Question

Arrange the following atoms in order of increasing ionization energy: Si, K, P, and Ca..

Short answer

K < Ca < Si < P.

Step-by-step solution

Understand Ionization Energy

Ionization energy is the energy required to remove an electron from an atom. Generally, ionization energy increases across a period (from left to right on the periodic table) and decreases down a group (from top to bottom).

Identify Periodic Table Positions

Locate the elements on the periodic table: Silicon (Si, Group 14, Period 3), Potassium (K, Group 1, Period 4), Phosphorus (P, Group 15, Period 3), and Calcium (Ca, Group 2, Period 4).

Determine General Trends

Si and P are in the same period (Period 3), with P further to the right, indicating a higher ionization energy than Si. K and Ca are in the same period (Period 4), with Ca being to the right of K, so Ca has higher ionization energy than K.

Compare Across Periods

Period 3 elements (Si and P) typically have higher ionization energies than Period 4 elements (K and Ca) because ionization energy increases across periods while decreasing down groups.

Arrange Based on Rules

First, arrange Period 4 elements: K has lower ionization energy than Ca. Then arrange them relative to Period 3: K < Ca. Now compare this with Period 3: Si < P. The overall arrangement considering all the atoms is K < Ca < Si < P.

Key concepts

Periodic Table Trends

In chemistry, understanding the trends of the periodic table is crucial for predicting and explaining the behaviors of elements. One of these important trends is ionization energy, which generally increases as you move from left to right across a period and decreases as you move down a group. These trends occur because of the electronic configurations of atoms and the effect of nuclear charge.

• Across a period: Atomic numbers increase, and electrons fill up the same energy level (orbital), which means the electron cloud becomes more compact. This results in greater attraction between the nucleus and the outer electrons, thus higher ionization energy.
• Down a group: Additional electron shells are added, increasing the distance between the outermost electrons and the nucleus, thus making the electrons easier to remove and decreasing ionization energy.

These trends help predict which elements will have higher ionization energy, crucial for understanding their reactivity and bonding characteristics.

Atomic Structure

A fundamental aspect of chemistry is atomic structure, which defines how electrons are arranged around the nucleus of an atom. Each atom consists of a dense nucleus made of protons and neutrons, surrounded by electrons organized in various energy levels or electron shells.

• The number of protons, which defines the atomic number, remains constant for any given element and essentially determines the nature and identity of the element.
• Electrons are arranged around the nucleus in shells or energy levels, with each higher shell being further away from the nucleus. As the shell level increases, the energy needed to remove an electron decreases due to the increased distance and reduced interaction with the positively charged nucleus.
• Valence electrons, located in the outermost shell, are most relevant for chemical reactions and ionization processes because they are subject to removal or sharing.

Understanding atomic structure is essential to grasp how elements interact and why ionization energy varies across different elements.

Electron Removal

The concept of ionization energy revolves around the process of electron removal from an atom. This requires energy due to the electrostatic attraction between the negatively charged electron and the positively charged nucleus. Higher ionization energies indicate a stronger attraction between the nucleus and the electrons and, therefore, make electron removal more difficult.

• First ionization energy refers to the energy needed to remove the outermost (valence) electron from an atom in its gaseous state.
• Subsequent ionization energies usually increase because removing additional electrons disturbs the stable electronic arrangement and increases the effective nuclear charge felt by the remaining electrons.
• In general, the ease or difficulty of removing an electron is influenced by the atomic size, nuclear charge, and electron configuration.

These concepts are essential for understanding how elements form ions and participate in chemical bonding.

Periodicity in Chemistry

Periodicity refers to the recurring trends that we observe in the elements of the periodic table. One critical manifestation of periodicity is found in ionization energy. This trend, as with others, helps chemists understand the structure, properties, and reactivity of different elements.

• The systematic increase of ionization energy across a period can be attributed to the progressive increase in nuclear charge without any additional shielding effect, leading to a stronger attraction between the nucleus and the electrons.
• The decrease in ionization energy down a group occurs due to increased electron shielding and greater atomic radii, which outweighs the increase in nuclear charge, making electrons easier to remove.

These patterns are critical for predicting the behavior of elements in chemical reactions and for understanding various physical and chemical properties, such as reactivity and bonding capabilities. Being familiar with these periodic trends allows chemists to draw insights into atomic and molecular behavior.