What will happen if aqueous solutions of copper II nitrate and sodium sulfate are mixed together?

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1-Does a reaction occur when aqueous solutions of calcium chloride and lead(II) nitrate are combined? If a reaction does occur, write the net ionic equation. Does a reaction occur when aqueous solutions of potassium sulfide and ammonium nitrate are combined? If a reaction does occur, write the net ionic equation. Does a reaction occur when aqueous solutions of potassium carbonate and iron(II) nitrate are combined? If a reaction does occur, write the net ionic equation. Does a reaction occur when aqueous solutions of iron(II) nitrate and sodium sulfate are combined? If a reaction does occur, write the net ionic equation. Does a reaction occur when aqueous solutions of sodium carbonate and lead(II) nitrate are combined? If a reaction does occur, write the net ionic equation. Does a reaction occur when aqueous solutions of sodium carbonate and lead(II) nitrate are combined? If a reaction does occur, write the net ionic equation.

Having knowledge of precipitation reactions to predict when they will occur is a powerful tool in extraction of metals and salts from solution. If you know what ions are present, then you can design a series of precipitation reactions to extract the specific one that you want. For an example of this in action click to see the extraction and quantification of sulfate in fertiliser experiment.

Let's examine a precipitation reaction that occurs within aqueous solution. A solution is formed when a solute (in this case a salt, like CuSO₄) is dissolved in a solvent (water). The water molecules are small, and have polar regions which have a slight positive and negative charge on either side which enables the molecule to move between and surround the ions, stick to their charges, and separate them (Taylor, 2007, p.294-295).

Salts are ionic compounds; compounds that are made out of Cations (positively charged ions), and Anions (negatively charged ions). When the water molecules move in between the ionic bonds, these cations and anions separate from each other. This process is known as dissociation (Taylor, 2007, p.295).

CuSO₄ (s)      +       H₂O (l)      →        Cu²⁺ (aq)      +       SO₄²⁻ (aq)       
Copper Sulfate (s)   +    Water (l)    →   Copper (II) ion (aq)   +   Sulfate ion (aq)

This means, that when any combination of salt solutions are mixed, these compounds exist in their dissociated form as ions.

NaCl (aq)   +   CuSO₄ (aq)    →    Na⁺ (aq) + Cl⁻ (aq)    +    Cu²⁺ (aq) + SO₄²⁻ (aq)
Sodium Chloride + Copper Sulfate  →  Sodium ion + Chloride ion + Copper (II) ion + Sulfate ion

Ions remain in solution (meaning, dissolved in water) because the attraction forces between the ions and the water is greater than the attraction forces between the ions and other ions. Conversely, ions will come out of solution (forming a precipitate) if the attraction forces are greater between the ions and other ions. This is like a situation where your family (mum, dad, brothers and sisters) may dissociate from each other and mingle when at a family barbecue - whether you choose to stick together or not depends on how much you like to hang out with the other people there.But how can we check whether the ions present in our mixture of salts like to hang out with each other more than with the solvent? Tables of data have been collected from many experiments of all sorts of combinations of ions such that we can simply look them up to determine whether any of the ions present will produce a solid. These are known as solubility tables of ions.

Figure 1: Solubility Rules (Linstead, 2011, p.152)

Let's look at our example mixture earlier, NaCl + CuSO₄, but this time the physical states (solid, liquid, gas, aqueous) have been removed temporarily from the products until we can be certain of their solubility. NaCl (aq)   +   CuSO₄ (aq)    →    Na⁺   +   Cl⁻    +    Cu²⁺  +   SO₄²⁻ Like charges repel, so there will not be any chemical reactions occurring with the cations with other cations, and likewise the anions with other anions. On the other hand, opposite charges attract, so our possible products could be any of the following: NaCl       - we can rule this one out because it's the same as our reactant Na₂SO₄   (Sodium Sulfate) CuSO₄   - we can rule this one out too because it's the same as our reactant CuCl₂      (Copper Chloride) Having narrowed down the possibilities, let's look these up in the table. Sodium Sulfate contains the sulfate ion and looking across the row on the table from figure 1, sulfates are soluble with anything except when combined with calcium, strontium, barium and lead cations. So Sodium Sulfate will remain in solution.

