Periodic Trend: Cumulative: Videos & Practice Problems

Understanding periodic trends is crucial for grasping the behavior of elements within the periodic table. As we consider movement towards the top right corner of the periodic table, we observe two primary types of trends. Metallic character and atomic radius, which we can categorize as Type A trends, tend to decrease. On the other hand, properties that start with 'I' or 'E,' such as ionization energy, electron affinity, electronegativity, and effective nuclear charge, increase and are classified as Type B trends. It's important to note that ionic radius is an exception and does not fit neatly into these categories, as it is influenced by the number of electrons. Remembering these trends can significantly aid in predicting and understanding the chemical properties and reactivity of elements.

Overview of Periodic Trends

Periodic trends in the periodic table can be categorized into two main groups based on their behavior as you move towards the top right corner. The first group, referred to as Type A trends, includes metallic character and atomic radius. Both of these trends exhibit a decrease as you progress towards the top right corner of the periodic table. This means that elements become less metallic and their atomic size diminishes in this direction.

The second group, known as Type B trends, encompasses ionization energy, electron affinity, electronegativity, and effective nuclear charge. These trends show an increase as you move towards the top right corner. This indicates that elements require more energy to remove an electron, have a greater tendency to attract electrons, and possess a stronger effective nuclear charge in this region.

It is important to note that ionic radius does not fit neatly into either category, as it is influenced by the number of electrons and their arrangement. Understanding these trends allows for a comprehensive overview of how elements behave in relation to one another on the periodic table, facilitating easier predictions of their chemical properties.

Ionization Energy and Electron Affinity Example

In the study of elements, understanding the relationship between ionization energy and electron affinity is crucial for predicting the behavior of metals. Elements that exhibit low first ionization energies and more endothermic electron affinities are typically those that form cations, which are positively charged ions. This characteristic is predominantly found in metals, as they tend to lose electrons easily, facilitating the formation of cations.

Ionization energy refers to the energy required to remove an electron from an atom. For metals, this energy is generally low, indicating that it is relatively easy to remove an electron and form a positive ion. Consequently, when considering the trends in the periodic table, metals are located on the left side, where ionization energies are lower compared to nonmetals on the right side.

On the other hand, electron affinity measures the energy change when an electron is added to a neutral atom. As one moves towards the top right corner of the periodic table, electron affinity tends to increase and become more exothermic. However, for metals, which are situated more towards the left, the trend is reversed; they exhibit more endothermic electron affinities, meaning they do not readily gain electrons.

In summary, when evaluating elements that form cations, it is essential to recognize that metals, characterized by low ionization energies and endothermic electron affinities, are the primary candidates. This understanding aids in predicting the chemical behavior of various elements in reactions.

Which of the following statements is false? 

• A. The first ionization energy of chlorine is larger than magnesium.
• B. Anions are larger than their parent atoms.
• C. The removal of the first electron from boron is highly exothermic.
• D. The oxygen atom is smaller than the barium atom.
• E. All of the statements are true.

Answer: (C)

For main-group elements, the atomic radius increases going down a group because: 

• A. Effective nuclear charge decreases.
• B. The mass of an element increases.
• C. The highest principal quantum number of the valence orbitals increases.
• D. Electron affinity increases.
• E. Both effective nuclear charge and the principal quantum number of valence orbitals increases.

Answer: (A)

Which of the following statements is false? 
a) The electron affinity of phosphorus is less exothermic than those of both silicon and sulfur. 
b) The second ionization energy of oxygen is less than the first ionization of fluorine.
c) As the principal quantum number of an atom increases the effective nuclear charge will increase. 
d) Cations are smaller than their parent atoms. 
e) The second electron affinity of an atom will be more exothermic than the first electron affinity. 

• A. b only
• B. b and c
• C. a, c and e
• D. a and c
• E. b, c and e

Answer: (E)

Which of the following statements is true? 

• A. Valence electrons are more difficult to remove than inner core electrons.
• B. Argon and bromide are isoelectronic species.
• C. Francium possesses the greatest electronegativity.
• D. The second ionization energy of an atom is always larger than its first ionization energy.
• E. The aluminum atom will experience a larger jump in its third ionization energy than its fourth.

Answer: (D)

Here’s what students ask on this topic:

What are the periodic trends for metallic character and atomic radius?

Metallic character and atomic radius are categorized as Type A periodic trends. As you move towards the top right corner of the periodic table, both metallic character and atomic radius decrease. This means that elements become less metallic and their atomic size gets smaller as you go from left to right across a period and from bottom to top within a group. This trend is due to the increasing effective nuclear charge, which pulls the electrons closer to the nucleus, reducing the atomic radius and metallic character.

How do ionization energy, electron affinity, and electronegativity change across the periodic table?

Ionization energy, electron affinity, and electronegativity are classified as Type B periodic trends. As you move towards the top right corner of the periodic table, these properties increase. Ionization energy is the energy required to remove an electron from an atom, electron affinity is the energy change when an electron is added to an atom, and electronegativity is the ability of an atom to attract electrons in a chemical bond. These trends increase due to the higher effective nuclear charge, which makes it more difficult to remove electrons and easier to attract additional electrons.

Why does the ionic radius not fit into the Type A or Type B periodic trends?

The ionic radius does not fit neatly into the Type A or Type B periodic trends because it depends on the number of electrons in the ion. When an atom loses electrons to form a cation, its ionic radius decreases due to the reduced electron-electron repulsion and increased effective nuclear charge. Conversely, when an atom gains electrons to form an anion, its ionic radius increases due to the added electron-electron repulsion. Therefore, the ionic radius is influenced by the specific electron configuration of the ion, rather than following a simple trend across the periodic table.

How can understanding periodic trends help in predicting the chemical properties of elements?

Understanding periodic trends is crucial for predicting the chemical properties and reactivity of elements. By knowing how properties like atomic radius, ionization energy, electron affinity, and electronegativity change across the periodic table, you can infer how elements will behave in chemical reactions. For example, elements with low ionization energy and high metallic character are more likely to lose electrons and form cations, making them good reducing agents. Conversely, elements with high electronegativity and electron affinity are more likely to gain electrons and form anions, making them good oxidizing agents. These trends help chemists predict bonding patterns, reactivity, and the stability of compounds.

What is the effective nuclear charge and how does it affect periodic trends?

Effective nuclear charge (Z_eff) is the net positive charge experienced by an electron in a multi-electron atom. It is calculated by subtracting the shielding effect of inner electrons from the total nuclear charge. Z_eff increases across a period from left to right due to the addition of protons in the nucleus, which increases the attraction between the nucleus and the valence electrons. This increase in Z_eff causes a decrease in atomic radius and metallic character, and an increase in ionization energy, electron affinity, and electronegativity. Understanding Z_eff helps explain why these periodic trends occur and how they influence the behavior of elements.