Elements and the Periodic Table - GRE Subject Test: Chemistry

Nitrogen, phosphorous, antimony, and bismuth are all in the same group (column) of the periodic table.

The atomic radius increases from the top of a group to the bottom, due to increased principle shell number (n). As one travels down a group, another s shell is added, meaning that electrons are added in another orbit farther from the nucleus. This serves to increase the atomic radius of the atom.

This question refers to electron affinity, which is defined as the energy given off when a neutral atom in the gas phase gains an extra electron.

Electron affinity increases for elements towards the top and right of the periodic table, so the elements in the top right lose the most energy when gaining an electron. Another way of thinking is that they lose energy, but gain stability. Of the available answers, the element to the most upper right of the periodic table is fluorine.

Ionization energy is a property of an atom that describes the lowest energy needed to remove a electron from its valence shell in the ground state.

Ionization energy increases up and to the right on the periodic table. All the elements listed are halogens located in group 7A of the periodic table. Because fluorine is located at the top of the periodic table in group 7A, it will have the highest first ionization energy. Ionization energy can be though of as the inverse of atomic radius. That is, the closer the valence shell of electrons is to the nucleus, the larger the magnitude of electric attractive force, and thus, the more energy is required to strip that electron.

Ionization energy is a property of an atom or ion that describes the lowest energy needed to remove a electron from it in the ground state.

Ionization energy increases up and to the right on the periodic table. Because oxygen is the highest and rightmost element of those listed, it will have the highest first ionization energy. Ionization energy can be though of as the inverse of atomic radius. That is, the closer the valence shell of electrons is to the nucleus, the larger the magnitude of electric attractive force, and thus, the more energy is required to strip that electron.

Ionization energy is a property of an atom or ion that describes the lowest energy needed to remove a electron from it in the ground state.

Ionization energy increases up and to the right on the periodic table. All the elements listed are chalcogens located in group 6A of the periodic table. Because oxygen is located furthest to the top right on the periodic table in group 6A, it will have the highest first ionization energy.

Ionization energy is a property of an atom or ion that describes the lowest energy needed to remove a electron from it in the ground state.

Ionization energy increases up and to the right on the periodic table. Because fluorine is located furthest to the top right on the periodic table in group 7A, it will have the highest first ionization energy.

Lithium, boron, oxygen, and neon are all in the same row (period) of the periodic table.

The atomic radius decreases from left to right along a period due to increased effective nuclear force. From left to right the atomic number increases, indicating that more protons are added. The addition of protons increases the positive charge in the nucleus, pulling in the outer electrons by increasing the effective nuclear force, decreasing the radius.

In math terms, we can equate effective nuclear force using the force equation between two charged particles.

We can see that the farther apart the electrons and protons are, the less the force is between them.

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.

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 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 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.

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.

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).

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.

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.

Atomic radius increases down each group of the periodic table and toward the left of each period. Since the elements listed are all in the same group, iodine would have the greatest atomic radius because it farther down the period compared to the others.

Atomic radius can be determined using the periodic trends. Atomic radius increases to the left of a period and down a group of the periodic table. Electronegativity, in contrast, increases to the right of a period and up a group of the periodic table. Relating the two, we can see that the greater the atomic radius, the weaker its electronegativity because the electrons are farther away from the nucleus and are unable to feel the attractive force of the protons in the nucleus.

Energy level increases moving down a group of the periodic table. As energy level increases, the outer valence shell becomes more distant from the nucleus, causing atomic radius to increase.

Energy level remains constant across a period, but electrons are added within the same orbitals. When new electrons are added within the same orbital, additional protons are also added to the nucleus. This increases the effective nuclear charge, pulling the electrons closer to the nucleus. The trend for atomic radius is to decrease as we move right along a row.

This means that the general trend for atomic radius is to increase as one moves to the left and downward on the periodic table.

Atomic radius increases with increasing effective nuclear charge (Z). Elements toward the right and toward the top of the periodic table have the highest Z values. Protons and electrons are added in pairs as we traverse the periodic table from left to right. A attractive force is established between the positively-charged nucleus and the negatively-charged electron cloud, which increases as the number of particles grows.

When electrons are added or taken away without the same happening to a proton, an imbalance of charge accumulates. When more electrons are present than normal, the electron cloud sags farther away from the nucleus. When fewer electrons are present than normal, the electron cloud is drawn in more tightly toward the nucleus. Atoms with extra electrons (a negative charge) will have larger nuclei than their neutral counterparts. A chloride ion will thus has a larger atomic radius than argon, a potassium ion, or a calcium ion.

The trend for atomic radius is to increase going from top to bottom, as additional valence shells are added to the atom. Out of the answer choices, rubidium has the highest energy valence shell.

With a single electron in the fifth energy level, krypton will have the highest number of energy levels of the group I elements listed.

When moving across a period, atomic radius will decrease as the number of protons increases. These protons increase the attraction between the high-energy electrons and the nucleus, effectively "shrinking" the electron cloud.