Copper chloride, the other possibility, we read the row on the table for chlorides and find that they are soluble in anything except when combined with silver, lead, mercury, or copper cations. So Copper chloride will also remain in solution.

Now that all possibilities have been ascertained, there will be no precipitates formed, and hence no reaction would occur to change their chemical composition. So our original example from above was correct in stating that all the ions are aqueous.

NaCl (aq)   +   CuSO₄ (aq)    →    Na⁺ (aq)   +   Cl⁻ (aq)   +   Cu²⁺ (aq)   +   SO₄²⁻ (aq)

Let's try another example;
Determine the products, if any, of the mixture between Sodium Chloride (NaCl) and Silver Nitrate (AgNO₃) dissolved in water. From the question we know that these two ionic compounds are dissolved in water (another way of telling is if they describe these as 'solutions'), and so we can write their ions in their dissociated form:

NaCl (aq) + AgNO₃ (aq) → Na⁺(aq) + Cl⁻(aq) + Ag⁺(aq) + NO₃⁻(aq)

First we determine what are the possible combinations of these cations and anions could be:

• NaCl         same as reactant (which is aqueous, so we know it is soluble)
• NaNO₃
• AgNO₃     same as reactant (which is aqueous, so we know it is soluble)
• AgCl

1. According to figure 1 all nitrates are soluble so therefore NaNO₃ is soluble; we write this as NaNO₃ (aq), or as their dissociated form Na⁺ (aq) + NO₃- (aq).
   
2. According to figure 1, chlorides (Cl⁻) have low solubility when combined with Silver, Lead, Mercury, and Copper cations, therefore AgCl will form a precipitate. We write this as AgCl (s) to show that it has formed a solid.

NaCl (aq) + AgNO₃ (aq) → NaNO₃ (aq) + AgCl (s)
where AgCl is the precipitate!

At this point, you may have noticed that the products look very similar to the reactants, only that the anions have swapped partners. Sodium (Na⁺) has swapped its Chloride (Cl⁻) partner for a Nitrate (NO₃⁻) partner instead, and similarly the Silver has swapped the Nitrate partner for a Chloride ion. This observation can save you some time because if there is any reaction occurring at all, then the anions will have exchanged places - simply look these products up on the solubility table and determine whether they are soluble or not. 

1. A common error that I see when students write these equations down, is they forget to assign the physical states of matter to each formula that appears in the equation (a formula is things like NaCl, H₂O, Na⁺). This is especially important for precipitation reactions where there is likely to be one substance that has changed from an aqueous phase to a solid phase and this needs to be clearly labelled. For example;

• Mg(NO₃)₂ + 2NaOH → Mg(OH)₂ + 2NaNO₃ is wrong ✗
  
• Mg(NO₃)₂ (aq) + 2NaOH (aq) → Mg(OH)₂ (s) + 2NaNO₃ (aq) is right ✓
  

• Sodium Sulfate is not NaSO₄, because each sodium ion carries 1+ charge, and each sulfate ion carries 2⁻ charge. 
• There needs to be twice as many sodium ions as sulfate ions to balance this out, so the correct formula for sodium sulfate is Na₂SO₄.
  

• Mg(NO₃)₂ (aq) + NaOH (aq) → Mg(OH)₂ (s) + NaNO₃ (aq) is wrong ✗
  
• Mg(NO₃)₂ (aq) + 2NaOH (aq) → Mg(OH)₂ (s) + 2NaNO₃ (aq) is right because now it is balanced. ✓
  

Most chemistry textbooks will have practice problems like this so check those first. If you would like to read some alternative explanations as well as  practice problems then visit some of these websites:

Linstead, G, 2011. Pearson Science 10 Student Book. 1st ed. Melbourne: Pearson Australia.
Taylor, N, 2007. StudyON Chemistry 1. 1st ed. Brisbane: John Wiley & Sons Australia Ltd.