E5: Unit E Reading Notes Page 6 of 6

E5: Unit E Reading Notes Page 6 of 6

E-1 Reading

By 1860 scientists had developed a reasonably reliable method for determining average atomic mass values for more than 60 elements. At this time, a Russian scientist by the name of Dmitri Mendeleev decided to use this information, along with his knowledge of the properties of the known elements, to organize elements. He found that as the atomic mass increased, there existed, with very few exceptions, a periodic repetition of the properties of elements (Think about the demonstration of the 12-tone musical scale that we did in class). The exceptions that he noted included the reversal of iodine (atomic mass (128) and tellurium (atomic mass 127) because their properties matched up that way. Mendeleev recognized that some elements were missing because their were holes in the atomic mass values, but he was able to accurately predict the properties and masses of these as yet undiscovered elements because of his first Periodic Table of the Elements.

Mendeleev’s work left some unanswered questions—Why did some of the elements not follow the pattern of periodic repetition of properties according to atomic mass?, and Why did the properties repeat in this way? The answer to the first question came after the discovery of electrons and the nucleus. Henry Moseley worked with Ernest Rutherford (remember his gold foil experiments?), and he noticed that the nuclear charge (the number of protons) was actually the determining factor in organizing the elements. He suggested in 1911 that arranging elements in order of increasing atomic number yielded a reliable repetition of element properties. The modern periodic law is now expressed as when the elements are arranged in order of increasing atomic number, elements with similar properties appear at regular intervals. This means that the periodic table of the elements is an arrangement of elements in order of their atomic numbers so that elements with similar properties fall in the same column, or group.

Elements that were missing at the time of Mendeleev’s work included all of the non-reactive noble gases. Their discovery toward the end of the 19th century led to the placement of a new group that was placed between group 17 and group 1. In addition, the lanthanides (atomic number 58-71, period 6) and the actinides (atomic number 90-103, period 7) were found to belong between groups 3 and 4. They are often placed below the regular periodic table in order to preserve its shape.

There are many properties of the elements that are periodic—the first is differences in atomic number as you move down a group: the pattern is always 8, 8, 18, 18, 32 no matter what group you are looking at. You will notice that this bears a striking resemblance to the pattern of electrons in each principle energy level (2n2), and in fact the arrangement of electrons around the nucleus is responsible in large part for the periodic repetition of element properties.

E-2 Reading

Types of Elements:

The periodic table is divided into 2 main sections, metals and nonmetals by a stairway at middle-right of the periodic table. It is clearly seen in the periodic table in the text, as well as on the pretty periodic table, B2, that is part of your unit handout. Any element to the left of the stairway is a metal (except hydrogen) and anything to the right is a nonmetal. The exceptions to this rule of thumb are the semimetals, or metalloids, located above and below the stairway treads. Specifically, these are boron (B), silicon (Si), germanium (Ge), arsenic (As), antimony (Sb), and tellurium (Te). Polonium (Po) and astatine (At) also exhibit some semimetallic nature.. Even though aluminum (Al) is located under boron, it is considered to be a metal because of its properties, which are outlined below.

Metals:

With the exception of mercury (Hg), all metals are solids at room temperature. They usually have shiny or lustrous surfaces. They conduct heat and electricity well, and they can be easily hammered or rolled into thin sheets (malleable) and able to be drawn in thin wires (ductile).

Nonmetals:

Many nonmetals are gases at room temperatures (nitrogen, oxygen, hydrogen, fluorine, and chlorine, and all of the noble gases), but one is a liquid (bromine) and the remainder are solids (carbon, sulfur, selenium, iodine, and phosphorus). Nonmetals are brittle rather than malleable and ductile. They usually have poor thermal (heat) and electrical conductivity.

Metalloids (Semimetals):

The metalloids are elements that have some characteristics of metals, and some of nonmetals. Pure metalloids are always solid at room temperature, and they tend to be less malleable than metals, but not as brittle as nonmetals. They tend to be semiconductors of electricity, which means that they are somewhere between the conducting abilities of metals and nonmetals.

Noble Gases:

These elements are located in Group 18 on the periodic table, and I remember their name by thinking of them as “snooty,” and not mingling with the riff-raff of the rest of the table. They do not usually react chemically with other elements, but scientists were finally able to force them to form compounds with other elements in the last 40 years.

The Group 18 elements of the periodic table (noble gases) do not react with other elements because of their extremely stable ns2p6 valence shell electron configuration. In fact the electron configuration of the valence shell is often what determines many of an atom’s chemical properties.

Rows or periods of the periodic table have different lengths. Towards the top of the table, period 1 represents the filling of the n=1 energy level. Because there is only the 1s sublevel at n=1, there are only two elements. Period 2 represents the addition of the n=2 energy level, which has the 2s and 2p sublevels, which can together hold 8 electrons. Therefore, there are only 8 elements in this period. The same is true for the 3rd period. Filling of the 3d sublevel in the 4th period (n=4) adds 10 more elements to this period.

It is clear from this that it is possible to determine the period in which an element is placed from its electron configuration. The text uses the element arsenic as an example: Its kernel notation electron configuration is [Ar] 4s23d104p3. This indicates that the highest energy level is n=4, which means that the element is located in period 4.

The periodic table can be divided up into blocks that describe the type of energy sublevel being filled. Groups 1 and 2 are designated the s-block elements, groups 3-12 are called the d-block elements, groups 13-18 are called the p-block elements, and the lanthanide and actinide series elements are called f-block elements.

