Question

Write the Lewis structure for each of the following atoms. a. He \((Z=2)\) b. \(\operatorname{Br}(Z=35)\) c. \(\operatorname{Sr}(Z=38)\) d. \(\operatorname{Ne}(Z=10)\) e. I \((Z=53)\) f. Ra \((Z=88)\)

Short answer

a. He: \[\bullet \quad \bullet\] b. Br: \[\begin{array}{c} \bullet & \bullet & \bullet \\ \mathrm{Br} & \bullet & \bullet \\ \bullet & \bullet & \end{array}\] c. Sr: \[\bullet \quad \bullet\] d. Ne: \[\begin{array}{c} \bullet & \bullet & \bullet \\ \mathrm{Ne} & \bullet & \bullet \\ \bullet & \bullet & \bullet \end{array}\] e. I: \[\begin{array}{c} \bullet & \bullet & \bullet \\ \mathrm{I} & \bullet & \bullet \\ \bullet & \bullet & \end{array}\] f. Ra: \[\bullet \quad \bullet\]

Step-by-step solution

Find the number of valence electrons

For each element, look up its atomic number (Z) and determine the number of valence electrons based on its position in the periodic table. a. He (Z=2): Helium belongs to Group 18 (Noble Gases) and has 2 valence electrons. b. Br (Z=35): Bromine belongs to Group 17 (Halogens) and has 7 valence electrons. c. Sr (Z=38): Strontium belongs to Group 2 (Alkaline Earth Metals) and has 2 valence electrons. d. Ne (Z=10): Neon belongs to Group 18 (Noble Gases) and has 8 valence electrons. e. I (Z=53): Iodine belongs to Group 17 (Halogens) and has 7 valence electrons. f. Ra (Z=88): Radium belongs to Group 2 (Alkaline Earth Metals) and has 2 valence electrons.

Draw the Lewis structures

Draw the symbol of each element and place the valence electrons as dots around it. a. He (2 valence electrons): He: \[\bullet \quad \bullet\] b. Br (7 valence electrons): Br: \[\begin{array}{c} \bullet & \bullet & \bullet \\ \mathrm{Br} & \bullet & \bullet \\ \bullet & \bullet & \end{array}\] c. Sr (2 valence electrons): Sr: \[\bullet \quad \bullet\] d. Ne (8 valence electrons): Ne: \[\begin{array}{c} \bullet & \bullet & \bullet \\ \mathrm{Ne} & \bullet & \bullet \\ \bullet & \bullet & \bullet \end{array}\] e. I (7 valence electrons): I: \[\begin{array}{c} \bullet & \bullet & \bullet \\ \mathrm{I} & \bullet & \bullet \\ \bullet & \bullet & \end{array}\] f. Ra (2 valence electrons): Ra: \[\bullet \quad \bullet\]

Key concepts

Valence Electrons

Understanding valence electrons is crucial for constructing Lewis structures and predicting an element's chemical reactivity. Valence electrons are the outermost electrons of an atom and are primarily responsible for the chemical properties of the element. In simple terms, these electrons can be thought of as the 'hands' of an atom, reaching out and forming bonds with other atoms.

For instance, helium (He) has two valence electrons and is stable; it doesn't tend to form bonds with other atoms. Bromine (Br), on the other hand, with seven valence electrons, seeks to share or gain an electron to achieve a full octet like the noble gases, which represents a particularly stable configuration. This is reflected in the step-by-step solution, where the number of dots around the symbol of each element represents its valence electrons.

The stability of elements and their tendency to form compounds depends significantly on the configuration of these valence electrons. The Lewis structure, a visual representation, not only shows us the number of valence electrons but how atoms might interact with others to form molecules. The dots placed around the elemental symbol give us a clear picture of the available electrons for bonding.

Periodic Table Groups

The periodic table organizes elements into groups that share similar chemical properties, mainly due to the similar valence electron configurations within a group. Group numbers, found at the top of each vertical column, can help us quickly determine the number of valence electrons for elements of the main groups.

For example, all elements in Group 1 have one valence electron and tend to lose that electron to achieve a full outer shell. Group 17 elements, like bromine (Br) and iodine (I), typically have seven valence electrons and often gain one more to complete their octet. Group 18, the noble gases including helium (He) and neon (Ne), are already stable with a complete octet (or duplet, in the case of helium), which makes them relatively nonreactive.

When students are learning to draw Lewis structures, recognizing the group to which an element belongs simplifies predicting the number of valence electrons and, therefore, the structure. The exercise solution demonstrates this with various elements, outlining the predictable pattern of valence electrons according to their group placement.

Atomic Number

The atomic number of an element, often represented by the symbol 'Z', is a fundamental characteristic foundational for understanding the periodic table and its arrangements. It tells us the number of protons found in the nucleus of an atom and, by extension, the number of electrons when the atom is neutral.

The Basics

The periodic table is laid out in order of increasing atomic number. Helium, for instance, has an atomic number of 2, which means it has two protons and, when neutral, two electrons. These two electrons are the valence electrons for helium. For the element bromine, with an atomic number of 35, it has 35 protons and, when neutral, 35 electrons with seven fulfilling the role of valence electrons.

The atomic number essentially guides us in determining how many electrons an element has available for bonding, which is crucial when drawing Lewis structures, as seen in the provided exercise solutions. It's the 'ID number' for elements, allowing us to pinpoint its location on the periodic table and subsequently deduce important chemical properties, such as reactivity and valency.