Chapter 8

We will refer to the Periodic Table throughout this chapter and we will be using the model in the inside front cover of your book, which has the first 2 columns and the last 6 columns labeled 1A, 2A, 3A etc. The elements in the A columns are called Representative Elements Some families have special names, IA are Alkali Metals, IIA are Alkaline Earth Metals, VIIA are Halogens and VIIIA are Noble Gases. All have incomplete s or p subshells of the highest n value, except for 8A, the Noble Gases, which have all filled s and p sublevels.

Alkali Metals - Active, lose one electron readily React with water ( example: 2Na + H 2 O  2NaOH) These reactions can be very vigorous; in fact K and Rb react violently with water. An important point to note about families: Elements in a family have similar properties, but not the same. The differences tend to follow a trend. For example, the reactivity of Alkali metals with H 2 O increases from Li to Cs. We will be investigating general trends later in this chapter.

Alkaline Earth Metals - Not as active but still relatively active Some react with water as alkali metals do, others don’t ( Be and Mg), but not as readily. Halogens - All exist as diatomic elements - all are quite reactive non-metals Most halogens occur in nature as salts not as free elements - Halogen means salt former - this is because of their reactivity Also react readily with many non-metals

Hydrogen is unique - sometimes behave like a non-metal other times like a metal. Common thread in all acids - will discuss this more in General Chem 2. Transition Metals – All the elements in columns headed by a 1B, 2B, 3B etc. Won’t spend very much time on them now – Just note that all are metals but none are as reactive as IA or IIA metals. Many transition metal compounds are colored. Most gem stones owe their characteristic color to transition metals present in the crystals. Inner Transition Metals - Don’t worry about them now.

Another aspect of the trends within families, is that for the representative elements, in most cases, if you predict the formula of a compound between any 2 representative elements (as we have already done in Chapter 2), then any of the elements in those 2 families will form the same formula. For example, we have shown how Na and Cl form the compound NaCl. That means that Cs and I will form the compound CsI and K and F form KF. There will be some exceptions, but this works most of the time.

Valence electrons - electrons involved in chemical bonding - outermost s + p electrons - outershell electrons The # of valence electrons can be found on the Periodic Table. The number in front of the letter at the top of any column gives the # of valence electrons for every element in that Family.

Noble Gases do not react chemically, basically, especially He, Ne and Ar. Chemists asked why? If you write the ground state electron configuration, all Noble Gases have completely filled Valence Shells. We will return to this in the next chapter. Whenever we represent elements in chemical equations, we simply write the symbol of that element, with the exception of the 7 diatomic elements that we learned in Chapter 2, (H 2, N 2, O 2, F 2, Cl 2, Br 2 and I 2 ).

We can also write electron configuration of ions formed from atoms, simply by noting whether electrons were added or removed. This is straightforward for Representative Elements. For transition elements, the first electrons to be removed are almost always the outermost s electrons, followed by d electrons, if necessary.

Remember that Mendeleev called his table of elements the Periodic Table, because properties repeated themselves, like the swing (or period) of a pendulum. We will now look at some of these properties and note the periodicity and also the trends within families that we noted earlier in this chapter.

Atomic Radius – The distance from the nucleus to the outermost electrons. The trend we see is that as you move from Li to Ne, the size gets smaller and then jumps to a much larger size with the next element Na. From Na to Ar the size again gets smaller and then jumps with K. This is the general trend. Size decreases from left to right and increases from top to bottom. There are some exceptions, but this is the general trend. Most of the other properties we look at will have their trends related to the size trend.

The size gets larger from top to bottom, because the valence electrons in each row are one energy level higher (i.e., further from the nucleus). There should be no surprise here. On the other hand, why should the size decrease from left to right in any particular row? Each succeeding atom has a higher atomic number and in most cases a higher atomic weight.

The explanation that is accepted is that as we move across the table each additional electron is added to the same principle energy level (i.e., approximately the same distance from the nucleus), but in each succeeding element the # of protons in the nucleus increases by 1 (i.e., the + charge in the nucleus increases by 1). The size of the atom is caused by the attraction of the nucleus for the electrons, in particular, the outermost electrons. Since the outermost electron in each succeeding element is basically going to the same place, but the attractive charge in the nucleus is increasing, the overall size decreases.

Looking at Figure 8.6 on page 325, we can easily see the Periodic Trend:

When an ion is formed, its size is different than the atom from which it was formed. Anions are always bigger than the neutral atom. Makes sense, because electrons are being added to the same nuclear charge; more repulsion. Cations are always smaller than the neutral atom; makes sense, because electrons are being removed from the same nuclear charge; stronger attraction. Also, frequently, the valence shell electrons are all removed, so that the new valence shell is closer to the nucleus. This is shown in Figure 8.7 on page 326:

Ionization Energy(IE) - The amount of energy needed to remove an electron from any atom. It always requires energy to remove an electron. The more tightly held an electron is, the higher the IE will be. Generally, the further an electron is from the nucleus, the weaker the attraction. Thus, the trend in IE is that it gets lower from top to bottom and higher from left to right. Basically following the size trend, smaller atom, higher IE. Removing additional electrons always requires significantly more energy. Especially when the additional electron is being removed from a Noble Gas Electron configuration.

These trends can be seen in Table 8.2 on page 331.

Electron Affinity – The negative of the energy change that occurs when an electron is added to an atom in the gaseous form to form an anion. Basically, the higher the E.A., the easier it is to add an electron. Basically E.A. decreases as you go down a column and increases as you go across, with some exceptions they we won’t worry about. You should read the remainder of Chapter 8, which discusses some specific chemical behavior of each group of representative elements.