Properties of the s-block elements (Groups 1 and 2)

These elements are chemically reactive metals. Group 1 is called the alkali metals. They all have only 1 electron in the valence shell (the outermost energy level), so the valence electron configuration is ns1. You will see in the next section that the ease with which these elements lose this single electron to achieve a noble gas configuration causes them to be extremely chemically reactive. They are all silvery in their pure state, and are soft enough to cut with a knife. They are never found in elemental form in nature, however, because they are so reactive—they combine vigorously with most nonmetals, and they react strongly (explosively) with water to produce hydrogen gas and aqueous solutions of substances called alkalis (sodium hydroxide, NaOH, is an example). They are often stored in kerosene to protect them from atmospheric moisture and oxygen. As you move down the group, the melting points have successively lower melting points.

Group 2 elements are called alkaline earth metals. They have a pair of electrons in the outer s sublevel, so the valence electron configuration for all of these elements is ns2. The elements in this group are harder, denser (more mass per unit volume), and stronger than the alkali metals, and they have higher melting points. They are very reactive as well, but not as reactive as the alkali metals. Hydrogen and helium should technically both considered to be s-block elements, but their properties are so very different from the alkali metals and alkaline earth metals that they deserve special consideration. Both are gases, so they are both nonmetals. Because helium has a full valence shell, it is extremely stable, and behaves as a noble gas. It is therefore placed in Group 18. Hydrogen, however, is extremely reactive, but it is capable of either losing its lone electron, or gaining an electron to obtain the electron configuration of helium to become more stable. It is also capable of sharing its valence electron with another atom. This makes hydrogen unique in its properties, so it is place above Group 1 based purely on its valence electron configuration.

Properties of the d-block elements (Groups 3-12)

The d-block elements begin in the 4th period after the 4s block is filled. Because it is the 3d sublevel that is filling, we say that the electron configuration takes the form ns2(n-1)d. There are going to be some deviations from this configuration as you move across the period. At times, you will notice on the periodic table that on of the s electrons will jump out of that sublevel and be promoted to the d-sublevel (e.g. chromium, Cr, in Group 6 or copper, Cu, in Group 11). It’s important to note, however, that the sum of the s and d electrons in the period will always be equal to the group number. This is true for all d-block elements within a given group.

The d-block elements (also called transition elements) are all metals with typical metal properties—they are all good electrical and thermal conductors, shiny, malleable, and ductile. They are usually less reactive than the alkali and alkaline earth metals. Many, like gold, platinum, and palladium, are so non-reactive that they frequently exist in nature in their elemental form.

Properties of p-block elements (Groups 13-18)

These elements include all group 13-18 elements except for helium. Electrons will only add to a p-sublevel after the s-sublevel at that main energy level is filled. This means that all p-block elements will already have 2 electrons in the ns-sublevel. The s- and p-block elements together constitute the main-group elements. You should note that all Group 13 elements have the valence shell electron configuration of ns2np1. Group 14 elements will be ns2np2, Groups 15 elements will be ns2np3, Group 16 elements will be ns2np4, Group 17 elements will be ns2np5, and Group 18 will be ns2np6. You will notice that for p-block elements, the number of outer energy level (valence shell) electrons is equal to the group number minus 10.

The properties of the p-block elements are all over the place. On the far right side of the block, these elements will be nonmetals. On the stairway, you will find the metalloids (semimetals) boron, silicon, germanium, arsenic, antimony, and tellurium. Below and to the left of the stairway these elements will be metals. See Figure 4 on pp. 120-121 to make sure you understand where everything is.

The Group 17 elements are called halogens. They are the most reactive of the nonmetals, and will react with most metals to form salts (e.g., sodium chloride, which is table salt) The reactivity of this group is based on the fact that there are 7 electrons in the valence shell, and they react to gain an electron and subsequently obtain an noble gas electronic configuration in the valence shell.

The metalloid (semiconducting) elements are mostly brittle solids with some properties of both metals and nonmetals. The metals of the p-block tend to be harder and denser than the s-block metals, but are generally softer and less dense than the transition metals. They also tend to be found most often in nature in compound form, which means that they are also generally more reactive than the transition metals. Once they are purified as free metals, however, they tend to be fairly stable in the presence of air (unlike Groups 1 and 2).

Properties of f-block elements: Lanthanides and Actinides

These elements are wedged between the Group 3 and 4 elements in periods 6 and 7. The elements in period 6 (lanthanides) are adding electrons at the 4f sublevel, which means that there will be 14 elements between lanthanum and hafnium. They are shiny metals which have reactivity similar to those in Group 2. The elements in Group7 (actinides) are adding electrons at the 5f sublevel, which means there are 14 of them between element 89, actinium (Ac), and element 104, Rutherfordium (Rf). All of the actinides are radioactive, and only thorium, protactinium, uranium, neptunium, and Plutonum occur naturally. All of the others have been synthesized by scientists.

E-3 Reading

We have already discussed how the properties of elements fall into patterns when they are arranged in order of increasing atomic number, but we haven’t been very specific about why these similarities occur within groups, nor have we discussed periodic and group trends in these properties. We will now examine these trends with respect to atomic radius, ionization energy, ionic radius, and electronegativity.

Before we discuss any of these trends, we need to discuss attractive and repulsive forces and the shielding effect within the atom. I am sure that you understand the concept that opposite charges attract, and like charges repel. There are two factors which affect the strength of these attractive and repulsive forces:

1.  Quantity of charge: