Which pair would not be expected to form precipitate when solution are mixed?

Options

A. `K^(o+),SO_(4)^(2-)`
B. `Na^(o+),S^(2-)`
C. `Ag^(o+),NO_(3)^(Θ)`
D. `Al^(3+),HO^(Θ)`

Explanation

To determine which pair would not be expected to form a precipitate when solutions are mixed, we need to analyze the solubility of the compounds formed by the ions in each pair. Here’s a step-by-step solution:

### Step 1: Identify the Ions in Each Pair
We have four pairs of ions to consider:
1. K⁺ and SO₄²⁻
2. Na⁺ and S²⁻
3. Ag⁺ and NO₃⁻
4. Al³⁺ and OH⁻

### Step 2: Determine the Compounds Formed
Next, we will write down the compounds formed by the combinations of these ions:
1. K⁺ + SO₄²⁻ → K₂SO₄
2. Na⁺ + S²⁻ → Na₂S
3. Ag⁺ + NO₃⁻ → AgNO₃
4. Al³⁺ + OH⁻ → Al(OH)₃

### Step 3: Check the Solubility of Each Compound
Now, we will check the solubility of each compound in water:
1. K₂SO₄ is soluble in water.
2. Na₂S is soluble in water.
3. AgNO₃ is soluble in water.
4. Al(OH)₃ is not soluble in water and forms a precipitate.

### Step 4: Identify the Pairs That Do Not Form Precipitates
From the analysis:
- K₂SO₄ (K⁺ and SO₄²⁻) is soluble.
- Na₂S (Na⁺ and S²⁻) is soluble.
- AgNO₃ (Ag⁺ and NO₃⁻) is soluble.
- Al(OH)₃ (Al³⁺ and OH⁻) forms a precipitate.

### Conclusion
The pairs that would not be expected to form a precipitate when mixed are:
- K⁺ and SO₄²⁻
- Na⁺ and S²⁻
- Ag⁺ and NO₃⁻

Thus, the answer is that the pairs K⁺ and SO₄²⁻, Na⁺ and S²⁻, and Ag⁺ and NO₃⁻ do not form precipitates.