergone extensivechange since Mendeleev’s time. Chemists have discovered new elements and have synthesizedmany others. However, each of them can be placed in a group of other elements with similar properties. The modern periodic Table

16. Periodic Table– An arrangement of the elements in order of their atomic numbers so that elements with similar properties fall in the same column, or group The modern periodic Table

17. Perhaps the most significant addition to the periodic table came with the discovery of the noblegases. John William Strutt and Sir William Ramsay discovered argon in 1894. It had previously not been discovered because of its lack of reactivity. The noble gases

18. In 1868, another noble gas, helium was discovered as a component of the sun. Ramsay showed that helium also existed on Earth. In order to fit argon and helium onto the periodic table, Ramsay proposed a new group. He placed the group after group 17. The noble gases

19. The next major step in the development of the periodic table was the discovery of the of the lanthanideelements in the early 1900’s. Because these elements are so similarin chemical and physical properties, the process of separating and identifying them was very tedious. The lanthanides

20. Another major development came with the discovery of the actinideseries. To save space these two groups are usually placed at the bottom of the periodic table. The actinides

21. Periodicity with respect to atomic numbercan be seen in any group. Think about the noble gas group. Helium is the first element in that family. As we increase in atomic number, we do not see another element that behaves as helium does until we get to neon. periodicity

22. This is what periodicity is talking about. The reason that periodicity exists is because of the number of electronsthat surround the nucleus. Each of the members of a family have the same number of electrons in their outershell, which makes the atoms of that group behave in similar manners. periodicity

23. The periodic table gives you a ton of information about the natureand characteristicsof the elements if you are able to remember the explanationsthat you can use. There are only 4possible explanations of all of the periodic trends. Section 3:electron configuration and periodic properties

24. Effective nuclear charge, Zeff—essentially equal to the group number. Think of the atoms in the first column as having a Zeff of one while the halogens have a Zeff of 7! The idea is that the higher the Zeff, the more positive the nucleus, the more attractive force there is emanating from the nucleus drawing electrons in or holding them in place. Section 3:electron configuration and periodic properties

25. Distance—attractive forces dissipate with increased distance. Distant electrons are held loosely and thus easily removed. Section 3:electron configuration and periodic properties

26. Shielding—electrons in the inside of the atom effectively shield the nucleus’ attractive force for the valence electrons. Use this only when going up and down the table, not across. There is ineffective shielding within a sublevel or energy level. Section 3:electron configuration and periodic properties

27. Minimize electron/electron repulsions—this puts the atom at a lower energy state and makes it more stable. • Now, we will go through all of the periodic trends and use these explanations to discuss why each trend works the way it works as you go across the periodic table and down the periodic table. Section 3:electron configuration and periodic properties

28. Ideally, the size of an atom is defined by the edge of its orbital. However, this boundary is fuzzyand varies under different conditions. Therefore, to estimatethe size of an atom, the conditions under which the atom exists must be specified. Atomic Radii

29. One way to express an atom’s radius is to measure the distance between the nucleiof two identical atoms that are chemically bonded together, then divide the distance by two. • Atomic radius– one-half the distance between the nuclei of identical atoms that are bonded together Atomic radii

30. So, as you go across the periodic table from left to right, the atomic radius gets smaller. Which, if you think about it, doesn’t seem to make sense. As you move that direction there are more protons, more neutrons, and more electronsin each of the atoms. Therefore, it makes sense that the atom should be getting larger. Atomic radii

31. However, that is not one of our arguments. Let’s think about Zeff. As we are going across the table, the positive charge in each subsequent nucleus is getting larger. Therefore, the Zeff is becoming larger. The electrons, which are negativelycharged are more attracted to a more positive nucleus. That is why the atomic radius gets smalleras you go across the periodic table. Atomic radii

