Acid-Base Reactions and Ionic Equations

The video tutorial explains how to write and balance the net ionic equation for the reaction between nitric acid (HNO3) and ammonium carbonate ((NH4)2CO3). It covers the initial double displacement reaction, the decomposition of carbonic acid, and the steps to balance the molecular equation. The tutorial also discusses the solubility of compounds, the formation of the complete ionic equation, and the identification and removal of spectator ions to derive the net ionic equation. The final balanced net ionic equation is presented, ensuring both charge and atom balance.

What type of reaction occurs between nitric acid and ammonium carbonate initially?

• Decomposition

• Synthesis

• Double displacement

• Single displacement

What happens to carbonic acid in the reaction?

• It reacts with nitrate ions

• It remains unchanged

• It decomposes into water and carbon dioxide

• It forms a precipitate

What is the first step in balancing a molecular equation?

• Assigning states to compounds

• Adjusting the coefficients

• Writing the net ionic equation

• Identifying spectator ions

Which of the following is a strong acid and thus soluble in water?

• Nitric acid

• Ammonium carbonate

• Water

• Carbon dioxide

What is the state of carbon dioxide in the reaction?

• Liquid

• Solid

• Gas

• Aqueous

What do you do with spectator ions in a complete ionic equation?

• Double their coefficients

• Convert them to gases

• Cross them out

• Highlight them

Which ions are considered spectator ions in this reaction?

• Water molecules

• Nitrate ions

• Carbonate ions

• Hydrogen ions

What is the net charge on both sides of the net ionic equation?

• Positive

• Variable

• Negative

• Zero

What is the final step in writing a net ionic equation?

• Balancing the atoms

• Writing the states

• Crossing out spectator ions

• Ensuring charge balance

Why is it important to balance the charge in a net ionic equation?

• To maintain electrical neutrality

• To ensure the reaction is spontaneous

• To increase reaction rate

• To decrease reaction rate