Periodic Table Groupings - MCAT Chemical and Physical Foundations of Biological Systems

Halogens are in the group next to the noble gasses. They have seven valence electrons, and therefore have a high electronegativity. The addition of only a single electron (production of an anion) generates a full valence octet.

Their diameters vary within the group. The diameter can be very small, like fluorine, or large, like iodine. They do not conduct electricity well, as they are non-metals.

Transition elements have multiple oxidation states because of the d-orbitals they possess. This allows them to lose or gain electrons in a variety of ways, often leading to the standard characteristics of metals, such as electrical conductivity. Lanthanides and actinides are less commonly tested, but also have the ability to form multiple oxidation states due to their large and variable orbitals. Ytterbium is one of the lanthanides, and has Yb(II) and Yb(III) oxidation states.

The correct answer is cobalt, since it is the only metal among the answer choices. Metals have all the properties described (malleability, ductility, and conductivity). Sulfur, boron, and silicon do not exhibit these properties to the same extent.

The halogens are the second to last column in the periodic table, meaning that they have an affinity for a single additional electron. Halogens would be most likely to react with alkali metals, which contain only one loosely bound electron in the valence shell. Alkali metals have very low ionization energy, readily losing an electron, while halogens have very high electronegativity, readily gaining an electron. This interaction allows the alkali metals to form ionic bonds with the halogens.

The properties described fit well with the noble gases. Alkali and alkaline earth elements are solid at room temperature, meaning that they have a high boiling point. Halogens can be gaseous at room temperature, but are very reactive. Noble gases have low boiling points and rarely act in spontaneous reactions. Their properties are due to their full valence shell, which is the source of their stability. Changes to their electron configuration (such as removing an electron) require large amounts of energy.

Alkali metals are much less dense than other metals due to their large radii, which results from having a single loosely bound valence electron. Some of the alkali metals have such low densities that they can float on water.

First off, it is important to know that the ability to be hammered into a sheet and conduct electricity are characterisitcs typically reserved for metals. With this in mind, we can check the periodic table, and see where on the table metals reside. Metallic character generally decreases as you go from left to right on the table, which results in nonmetals being found on the right side of the table. As a result, we would not expect to find this element and its metallic characteristics on the right side.

In general, the metals fall on the left side of the periodic table and are separated from the non-metals by the metalloids. Transition metals fall in the d block of the periodic table, in groups (columns) 3-12. Examples of metals, non-metals, transition metals, and metalloids are calcium, oxygen, zinc, and arsenic, respectively.

Alkali metals are a special class of metal only found in group 1 of the periodic table.

The non-metals in group 8 are called noble gases because they tend to resist reactions with other atoms. Noble gases are the only elements to have valence octets in their ground states ().

Halides are the gases in group 7, which are most stable as negative ions to reflect the octet configurations of the noble gases. Lanthanides are the elements in period 6 that have an incomplete f shell. The metalloids are arranged along the diagonal between boron and polonium and divide the periodic table between the metals (to the left) and nonmetals (to the right).

Elements are generally more reactive the closer they are to the stable noble gas configuration (eight valence electrons). Halogens have seven valence electrons, so only need one more to achieve a full valence shell, thus they are said to have high "electron affinity" since they react easily to gain the final valence electron. All the other options are correct statements about halogens.

Alkaline earth metals are found in the second group of the periodic table and include beryllium, magnesium, calcium, strontium, barium, and radium. These compounds are not as reactive as the alkali metals (found in group 1), but still participate in many reactions due to their electron configuration. Alkaline earth metals carry two valence electrons, located in the s orbital. Loss of these two electron leaves the alkaline earth metals with a full octet, giving them a stable oxidation state of +2. In contrast, the alkali metals have a stable oxidation state of +1. Both compounds have very low first ionization energies, but the second ionization energies of the alkaline earth metals are much lower than those of corresponding alkali metals. Removing a second electron from an alkali metal removes it from a stable octet, while removing an additional electron from an alkaline earth metal results in a stable octet.

While all alkali metal salts are soluble, the alkaline earth metals result in several exceptions to the solubility rules. For example, is not soluble in aqueous solutions.

The transition metals are defined in the region of the periodic table in which atoms are being added to the d subshell. As a result, the transition metals have unfilled or incomplete orbitals within the d shell. Since each orbital is filled with one electron before orbitals start to become completely filled, there are increasing numbers of unpaired d shell orbitals. This allows transition metals to give up variable numbers of electrons, while maintaining stability, as electrons move between d orbitals.

A common example is iron, which is stable in both the and electron configurations.

Valence orbitals are filled with one electron at a time until all orbitals of the same energy levels have one electron. Valence orbitals will then begin to be completely filled. The p orbital has three subshells, so three electrons will be in each p orbital before any one becomes completely filled.

Oxygen has the electron configuration . This means that oxygen will have two completely full s orbitals and four electrons distributed between the three p orbitals. These four electrons will first fill each orbital with a single electron, then add the fourth electrons to on p orbital to form a lone pair. The result is three lone pairs (2 s orbitals and 1 p orbital) and two unbonded electrons. This configuration is shared by all elements in the oxygen group, group 16.

The question states that the element is a solid at room temperature and conducts electricity. These are characteristics of a metal; therefore, the element is likely a metal. Recall that metals are usually found on the left side of the periodic table. The groups that are classified as metals include the alkali metals (group 1), alkaline earth metals (group 2), the aluminum family (group 3), and the transition metals (D block).

