1. Periodic Law

2.  During the 1800’s, chemists had discovered more than half of all of the elements and were forced with the task of remembering all of their properties and what happened when they were combined with other elements into compounds. There was no table of elements developed to indicate what these elements behaved like.

3.  Also, the method for determining the atomic masses of atoms was not standardized. Therefore, several different chemists were reporting different 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.

4.  Italian chemist Stanislao Cannizzaro presented a method that allowed the masses to be determined very accurately. This paved the way for standard values for atomic mass and initiated a search for relationships between atomic mass and other properties of the elements.

5.  After Russian chemist Dmitri Mendeleev heard 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.

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 properties and looked for any trends.

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

8.  Mendeleev then created a table in which elements with similar properties were 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.

9.  He was correct in this assumption. He also left several empty spaces 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.

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?

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

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 organization of the periodic table.

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

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

15.  The periodic table has undergone extensive change since Mendeleev’s time. Chemists have discovered new elements and have synthesized many others. However, each of them can be placed in a group of other elements with similar properties.

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

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

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.

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

20.  Another major development came with the discovery of the actinide series. To save space these two groups are usually placed at the bottom of the periodic table.

21.  Periodicity with respect to atomic number can 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.

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

23.  The periodic table gives you a ton of information about the nature and characteristics of the elements if you are able to remember the explanations that you can use. There are only 4 possible explanations of all of the periodic trends.

24.  Effective nuclear charge, Z eff —essentially equal to the group number. Think of the atoms in the first column as having a Z eff of one while the halogens have a Z eff of 7! The idea is that the higher the Z eff, the more positive the nucleus, the more attractive force there is emanating from the nucleus drawing electrons in or holding them in place.

25.  Distance—attractive forces dissipate with increased distance. Distant electrons are held loosely and thus easily removed.

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.

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.

28.  Ideally, the size of an atom is defined by the edge of its orbital. However, this boundary is fuzzy and varies under different conditions. Therefore, to estimate the size of an atom, the conditions under which the atom exists must be specified.

29.  One way to express an atom’s radius is to measure the distance between the nuclei of 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

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 electrons in each of the atoms. Therefore, it makes sense that the atom should be getting larger.

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

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 principal energy level. From last chapter, we talked about electrons in higher energy levels being farther away from the nucleus.

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

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

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

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

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

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

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.

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

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 release energy when they acquire an electron. However, some require energy to accept an electron. When you see the graph in your book, those will have an electron affinity of zero.

42.  Electron affinity increases as you go across the table. Increasing Z eff causes the nucleus to attract the electron more. The more the electrons are pulled on by the nucleus, the more energy will be given off when the atom gains the electron. However, it should be noted that electron/electron repulsion plays a heavy role in the anomalies that occur in this trend.

43.  Electron affinity decreases down 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 reducing the energy that is given off when the electron joins the atom.

44.  Cation – a positive ion  Anion – a negative ion  The formation of a cation by the loss of one of more electrons always leads to a decrease in atomic radius because the removal of the highest energy level electrons results in a smaller electron cloud. Also, the remaining electrons are drawn closer to the nucleus by its unbalanced positive charge.

45.  The formation of an anion by the addition of one or more electrons will always lead to an increase in atomic radius. This is because the total positive charge of the nucleus remains the same when an electron is added to an atom or an ion.

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 spreads out because of greater repulsion between the increased amount of electrons.

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 smaller because Z eff is increasing. With the nuclear charge being larger, the nucleus pulls on the electrons more effectively, thereby shrinking the ion.

48.  As you go down a group, the ions increase because you have added an entire new principal energy level.

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

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

51.  As you go across the period, electronegativity increases because of Z eff. An increase in Z eff causes the atoms to become smaller. A smaller atom can then get their nucleus closer to the electrons of the other elements in the compound. If you get your nucleus closer to the electrons of other elements, you are able to attract those electrons more effectively.

52.  As you go down the group, electronegativity decreases because the atom is becoming larger. If the atom is larger, there is more distance between the nucleus of one atom and the electrons of the other atoms in the compound. Thus, the attraction is diminished.