Unit 6 [B] Section 4A

Solubility and Precipitation

Reviewing double replacement reactions

The ______from two reactants replace each other.

General format: AX + BY  AY + BX

Ionic compounds are neutral compounds that are solids at room temperature. In order to be electrolyte & conduct electricity, they must be soluble in water. They dissolve to produce cations and anions.

A ______reaction is when 2 soluble substances are reacted and they form an ______substance called a precipitate.

Precipitation reactions are a type of double replacement reactions.

Must use solubility rules chart to decide if the new substance produced is insoluble.

The chart is based on anions. Watch out for exceptions!

Example:

Decide whether each is soluble. S for soluble and I for insoluble(precipitate)

[1]NaNO3

[2]CaBr2

[3]Fe(C2H3O2)2

[4](OH)2

[5]Cu(OH)2

Example:

Predict the products of the reaction and then identify the precipitate.

AgNO3 (aq) + NaCl(aq) ______+ ______

You Try:

Predict the products and identify the precipitate.

BaCl2 + K2CO3 ______+ ______

Unit 6 [B] Section 4B

Net Ionics of Precipitation Reactions

Writing Net Ionic Equations

To show the details of aqueous reactions that involve ions in aqueous solutions, we use ionic equations.

  • ______Equation: the typical chemical equation you are used to writing, keeping all molecules together
  • ______Equation: shows all the particles in a solution as they really exist, as ions or molecules

Anything aqueous needs to be split apart as cations & anions.

Any solid substance should stay intact.

Coefficients need to be multiplies by subscripts to determine the exact amount of cations and anions

  • A ______Ion: an ion that is not participating in the reaction; you can identify it because it is found on both the reactant and product side of an equation in the same amounts.
  • ______Equation: the final equation showing the major players. All spectator ions are removed.

Examples:

1 / Molecular Equation: / KI(aq) + AgNO3(aq)  KNO3(aq) + AgI(s)
Complete Ionic:
Net Ionic:
2 / Molecular Equation: / 2NaOH(aq) + CuCl2(aq)  2NaCl(aq) + Cu(OH)2(s)
Complete Ionic:
Net Ionic:
3 / Molecular Equation: / FeCl3(aq) + 3NaOH(aq)  Fe(OH)3(s)+ 3NaCl(aq)
Complete Ionic:
Net Ionic:

Unit 6 [B] Section 5

Stoichiometry

The ______of the balanced equation tell how many moles of each substance is used in the reaction.

A Mole ratio is a ______that relates 2 substances in moles. You must use a balanced chemical equation to create it.

What are all the possible mole ratios of:

2 H2 + O2 2 H2O

Stoichiometry: Calculations Using the mole ratio from the ______& information about one compound in the reaction to determine information about another compound in the equation

Example:

What is the mole ratio of chlorine to sodium in the following reaction?

2Na + Cl2 2NaCl

Unit 6 [B] Section 5A:

General Stoichiometry

Mole-Mole : 1 step problem

1. If 4.2 moles of H2reacts completely with O2, how many moles of O2 are needed?

2 H2 + O2 2 H2O

2. Given the BALANCED EQUATION: 2KNO3 2KNO2+ O2, how many moles of oxygen are produced by the decomposition of 0.67 moles of potassium nitrate, KNO3?

Moles and Mass : 2 step Problem

We can’t measure moles in the lab. We can only measure grams.

Molar Mass (grams) = 1 mole of a compound

Mole-Mass (2 step problem)

1. How many grams of AgCl will be precipitated if 0.45 mole AgNO3 is reacted as follows:

2 AgNO3 + CaCl22 AgCl + Ca(NO3)2

2. Given the BALANCED EQUATION:N2 + 3H2 2NH3

How many moles of ammonia, NH3 are produced from 4.42 g of hydrogen gas, H2?

Stoichiometry & Gases:

Recall: Molar Volume of a Gas– at STP, 1 mole of any gas = 22.4 liters

  1. If you need to react 1.5 g of zinc completely, what volume of hydrogen gas will be produced at STP?

2 HCl (aq) + Zn (s) ZnCl2 (aq) + H2 (g)

  1. Given the balanced equation:C3H8 + 5O2 → 3CO2 + 4H2O

How many moles of water will be produced from the complete combustion of 7.3 L of oxygen gas? Assume STP

Keeping all these MOLE equalities straight!

To Go Between… / Use the Equality…
Particles and Moles / 1 mole = 6.02 x 1023 atoms, molecules or FU
Grams and Moles / 1 mole = molar mass (grams)
Moles and liters of a gas at STP / 1 mole =22.4 liters at STP
2 different chemicals in a reaction / Coefficient ratio (mole ratio) from balanced equation

Let’s Practice:

  1. Given the UNBALANCED EQUATION: __MgCO3  __MgO + __CO2, how many L of carbon dioxide gas (CO2)are produced from the reaction of 15 grams of MgCO3? Assume STP and balance the equation.

Unit 6 [B] 5 B:

Percent Yield

  • A “Yield” is a product
  • An “Actual Yield” (A): the actual amount of product produced in the lab
  • A “Theoretical Yield” (T) : the amount of product you should produce if nothing went wrong; use the balanced chemical equation to calculate it
  • A “Percent Yield”: ratio of the actual yield compared to the theoretical yield

% Yield = /

Let’s Practice:

  1. If 4.20 moles of H2 reacts completely with O2, how many grams of H2O are produced (otherwise known as the Theoretical Yield)?

2H2 + O2 2 H2O

Theoretical yield of water = ______g

a)What is the percent yield if 60.0 grams of H2O are actually produced?

  1. You have precipitated 8.5 g of Ba(OH)2. If you started with 4.57 grams NaOH, what is the percent yield?

2 NaOH + BaCl2 Ba(OH)2 + 2 NaCl

Hint:
Identify the actual yield.
Calculate the theoretical yield of Ba(OH)2 using the reactant amount

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