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 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.

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 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.

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).

Nonmetals are generally brittle elements that have lower melting points than metals. They also form covalent bonds with one another due to minimal difference in electronegativity between nonmetal elements.

Nonmetals most frequently form negative ions in solution, called anions. In contrast, metals most frequently form positive ions in solution, known as cations.

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.

Elemental notation can provide you with a variety of information. The letter(s) indicate the element in question, the bottom number represents the atomic number (number of protons), while the top number represents the mass number (number of protons and neutrons). In order to find the number of neutrons in an atom, simply subtract the atomic number from the mass number. Doing this with our oxygen atom reveals that it has eight neutrons. The only answer option that has the same number would be the nitrogen atom, because 15 minus 7 is equal to 8.

Elements within a group have the same number of valence electrons, but in increasing energy levels. Elements toward the bottom of a group have valence electrons with higher energies in larger orbitals. This results in a larger radius and a weaker attractive force between the nucleus and outer electrons. The ionization energy decreases as the electrons are more removed from the attraction of the nucleus.

When moving down a group, atomic radius increases and ionization energy decreases.

Ionization energy is the energy that an atom must absorb in order to release an electron. Metals have fairly low ionization energies, up until they develop a noble gas electron configuration. At that point, the ionization energy will spike due to the stability of the noble gas configuration. For example, sodium has a very low first ionization energy because it has only one valence electron, but a very high second ionization energy because removal of a second electron disrupts the noble gas configuration of the ion.

Based on its position on the periodic table, scandium shows that it can release three electrons before it develops a noble gas electron configuration (hence its tendency to become an ion with a charge of +3).

At this point, the atom will require a massive amount of absorbed energy in order to release a fourth electron. This means that the 4th ionization energy for scandium is the highest of the given options.