Question

Using the periodic table, predict the charges of the ions of the following elements: (a) Ga, (b) Sr, (c) As, (d) Br, (e) Se.

Short answer

(a) Ga: \(+3\) charge, (b) Sr: \(+2\) charge, (c) As: \(-3\) charge or covalent compounds with no charge, (d) Br: \(-1\) charge, (e) Se: \(-2\) charge.

Step-by-step solution

Locate the elements on the periodic table

First, find the given elements on the periodic table. The periodic table is organized into periods (horizontal rows) and groups (vertical columns). Elements in the same group generally have the same number of valence electrons, which determine their chemical properties and ion charges. (a) Ga: Gallium is in Group 13 (or 3A) and Period 4. (b) Sr: Strontium is in Group 2 (or 2A) and Period 5. (c) As: Arsenic is in Group 15 (or 5A) and Period 4. (d) Br: Bromine is in Group 17 (or 7A) and Period 4. (e) Se: Selenium is in Group 16 (or 6A) and Period 4.

Determine the common ion charges based on their group

Next, we can determine the most common ion charge based on the group number of the elements. For elements in Groups 1-2 and 13-18: - Group 1 (or 1A) elements usually form ions with a +1 charge. - Group 2 (or 2A) elements usually form ions with a +2 charge. - Group 13 (or 3A) elements usually form ions with a +3 charge. - Group 14 (or 4A) elements typically do not form ions (e.g., carbon), or may form ions with various charges (e.g., silicon). - Group 15 (or 5A) elements usually form ions with a -3 charge or compounds as covalent structures with no charge. - Group 16 (or 6A) elements usually form ions with a -2 charge. - Group 17 (or 7A) elements usually form ions with a -1 charge. - Group 18 (or 8A) elements are the noble gases which are mostly inert and do not usually form ions.

Predict the charges of the given elements

Based on the group numbers and the common ion charges, we can predict the charges of the given elements. (a) Ga: Gallium is in Group 13, so it usually forms ions with a +3 charge. (b) Sr: Strontium is in Group 2, so it usually forms ions with a +2 charge. (c) As: Arsenic is in Group 15, so it usually forms ions with a -3 charge or form covalent compounds with no charge. (d) Br: Bromine is in Group 17, so it usually forms ions with a -1 charge. (e) Se: Selenium is in Group 16, so it usually forms ions with a -2 charge.

Key concepts

Ion Charges

Ion charges refer to the electric charge that an atom acquires when it loses or gains electrons, forming an ion. This happens because atoms want to have a full outer shell of electrons, which provides stability. The number of electrons an atom will gain or lose—and hence the charge of the ion—is determined by its position on the periodic table.

• Elements in groups 1, 2, and 13 usually lose electrons to form positive ions, known as cations.
• For instance, elements in Group 1 lose one electron and form a +1 charge. Those in Group 2 lose two electrons, forming a +2 charge, and in Group 13, they lose three electrons, forming a +3 charge.
• Conversely, elements in Groups 15, 16, and 17 tend to gain electrons to form negative ions, known as anions.
• For example, Group 15 elements typically gain 3 electrons, giving them a -3 charge, Group 16 gains 2 electrons for a -2 charge, and Group 17 gains 1 electron for a -1 charge.

Understanding ion charges is crucial in predicting how different elements will react with each other, forming compounds.

Group Number

The group number on the periodic table plays a vital role in determining an element's characteristics, including its ion charge. Groups are the vertical columns on the periodic table, ranging from Group 1 to Group 18. Elements within a same group have the same number of valence electrons, influencing their chemical behavior.

• Groups 1 and 2, along with Groups 13 to 17, are known for their characteristic ion charges as they readily form ions, discussed in the previous section.
• For example, Group 1 elements are alkali metals that readily lose one electron to achieve a stable configuration, resulting in a +1 ion charge.
• On the other hand, noble gases in Group 18 are renowned for having full valence shells, making them largely unreactive and not forming ions under standard conditions.

This correlation between group number and ion behavior enables easier prediction of chemical reactions and bonding patterns.

Valence Electrons

Valence electrons are the outermost electrons of an atom and are responsible for forming chemical bonds. The number of valence electrons is often determined by the group number for main-group elements on the periodic table. Knowing this helps predict the ion charges and the types of ions an element can form.

• Generally, elements in Group 1 have one valence electron, which they readily lose to form a +1 ionic charge.
• Group 2 elements have two valence electrons and lose both to form a +2 charge.
• In contrast, Group 16 elements have six valence electrons and need to gain two more to fill their valence shell, typically forming a -2 charge.
• This pattern continues across the periodic table, emphasizing the link between valence number and the likelihood of electron gain or loss.

Recognizing valence electrons is fundamental in chemistry as it dictates how atoms will interact and bond with one another.

Chemical Properties

The chemical properties of elements are largely dictated by their valence electrons and ion charges. These properties govern how an element will engage in chemical reactions and the kinds of bonds it will form.

• Elements with similar valence electron counts—and thus in the same group—often demonstrate similar chemical behaviors.
• For example, alkali metals (Group 1) are highly reactive with water due to their single valence electron, which they eagerly lose.
• Halogens (Group 17), with their seven valence electrons, are highly reactive non-metals, often forming salts with metals by gaining an electron.
• Noble gases, however, are characterized by their complete valence shells, making them largely inert as they do not easily form ions or participate in chemical reactions.

Understanding these chemical properties helps in predicting how elements will behave in both isolated conditions and when combined with other substances.