WRITING MOLECULAR, IONIC AND NET IONIC EQUATIONS

John Dalton picture


John Dalton
English scientist

One of Dalton's hypotheses about matter was that the smallest "unit" (the atom) cannot be subdivided, created or destroyed . . . this was based on an existing concept which is now known as the . . . . it is the basis for balancing chemical equations and performing much of the math in chemistry.  In fact, the word "atom" comes from the Greek atomos meaning indivisible and uncuttable. It is this law that affords chemists and chemistry students the exhilarating pleasure of balancing chemical equations.

Simply make the mass of the . . . . compounds formed during a reaction. Products are located to the right of the reaction arrow (→) equal to the mass of the . . . . compounds broken apart during a reaction. Reactants are located to the left of the reaction arrow (→) and the equation will be balanced. In practice, it is actually easier to make the type and number of reactant atoms equal to the type and number of the product atoms. To do this you will need to accurately

  1. sum the number of atoms present in a chemical formula
  2. write a chemical formula from a compound's name
  3. predict the reaction products if they are not given

Writing Molecular Equations

While balancing chemical equations is mostly "trial and error", the following steps will reduce the number of "errors" you will encounter.

  1. If the chemical equation is given in "word form", convert the chemical names to formulas - this must be done accurately, or the equation cannot be balanced.
  2. Never change a compound's formula to balance an equation . . . . only a reactant's (or product's) . . . . the number written to the left of a formula that balances a chemical equation. is changed.

    For example, in the following decomposition reaction of water, there are 2 Oxygen atoms on the product side but only 1 Oxygen atom on the reactant side . . . .

    H2O → H2 + O2

    ✗  H2O2 → H2 + O2
    ✓  2 H2O → 2 H2 + O2

    The ✗incorrect answer is a balanced equation, but adding the "2" as a subscript makes the reactant hydrogen peroxide (not water). So, only the ✓ correct answer is the balanced equation for the decomposition of water.


  3. Identify the most complex substance and choose an element in it that appears in only one other compound. Adjust the coefficients to obtain the same number of atoms of this element on both sides.
  4. Balance . . . . ions containing more than one element . . . NH4+, C2H3O2-, MnO4-, CO32-, CrO42-, SO42-, PO43-, etc. (if present on both sides of the chemical equation) as a unit.
  5. If a reactant or product is an element, balance it last.
  6. Balance the remaining atoms . . . end with the least complex substance . . . use fractional coefficients if necessary.
  7. Convert fractional coefficients to a whole number by multiplying ALL coefficients by the denominator of the fraction.
  8. Check your work by counting the atoms of each type on both sides of the equation.

Write the   balanced molecular equation   for the reaction between Acetic acid and Sodium hydroxide.

 HC2H3O2 (aq) + NaOH(aq) → H2O(l) + NaC2H3O2 (aq) 

Note that if a coefficient is "1" it is not written. Follow the directions below to write the ionic and net ionic equations for this reaction.

Writing Ionic Equations

The    ionic equation   breaks some aqueous compounds (they have an (aq) after the formula) into their ions as shown below:

 HC2H3O2 (aq) + Na+(aq) + OH–(aq) → H2O(l) + Na+(aq) + C2H3O2–(aq) 

Note that acetic acid was not separated into its ions while sodium hydroxide and sodium acetate were written as separated ions. How do you know when to write ionic compounds as separated ions? Use the steps below to guide you:

  1. Break apart into ions any strong acid, strong base or soluble salt that has an (aq) after its formula.
    • NaOH is a strong base, so break it apart into   Na+(aq) + OH–(aq) 
    • NaC2H3O2 is a soluble salt, so break it apart into   Na+(aq) + C2H3O2–(aq) 
  2. Place the coefficient (from the balanced equation) TIMES the ion's subscript in front of the ion as its new coefficient. In the equation above, all coefficients and ionic subscripts are "1", so all the coefficients in the ionic equation are also "1". Below are examples of ionic compounds from a molecular equation where their coefficients are greater than "1":
    • 2 Fe2(SO4)3 (aq) from a molecular equation is written as 4 Fe3+(aq) + 6 SO42-(aq)
    • Hg2(NO3)2 (aq) from a molecular equation is written as Hg22+(aq) + 2 NO3-(aq)

      • You must know the common cations and ions and only "separate" ionic compounds into a common cation / anion. For the ionic compound Hg2(NO3)2 (aq), how did we know the cation was Hg22+(aq) and not 2 Hg2+(aq)?

      Since we know nitrate has a -1 charge and there are 2 nitrates in Hg2(NO3)2 (aq), the mercury ion must have a +2 charge. Hg22+(aq) has a +2 charge while 2 Hg2+(aq) has an overall +4 charge.

  3. Do Not break apart any reactant or product that
    • has (s) or (l) or (g) . . . . water   H2O(l)   is not broken apart.
    • has (aq) and is a weak acid or base . . . . acetic acid   HC2H3O2 (aq)   is not broken apart since it is a weak acid.
    • has (aq) and is a . . . . a substance that doesn't produce ions when dissolved in water.
      If the compound can't be separated into two ions (a cation and an anion) found in the list of common ions, then it is a non-electrolyte.      . Examples not found in the ionic equation above, include . . . .
      • ethanol, CH3CH2OH (aq), is not broken apart
      • glucose, C6H12O6 (aq), is not broken apart
  4. Make sure to include the charge and state for the ions and compounds - all compounds have a charge of zero.
  5. Check that the ionic equation is balanced.
  6. As a final check, the total charge on the reactant side must equal the total charge on the product side

Writing Net Ionic Equations

The   net ionic equation   is a simplified version of the ionic equation where . . . . ions that are present in exactly the same quantity and state on both the reactant and product side of an ionic equation. are removed.

HC2H3O2 (aq) + Na+(aq) + OH–(aq) → H2O(l) + Na+(aq) + C2H3O2–(aq)  (ionic equation)

 HC2H3O2 (aq) + OH–(aq) → H2O(l) + C2H3O2–(aq) 


What's the purpose of net ionic equations? They show the "players" in a chemical reaction and help chemists organize reactions. For example, the net ionic equation between any strong acid (i.e. hydrochloric acid) and any strong base (i.e. potassium hydroxide) is exactly the same . . . .

H+(aq) + Cl–(aq) + K+(aq) + OH–(aq) → H2O(l) + K+(aq) + Cl–(aq)  (ionic equation)

 H+(aq) + OH–(aq) → H2O(l) 



    Acid     +     Base       →       Salt       +   Water

H2SO4(aq) + 2 CsOH(aq) → Cs2SO4(aq) + 2 H2O(l)

2 H+(aq) + SO42–(aq) + 2 Cs+(aq) + 2 OH–(aq) → 2 Cs+(aq) + SO42–(aq) + 2 H2O(l)

2 H+(aq) + 2 OH–(aq) → 2 H2O(l)

The coefficients in a net ionic equation are
the smallest whole number ratio . . . .

 H+(aq) + OH–(aq) → H2O(l) 




Any of the acids (below) will react with any of the bases (below) in an acid‑base reaction to form a salt + water with the same net ionic equation. Sulfuric acid and cesium hydroxide provide a representative example.

Acid: HF, HCl, HBr, HI, HC2H3O2, HNO3, H2SO4, H3PO4
Base: LiOH, NaOH, KOH, CsOH, RbOH, Mg(OH)2, Ca(OH)2, Ba(OH)2, Sr(OH)2, Al(OH)3



With just the acids and bases listed, there are 8 × 10 (80) acid-base reactions that share the same   net ionic equation   .