Ch.6 - Ionic Compounds: Periodic Trends and Bonding Theory

McMurry8th EditionChemistryISBN: 9781292336145Not the one you use?Change textbook

McMurry 8th Edition

Ch.6 - Ionic Compounds: Periodic Trends and Bonding Theory

Problem 120

Chapter 6, Problem 120

Order the following atoms according to increasing atomic radius: S, F, O.

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Identify the periodic trend for atomic radius: Atomic radius generally increases as you move down a group and decreases as you move across a period from left to right.

Locate the elements on the periodic table: Sulfur (S), Fluorine (F), and Oxygen (O) are all in the same period (Period 2).

Compare the positions of the elements: In Period 2, the order from left to right is O, F, S.

Apply the periodic trend: Since atomic radius decreases from left to right across a period, the order of increasing atomic radius is F, O, S.

Conclude the order: Therefore, the order of increasing atomic radius is F < O < S.

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Key Concepts

Here are the essential concepts you must grasp in order to answer the question correctly.

Atomic Radius

Atomic radius is a measure of the size of an atom, typically defined as the distance from the nucleus to the outermost electron shell. It can vary based on the atom's position in the periodic table, influenced by factors such as nuclear charge and electron shielding. Generally, atomic radius increases down a group and decreases across a period.

Periodic Trends

Periodic trends refer to the predictable patterns observed in the properties of elements as you move across or down the periodic table. For atomic radius, as you move from left to right across a period, the radius decreases due to increased nuclear charge attracting electrons more strongly. Conversely, as you move down a group, the radius increases due to the addition of electron shells.

Electronegativity

Electronegativity is the tendency of an atom to attract electrons in a chemical bond. It is closely related to atomic radius; smaller atoms with higher electronegativity tend to have a stronger pull on electrons. In the context of the question, understanding electronegativity helps explain why fluorine, being the most electronegative and smallest atom among S, F, and O, has the smallest atomic radius.

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Electronegativity Trends

Related Practice

Textbook Question

Take a guess. What do you think is a likely ground-state electron configuration for the sodium ion, Na+, formed by loss of an electron from a neutral sodium atom?

Textbook Question

Iron is commonly found as Fe, Fe²⁺+, and Fe³⁺. (c) The third ionization energy of Fe is E_i3 = +2952 kJ/mol. What is the longest wavelength of light that could ionize Fe²⁺(g) to Fe³⁺(g)?

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Textbook Question

The ionization energy of an atom can be measured by photo-electron spectroscopy, in which light of wavelength l is directed at an atom, causing an electron to be ejected. The kinetic energy of the ejected electron 1EK2 is measured by determining its velocity, v since EK = 1/2 mv2. The Ei is then calculated using the relationship that the energy of the inci-dent light equals the sum of Ei plus EK. (a) What is the ionization energy of rubidium atoms in kilo-joules per mole if light with l = 58.4 nm produces elec-trons with a velocity of 2.450 * 106m/s? (The mass of an electron is 9.109 * 10-31 kg.)

Textbook Question

Which atom in each of the following pairs has a larger radius?(a) Na or K(b) V or Ta(c) V or Zn(d) Li or Ba

Textbook Question

Which atom in each of the following pairs has a larger radius?(a) C or Ge(b) Ni or Pt(c) Sn or I(d) Na or Rb

Textbook Question

One method for calculating Zeff is to use the equation

where E is the energy necessary to remove an electron from an atom and n is the principal quantum number of the elec-tron. Use this equation to calculate Zeff values for the highest-energy electrons in potassium (E = 418.8 kJ/mol) and krypton (E = 1350.7 kJ/mol).

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