The answer choices state that it could be either sodium or silicon. Silicon is in group 14 and is considered a non-metal; therefore, the element has to be sodium. Metals are good electrical conductors and solid at room temperature. They are also ductile, which means that a metal can be stretched to create thin wires.

Metals are good reducing agents, not oxidizing agents. Recall that metals usually have one, two, or three valence electrons. Since they only have a few valence electrons, metals prefer to lose their valence electrons to complete an octet. When an element loses electrons it is considered to be oxidized and can act as a reducing agent; therefore, sodium is not an oxidizing agent.

One of the key distinguishing characteristics of metals and non-metals is the acidity of their oxides. Metal oxides are basic oxides, whereas non-metal oxides are acidic oxides. Remember that acids have low pH (high hydrogen ion concentration), whereas bases have high pH (low hydrogen ion concentration); therefore, non-metal oxides will have the lower pH.

A neutralization reaction is a special type of reaction that occurs between an acid and a base. This type of reaction will produce a pH of 7 if the reacting species are a strong acid and a strong base. Metal and non-metal oxides do not involve acid-base reactions; therefore, neutralization reaction is irrelevant to this question.

Metals are found on groups on the left side of the period table (groups 1, 2, and 3, and transition metals). The groups on the left side have fewer valence electrons than the groups on the right side. Groups 1, 2, and 3 have one, two, and three valence electrons, respectively. Metals thus have fewer valence electrons than non-metals.

Electronegativity is a chemical property that is defined as the ability of an element to attract electrons towards itself. Since they have fewer valence electrons, metals find it easier to lose electrons to generate a complete octet (have eight valence electrons). Rather than attract electrons to fill octet, metals give them away. On the other hand, it is easier for non-metals to gain electrons to complete octet because they have larger amounts of valence electrons in their ground state. This means that non-metals have higher attraction for electrons and, consequently, have higher electronegativities.

Transition metals are classified as metals; therefore, the oxides they form are called metal oxides. Metal oxides are basic compounds, whereas non-metal oxides are acidic compounds. Metal oxides, such as and , give rise to high pH values, and non-metal oxides, such as and , give rise to low pH values.

An oxidation state is defined as the degree of oxidation a compound can achieve. It is often calculated by observing the gain and loss of electrons in an atom. For example, an atom that loses two electrons will have an oxidation number of , whereas an atom that gains two electrons will have an oxidation number of . Some transition metals have multiple oxidation states because they can lose varying amounts of electrons. This occurs because the energy difference between the outermost orbital and the outermost orbital is small; therefore, the energy required to remove the electron from the subsequent orbital is comparable to removing electrons from the orbital.

For example, iron () can have an oxidation state of or . Iron’s electron configuration is . When it loses two electrons, iron will have an empty orbital. The electron configuration of will be . The amount of energy difference between the orbital and the orbital is very small; therefore, it is easy for iron to lose another electron from the orbital and become . On the other hand, it is very hard to remove electrons from a filled orbital. This explains why non-metals, such as the oxygen and halogen groups, and metals, such as alkali and alkaline earth metals, have only one oxidation state.

Reduction involves gaining electrons and oxidation involves losing electrons. This means that reduction will decrease the oxidation number and oxidation will increase the oxidation number. It is generally favorable for a transition metal to lose electrons and become oxidized, though reduction can be achieved by adding enough energy to the system (such as in a electrolytic cell).

Halogens are elements found in group 17 of the periodic table. They are characterized by their seven valence electrons. To complete an octet, halogens only need to gain a single electron. Since they gain electrons, halogens are usually reduced and serve as good oxidizing agents. Recall that when a substance is reduced it is also classified as an oxidizing agent because it can oxidize another atom (remove an electron from another atom).

Halogens gain an electron to complete an octet because they have seven valence electrons. The elements in the oxygen group (group 16) gain two electrons to complete an octet because they have six valence electrons.

Halogens are highly electronegative. Electronegativity is defined as the ability of an atom to pull electrons towards itself. Since they need only one electron to complete an octet, halogens have high attraction towards electrons and, consequently, have high electronegativity values. Note that electronegativity increases as you go from left to right on the periodic table; therefore, halogens have the highest electronegativity of any group. Noble gases (group 18) are inert and do not have any attraction for electrons (low electronegativity).

The question states that a stronger acid will possess a weaker bond between the hydrogen atom and the acid. To have a weak bond, it is essential to have a non-metal atom that will repel electrons from the hydrogen. This will enable the hydrogen atom to distance itself from the acid, which will make the bond weaker (it will make it easier to remove the hydrogen atom).

Recall that electronegativity of an atom is defined as the ability of the atom to pull electrons towards itself. An atom with a high electronegativity will have high attraction for electrons, whereas an atom with low electronegativity will have low attraction for electrons; therefore, to have a strong acid, the non-metal atom must have a low electronegativity. In the periodic table, the electronegativity decreases as you go from right to left and top to bottom. This means that iodine will have a lower electronegativity than fluorine and will form the strongest acid.

Remember that fluorine is the most electronegative atom in the entire periodic table. Every halogen, except fluorine, forms a strong acid. , , and are all strong acids ( is the strongest), whereas is a weak acid due to the strength of the bond between hydrogen and fluorine.