Periodic Trends - Chapter 8 Quiz For AP Chemistry

The atom with the greatest atomic radius is C. Atomic radius refers to the size of an atom, and it generally increases as you move down a group and decreases as you move across a period in the periodic table. C is located to the left of N in the same period, so it has a larger atomic radius.

The atomic radius of an atom is determined by the number of electron shells and the effective nuclear charge. Ga has a larger atomic radius compared to the other atoms because it has more electron shells. As we move down a group in the periodic table, the number of electron shells increases, leading to larger atomic radii. Therefore, Ga, being located below Ge, As, and Br in the periodic table, has the largest atomic radius.

Silicon (Si) would most likely have some properties of both metals and nonmetals. It is a metalloid, which means it has characteristics of both metals and nonmetals. Silicon is a semiconductor and has properties such as high thermal and electrical conductivity like metals, but it also has some nonmetallic properties like brittle solid structure and ability to form covalent bonds. Therefore, Si is the element that would most likely have some properties of both metals and nonmetals.

The first ionization energy is the energy required to remove an electron from an atom. As we move across a period from left to right, the first ionization energy generally increases because the atomic radius decreases and the effective nuclear charge increases. Among the given elements, potassium (K) has the lowest first ionization energy because it is the furthest to the left in the periodic table. Calcium (Ca) has a higher first ionization energy than potassium because it is one step to the right. Phosphorus (P) has a higher first ionization energy than calcium because it is further to the right. Fluorine (F) has a higher first ionization energy than phosphorus because it is even further to the right. Neon (Ne) has the highest first ionization energy among the given elements because it is the furthest to the right in the periodic table. Therefore, the correct order of increasing first ionization energy is K

The electron affinity is a measure of an atom's tendency to gain an electron. It represents the energy change that occurs when an atom gains an electron to form a negative ion. Among the given elements, nitrogen (N) has the greatest electron affinity. This is because nitrogen has a small atomic radius and a high effective nuclear charge, making it more likely to attract and hold an additional electron compared to the other elements listed.

The compound P2O5 is a molecular compound because it consists of two phosphorus atoms and five oxygen atoms bonded together. It also has covalent bonds because the atoms share electrons in the bonding. However, this compound is not basic when added to water. In fact, it is acidic and forms phosphoric acid (H3PO4) when dissolved in water.

Potassium is the most likely atom R because it readily forms ionic compounds, such as potassium chloride, by losing its single valence electron. Phosphorus is in Group 5A, so it requires three electrons to achieve a stable electron configuration. Potassium can easily donate one electron to phosphorus, resulting in the formation of the ionic compound R3P. Aluminum, magnesium, and nitrogen are less likely because they do not readily form ionic compounds with phosphorus.

The electron configuration of element X indicates that it has 6 valence electrons (2 in the 3s orbital and 4 in the 3p orbital). Aluminum (Al) has 3 valence electrons. In order to achieve a stable octet (8 valence electrons), 2 atoms of X would need to combine with 3 atoms of Al. Therefore, the formula of the compound formed would be Al2X3.

The correct answer is O because it has the fewest number of electron shells compared to the other atoms listed. As you move across a period in the periodic table, the atomic radius generally decreases because there is an increase in the number of protons in the nucleus, resulting in a stronger attractive force on the electrons. Additionally, within a period, the atomic radius tends to decrease as you move from left to right. Therefore, since O is located towards the right side of the periodic table, it has a smaller atomic radius compared to Br, Cl, and At.

Strontium would have the greatest difference between its second and third ionization energies. This is because ionization energy generally increases as you remove electrons from an atom. Since strontium has a larger atomic number compared to the other elements listed, it has more electrons in its outermost energy level. As a result, the second ionization energy of strontium, which requires removing a second electron, would be significantly higher than the first ionization energy. The third ionization energy, which requires removing a third electron, would be even higher, resulting in the greatest difference between the second and third ionization energies for strontium.

Ar, or Argon, has the highest effective nuclear charge among the given elements. Effective nuclear charge refers to the positive charge experienced by an electron in the outermost energy level of an atom. It is determined by the number of protons in the nucleus and the shielding effect of inner electrons. Ar has 18 protons in its nucleus, and its outermost electrons experience a higher effective nuclear charge due to the absence of inner electron shielding. Therefore, Ar has the highest effective nuclear charge among Ga, Ca, Ar, and S.

The element with the highest SECOND ionization energy is Li. This is because Li has a relatively low atomic number and a small atomic radius, which means that its outermost electron is held more tightly by the nucleus compared to the other elements listed. As a result, it requires more energy to remove the second electron from Li compared to the other elements, making it have the highest second ionization energy.

Phosphide (P3-) has the highest ionic radius among the given ions. This is because phosphorus (P) is a larger atom with more electron shells compared to the other ions. As a result, the P3- ion will have a larger radius due to the increased number of electrons and electron-electron repulsion, causing the electron cloud to spread out more.

The correct answer is because chlorine has less of a shielding effect. Electronegativity is the measure of an atom's ability to attract electrons towards itself in a chemical bond. It is influenced by factors such as effective nuclear charge and shielding effect. The shielding effect refers to the ability of inner electron shells to shield the outer electrons from the full positive charge of the nucleus. Chlorine has a smaller atomic radius and fewer inner electron shells compared to iodine. This means that the outer electrons in chlorine experience a greater pull from the nucleus, resulting in higher electronegativity.

Rb (Rubidium) would be the most reactive element among Fe (Iron), Al (Aluminum), Rb (Rubidium), and Na (Sodium). Reactivity of an element is determined by its ability to gain or lose electrons. Rb has a single valence electron in its outermost shell, which makes it highly reactive as it readily loses this electron to form a positive ion. On the other hand, Fe, Al, and Na have multiple valence electrons and are less likely to undergo a reaction. Therefore, Rb is the most reactive element among the given options.