15 Aug, 2024
· Chemistry

Arrange these elements according to first ionization energy

The elements provided are:

  • Ar
  • Cl
  • S
  • P
  • Si
  • Al
  • Mg
  • Na
Short Answer
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Long Explanation

Explanation

Arrangement by First Ionization Energy

The first ionization energy generally increases across a period from left to right and decreases down a group in the periodic table. Based on this trend, we can arrange the given elements accordingly.

Important formula:

Ionization Energy(E)=hcR∞Z2n2\text{Ionization Energy} (E) = \frac{hcR_{\infty}Z^2}{n^2}

Where:

  • EE is Ionization Energy
  • hh is Planck's constant, c is the speed of light
  • R∞R∞ is Rydberg constant
  • ZZ is atomic number
  • nn is the principal quantum number

Given elements are: Ar, Cl, Si, P, AI, Mg, Na

Arranged according to first ionization energy:

Ar>Cl>P>Si>Mg>AI>Na\begin{array}{c} \text{Ar} > \text{Cl} > \text{P} > \text{Si} > \text{Mg} > \text{AI} > \text{Na} \end{array}

Breakdown

  1. Ar (Argon): Being a noble gas, it has the highest ionization energy.
  2. Cl (Chlorine): High due to non-metal and closer to noble gas.
  3. P (Phosphorus): Higher than Si due to its position in the same period and less shielding effect.
  4. Si (Silicon): Next in the period, hence lower than P.
  5. Mg (Magnesium): Higher than AI as it precedes it in the periodic table.
  6. AI (Aluminium): Lower than Mg but higher than Na.
  7. Na (Sodium): The lowest in the list, belonging to the alkali metals which have low ionization energy.

Conclusion

The order reflects the trend in ionization energies with argon (Ar) having the highest and sodium (Na) having the lowest.

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Johnathan Clark

Chemistry Content Writer at Math AI

Johnathan Clark, with a Master's in Chemistry from the University of São Paulo, is a young high school chemistry teacher and part-time contract writer. His engaging classroom experiments translate into compelling written content that makes chemistry exciting and practical.

chemistry
Concept

Periodicity Trends In The Periodic Table

Understanding Periodicity Trends in the Periodic Table

Periodicity trends in the periodic table refer to the predictable patterns that occur in the properties of elements as you move across a period (row) or down a group (column). These trends arise due to the periodic nature of the elements' atomic structures.

Atomic Radius

The atomic radius is the distance from the nucleus of an atom to the outermost electron shell. Across a period, the atomic radius decreases from left to right due to increasing nuclear charge, which pulls electrons closer to the nucleus.

Atomic Radius∝1Effective Nuclear Charge\text{Atomic Radius} \propto \frac{1}{\text{Effective Nuclear Charge}}

Down a group, the atomic radius increases because each successive element has an additional electron shell.

Ionization Energy

Ionization energy is the energy required to remove an electron from a gaseous atom or ion. Generally, ionization energy increases across a period due to the increased nuclear charge.

Ionization Energy∝Effective Nuclear Charge\text{Ionization Energy} \propto \text{Effective Nuclear Charge}

Conversely, ionization energy decreases down a group since the outermost electrons are farther from the nucleus and thus easier to remove.

Electron Affinity

Electron affinity is the energy change when an electron is added to a neutral atom to form a negative ion. Across a period, electron affinity generally becomes more negative (higher electron affinity), indicating that the atom more readily accepts an electron.

Electron Affinity=−ΔH\text{Electron Affinity} = -\Delta H

Down a group, electron affinity generally becomes less negative as the added electron is further from the nucleus and experiences less attraction.

Electronegativity

Electronegativity is a measure of an atom's ability to attract and bond with electrons. It generally increases across a period due to higher nuclear charges and decreases down a group as atomic size increases.

Electronegativity≈Ionization Energy+\text{Electronegativity} \approx \text{Ionization Energy} + +Electron Affinity+ \text{Electron Affinity}

Understanding these trends helps in predicting the chemical and physical properties of elements and in elucidating the behavior of elements in chemical reactions.

Concept

Atomic Structure And Shielding Effect

Explanation

Atomic structure refers to the arrangement of subatomic particles—protons, neutrons, and electrons—within an atom. Protons and neutrons are located in the atom's nucleus, while electrons orbit the nucleus in energy levels or shells. The atomic number is equal to the number of protons, which defines the element, and the mass number is the sum of protons and neutrons.

Electron Configuration

The distribution of electrons in an atom's orbitals follows the principles of quantum mechanics and is described by the electron configuration. Electrons occupy orbitals in such a way to minimize the energy of the atom, filling up lower energy levels before moving to higher ones (Aufbau principle). The configuration is written as:

1s2 2s2 2p6 3s2 3p6 …1s^2 \ 2s^2 \ 2p^6 \ 3s^2 \ 3p^6 \ \ldots

Shielding Effect

The shielding effect (also called screening effect) refers to the reduction in the effective nuclear charge on the electron cloud, due to a repulsion provided by other electrons. This effect is significant in multi-electron atoms, where inner-shell electrons partially shield outer-shell electrons from the full charge of the nucleus.

Calculating Effective Nuclear Charge

The effective nuclear charge (ZeffZ_{\text{eff}}) experienced by an electron can be approximated by:

Zeff=Z−SZ_{\text{eff}} = Z - S

Where:

  • ZZ is the atomic number
  • SS is the shielding constant, which represents the average number of electrons between the nucleus and the electron in question.

Importance of Shielding Effect

  • Periodic Trends: Explains trends in atomic size, ionization energy, and electron affinity across the periodic table.
  • Chemical Reactivity: Influences how easily atoms form bonds; for instance, it affects the reactivity of alkali and alkaline earth metals.

By understanding the interplay of atomic structure and shielding effect, chemists can predict and explain a wide variety of chemical behavior and properties.