Problem 69 - Chemistry Central Science

Question

Compare the elements sodium and magnesium with respect to the following properties: (a) electron configuration, (b) most common ionic charge, (c) first ionization energy, (d) reactivity toward water, (e) atomic radius. Account for the differences between the two elements.

Step-by-Step Solution

Answer

Sodium (Na) and Magnesium (Mg) have different electron configurations, with Na being 1s² 2s² 2p⁶ 3s¹ and Mg being 1s² 2s² 2p⁶ 3s². This leads to different ionic charges: +1 for Na and +2 for Mg. Magnesium has a higher first ionization energy (737.7 kJ/mol) than sodium (495.8 kJ/mol), indicating a greater difficulty in removing an outer electron. Sodium is more reactive with water than magnesium, with sodium reacting vigorously and magnesium reacting slowly. Lastly, sodium has a larger atomic radius (186 pm) than magnesium (160 pm) due to an increase in electron-electron repulsion as the number of electrons increases across a period. These differences are a result of their positions in the periodic table and electron distributions in energy levels and orbitals.

Step 1: Identify the electron configuration of sodium and magnesium.

Sodium (Na) has an atomic number of 11, and its electron configuration is 1s² 2s² 2p⁶ 3s¹. Magnesium (Mg) has an atomic number of 12, and its electron configuration is 1s² 2s² 2p⁶ 3s².

Step 2: Identify the most common ionic charge for sodium and magnesium.

Sodium typically loses one electron to achieve a full outer electron shell, resulting in a +1 ionic charge (Na⁺). Magnesium typically loses two electrons to achieve a full outer shell, resulting in a +2 ionic charge (Mg²⁺).

Step 3: Compare the first ionization energy of sodium and magnesium.

The first ionization energy is the energy required to remove the outermost electron from a neutral atom. Sodium's first ionization energy is 495.8 kJ/mol, whereas magnesium's first ionization energy is 737.7 kJ/mol. This indicates that it is more difficult to remove an electron from a magnesium atom than from a sodium atom.

Step 4: Compare the reactivity of sodium and magnesium towards water.

Both sodium and magnesium are very reactive toward water; however, sodium is more reactive than magnesium. Sodium reacts vigorously with water to form sodium hydroxide (NaOH) and hydrogen gas (H₂), while magnesium reacts more slowly, forming magnesium hydroxide (Mg(OH)₂) and hydrogen gas.

Step 5: Compare the atomic radius of sodium and magnesium.

In general, atomic radius decreases across a period and increases down a group in the periodic table. Sodium has a larger atomic radius (186 pm) than magnesium (160 pm). This is because the number of electrons increases across a period, leading to greater electron-electron repulsion which causes the atomic size to decrease.

Step 6: Explain the differences between the two elements.

Sodium and magnesium have different electron configurations, resulting in different ion charges, reactivities, and atomic radii. Sodium has a lower first ionization energy and a larger atomic radius than magnesium, resulting in a higher reactivity with water. Furthermore, magnesium's +2 ionic charge occurs due to the loss of two valence electrons, while sodium's +1 ionic charge is due to the loss of only one valence electron. These differences can be attributed to their positions in the periodic table and the distribution of their electrons among their energy levels and orbitals.

Key Concepts

Electron Configuration

When comparing the electron configurations of sodium (Na) and magnesium (Mg), it's essential to understand how electrons are arranged around the nucleus of these atoms. Sodium, with its atomic number of 11, has an electron configuration of 1s² 2s² 2p⁶ 3s¹, which means its outermost shell contains just one electron. In contrast, magnesium, with an atomic number of 12, has its electrons arranged as 1s² 2s² 2p⁶ 3s², with two electrons occupying its outer shell.

The differing numbers of electrons in the outermost shell influence the chemical behavior of these elements. This is because elements tend to gain or lose electrons to reach a full outer shell, resembling the nearest noble gas configuration. For sodium, shedding its one outer electron is a relatively simple process which is reflected in its chemical properties. Magnesium, on the other hand, requires more effort to remove two outer electrons and this further impacts how it interacts with other elements.

Insight into Valence Electrons

Recognizing that the outermost electrons are called valence electrons, we understand why sodium has a valence of one and magnesium a valence of two. These valence electrons are critical in determining how elements bond and react with others.

Ionic Charge

Focusing on the most common ionic charges of sodium and magnesium gives us insight into their chemical stability. Sodium, due to its single valence electron, readily loses that electron in chemical reactions, leading to a common ionic charge of +1 (Na⁺). This makes sodium a cation, an ion with a positive charge, eager to combine with non-metals or other negatively charged ions.

In the case of magnesium, it sheds both of its valence electrons to achieve a stable configuration resulting in a +2 (Mg²⁺) ionic charge. This confers magnesium a higher charge compared to sodium, making it more capable of forming ionic bonds with a wider variety of ions, as it can balance out two negative charges.

Equilibrium and Stability

Understanding the concept of electrical neutrality is crucial. Atoms become ions to ensure electrical neutrality through a full outer electron shell, thus achieving a more stable, low-energy state which is often found in noble gases.

First Ionization Energy

The first ionization energy signifies the energy required to remove the outermost electron from a neutral atom in its gaseous state. With sodium's first ionization energy at 495.8 kJ/mol and magnesium's significantly higher at 737.7 kJ/mol, this points to the greater difficulty in removing an electron from magnesium.

The discrepancy is largely due to the effective nuclear charge, which is greater in magnesium due to its additional proton. This means that sodium's outer electron is less attracted to the nucleus and thus easier to remove compared to magnesium's.

Nuclear Attraction

The larger the atomic number, the greater the number of protons in the nucleus, and consequently, the stronger the pull on the electrons. Sodium has one less proton than magnesium, so therefore, it has a lower first ionization energy.

Reactivity Toward Water

Reactivity toward water is a vivid indicator of an element's chemical activity. Sodium is infamous for its vigorous reaction with water, creating sodium hydroxide (NaOH) and hydrogen gas (H₂). Magnesium's reaction with water is notably slower and less exothermic, producing magnesium hydroxide (Mg(OH)₂) along with hydrogen gas.

The rate of reaction is influenced by the energy needed to displace hydrogen from water and the stability of the metal hydroxide formed. The lower ionization energy of sodium makes it more reactive, as it can more easily lose its outer electron to form a strong base, whereas magnesium requires extra energy to remove two electrons, resulting in a more gradual reaction.

Chemical Kinetics

Not only do the inherent properties of the elements affect their reaction rate, but so do surface factors, such as state of division and the presence of coatings on the metal that might impede the reaction.

Atomic Radius

The atomic radius is a measure of the size of an atom from its nucleus to the boundary of its electron cloud. Sodium's larger atomic radius of 186 pm compared to magnesium's 160 pm is attributed to its fewer number of electrons and the consequent reduced electrostatic pull by its nucleus.

The difference highlights the trend within the periodic table where, moving from left to right across a period, the atomic radius decreases due to increased nuclear charge pulling electrons closer. As sodium is placed just before magnesium in the periodic table, it experiences less nuclear pull, resulting in a larger radius.

Periodic Trends

The periodic table's structure allows for predictions about an element's size. As you move across a period, the effective nuclear charge increases, pulling electrons inwards and thus decreasing the atomic radius.