Lewis Dot Structures: Acids: Videos & Practice Problems

Acids are typically covalent compounds composed of hydrogen ions (hydrogen or hydronium ions) bonded to nonmetals. Common examples include hydrobromic acid (HBr), nitric acid (HNO₃), hydrochloric acid (HCl), sulfuric acid (H₂SO₄), and phosphoric acid (H₃PO₄). Acetic acid (CH₃COOH) is an exception, as it can be represented with a hydrogen ion at the beginning or within the molecular formula. When dissolved in solution, acids dissociate into H⁺ ions and anions, such as HBr dissociating into H⁺ and Br^-. This property is crucial for understanding acid behavior in various chemical reactions and acid–base titrations.

Lewis Dot Structures: Acids

Now we're going to say here, generally, an acid is a covalent compound, meaning it's only connected to nonmetals, and it begins with a hydrogen ion called the hydronium ion. Common acids. We have HBr, HNO3, HCl, 4H2SO4, and H3PO4. These are some common types of acids. All of them are covalent. All of them begin with hydrogen.

Now how could it be? Chemistry? Without some exceptions, of course. So acetic acid is an exception. This form shows it as covalent and hydrogen at the beginning. But acetic acid can also be written as CH3COOH. This is also its acidic form.

Now we're going to say acids. We're going to say they can break up when they're dissolved into H. So we're going to say that they are soluble covalent compounds that dissolve into H and an ion when placed in solution. So HBr here can break up into H and Br-. So just remember these key things about acids.

Lewis Dot Structures: Acids Example 1

Here it says to draw the Lewis dot structure for the nitric acid molecule. So HNL₃ we're going to step zero, we're going to break up the acid into the H ion and an anion. Now we're going to say here that for the anion, follow steps one to six to draw its Lewis dot structure. All right, so nitric acid is H⁺ and then NO₃^-, NO₃^- is our anion.

Nitrogen is in Group 5A, so it has five valence electrons. Oxygen has six because it's in Group 6A, so it's going to be 5 + 18 -. 1 means we've gained an electron from an outside source, so that's 24 valence electrons. Let's draw it. Nitrogen goes in the center, it's going to form single bonds with the oxygens. Now remember we need to follow the octet rule for the surrounding elements, so we're adding electrons from.

Now each oxygen has 8 electrons around them, but we've run into an issue with the nitrogen. It only has two 4⁶ electrons around itself. To get it to also follow the octet rule, what do we do? Well, we're going to erase one of these lone pairs and use it to make a double bond. Because nitrate ion has a negative charge, we're going to place it in brackets and the charge on the outside. So this is the drawing of the nitrate ion.

Step 7 says add the H ion to the surrounding elements with a form onto the surrounding elements with a formal charge of -1. Now here, who's gonna have a formal charge of -1? It'll be one of the oxygens that has a formal charge of -1. So here remember formal charge equals group number minus the bonds the element is making plus each individual non bonding electron. So if we looked at this oxygen here, it's 6 -, 1 bond plus six, which equals -1. So what I'm going to do here is I'm going to erase one of these lone pairs here and use it to connect to the hydrogen. So here this here represents the nitric acid molecule and would be the correct representation for this particular.

Determine the Lewis Dot Structure for the following acid:H₂SO₄

• A.
• B.
• C.
• D.

Answer: (C)

Determine the Lewis Dot Structure for hydrocyanic acid:HCN

• A.
• B.
• C.
• D.

Answer: (D)

Determine the Lewis Dot Structure for chloric acid:HClO₃

• A.
• B.
• C.
• D.

Answer: (B)

Determine the Lewis Dot Structure for phosophoric acid:H₃PO₄

• A.
• B.
• C.
• D.

Answer: (D)

Here’s what students ask on this topic:

What are common examples of acids and their chemical formulas?

Common examples of acids and their chemical formulas include:

• Hydrobromic acid: HBr
• Nitric acid: HNO₃
• Hydrochloric acid: HCl
• Sulfuric acid: H₂SO₄
• Phosphoric acid: H₃PO₄
• Acetic acid: CH₃COOH

These acids are covalent compounds composed of hydrogen ions bonded to nonmetals. Acetic acid is an exception as it can be represented with a hydrogen ion at the beginning or within the molecular formula.

How do acids dissociate in solution?

When acids are dissolved in solution, they dissociate into hydrogen ions (H⁺) and anions. For example, hydrobromic acid (HBr) dissociates into H⁺ and Br^- ions:

$H_2+{\mathrm{Br}}_{\mathrm{-}}$

This property is crucial for understanding acid behavior in various chemical reactions and acid–base titrations.

What is the difference between covalent and ionic compounds in the context of acids?

In the context of acids, covalent compounds are composed of nonmetals bonded together, and they typically begin with a hydrogen ion (H⁺). Examples include HBr, HNO₃, and H₂SO₄. These acids dissociate into H⁺ ions and anions when dissolved in solution. Ionic compounds, on the other hand, consist of metals and nonmetals bonded together, and they dissociate into cations and anions in solution. Acids are generally covalent compounds, but they exhibit ionic behavior when dissolved in water.

Why is acetic acid considered an exception among common acids?

Acetic acid (CH₃COOH) is considered an exception among common acids because it can be represented in two forms: with a hydrogen ion at the beginning (HCH₃COO) or within the molecular formula (CH₃COOH). This dual representation is unique compared to other common acids, which typically have the hydrogen ion at the beginning of their formulas. Despite this difference, acetic acid still dissociates into H⁺ ions and anions when dissolved in solution, similar to other acids.

What is the significance of hydrogen ions (H⁺) in acids?

Hydrogen ions (H⁺) are significant in acids because they are responsible for the acidic properties of these compounds. When acids dissolve in water, they release H⁺ ions, which contribute to the acidity of the solution. The concentration of H⁺ ions determines the pH level of the solution, with higher concentrations resulting in lower pH values (more acidic). Understanding the behavior of H⁺ ions is crucial for studying acid–base reactions, titrations, and the overall behavior of acids in various chemical processes.