32. Now, let’s go down the periodic table. As you go down the table, the atomic radius becomes larger. The reason that the atom grows larger is that you have added another principalenergy level. From last chapter, we talked about electrons in higher energy levels being fartheraway from the nucleus. Atomic radii

33. If the electrons are farther away from the nucleus, then Zeff can’t act on them as effectively and the atom becomes larger. Atomic radii

34. An electron can be removed from an atom if enough energyis supplied. Using A as a symbol for an atom of any element, the process can be expressed as follows. A + energy  A+ + e- Ionization energy

35. The A+ represents an ion of element A with a single positive charge, referred to as a 1+ ion. • Ion– an atom or group of bonded atoms that has a positive or negative charge • Ionization– any process that results in the formation of an ion Ionization energy

36. To compare the ease with which atoms of different elements give up electrons, chemists compare ionization energies. • Ionization energy– the energy required to remove one electron from a neutral atom of an element Ionization energy

37. As you go across the periodic table ionization energy increases. As Zeff increases across the table, the electrons are going to be held tighter and tighter to the nucleus. It will become more difficultto pull an electron, which is negative, from the nucleus, which is positive. Ionization energy

38. You could also argue that because the atomic radius decreasesas you go across the periodic table, the electrons are thereby closer to the nucleus, thus increasingthe amount of energy needed to pull an electron off. Ionization energy

39. As you go down the periodic table you will see that ionization energy becomes smaller. This comes from the fact that the atoms are larger, and the electron that would be taken is a valence electron. Ionization energy

40. Those electrons that are farther away are easierto pull off because the nucleus is not able to attract them as much because of the distance. Ionization energy

41. Neutral atoms can also acquire electrons. • Electron affinity– the energy change that occurs when an electron is acquired by a neutral atom • Most atoms releaseenergy when they acquire an electron. However, some requireenergy to accept an electron. When you see the graph in your book, those will have an electron affinity of zero. Electron affinity

42. Electron affinity increasesas you go across the table. Increasing Zeff causes the nucleus to attract the electron more. The more the electrons are pulled on by the nucleus, the moreenergy will be given off when the atom gains the electron. However, it should be noted that electron/electron repulsionplays a heavy role in the anomalies that occur in this trend. Electron affinity

43. Electron affinity decreasesdown a group because the electron is being added to atoms that are getting larger. The increased distance from the nucleus means that the nucleus will not be able to pull as much on the electron, thereby reducingthe energy that is given off when the electron joins the atom. Electron affinity

44. Cation– a positive ion • Anion– a negative ion • The formation of a cation by the lossof one of more electrons always leads to a decreasein atomic radius because the removal of the highest energy level electrons results in a smallerelectron cloud. Also, the remaining electrons are drawn closerto the nucleus by its unbalanced positive charge. Ionic Radii

45. The formation of an anion by the additionof one or more electrons will always lead to an increasein atomic radius. This is because the total positive charge of the nucleus remains the samewhen an electron is added to an atom or an ion. Ionic Radii

46. So the electrons are not drawn to the nucleus as strongly as they were before the addition of the extra electron. The electron cloud also spreadsout because of greater repulsionbetween the increased amount of electrons. Ionic Radii

47. When discussing ionic radii, you need to only compare cations to cations and anions to anions. As you move across a period, the ions should get smallerbecause Zeff is increasing. With the nuclear charge being larger, the nucleus pulls on the electrons more effectively, thereby shrinkingthe ion. Ionic Radii

48. As you go down a group, the ions increasebecause you have added an entire new principal energy level. Ionic Radii

49. Valenceelectrons hold atoms together in chemical compounds. In many compounds, the negative charge of the valence electrons is concentrated closerto one atom than to another. This uneven concentration of charge has a significant effect on the chemical propertiesof a compound. electronegativity

50. It is therefore useful to have a measure of how strongly one atom attracts the electrons of another atom within a compound. • Electronegativity– a measure of the ability of an atom in a chemical compound to attract electrons from another atom in the compound electronegativity