This cheat sheet covers the core calculations used in general chemistry stoichiometry, solution chemistry, and gas behavior. College chemistry students need these tools to move between balanced equations, measurable laboratory quantities, and chemical interpretations. It is designed to support quick reference during problem solving, lab preparation, and exam review.
The focus is on setting up calculations clearly so units, coefficients, and assumptions stay visible.
Key Facts
- The mole conversion between mass and amount is , where is moles, is mass, and is molar mass.
- Stoichiometric mole ratios come only from coefficients in the balanced equation, such as giving .
- The limiting reactant is the reactant that produces the smaller calculated amount of product, and theoretical yield is based on that reactant.
- Percent yield is calculated by .
- Molarity is , where must be in liters and has units of .
- Dilution problems use because the moles of solute stay constant when only solvent is added.
- The ideal gas law is , with consistent units such as .
- Real gases deviate most from ideal behavior at high pressure and low temperature, and the van der Waals equation is .
Vocabulary
- Stoichiometry
- Stoichiometry is the use of balanced chemical equations to calculate relationships between reactants and products.
- Limiting Reactant
- The limiting reactant is the reactant that is consumed first and determines the maximum amount of product that can form.
- Theoretical Yield
- The theoretical yield is the maximum product amount predicted by stoichiometry from the limiting reactant.
- Molarity
- Molarity is the concentration of a solution defined as moles of solute per liter of solution, .
- Dilution
- Dilution is the process of lowering solution concentration by adding solvent while keeping the moles of solute constant.
- Ideal Gas
- An ideal gas is a model gas whose particles have negligible volume and no intermolecular attractions, so it follows .
Common Mistakes to Avoid
- Using grams directly in mole ratios is wrong because balanced equation coefficients compare moles, not mass. Convert mass with before applying stoichiometric ratios.
- Choosing the reactant with the smaller given mass as the limiting reactant is wrong because different substances have different molar masses and coefficients. Calculate product amount from each reactant to identify the true limiting reactant.
- Using milliliters directly in is wrong because molarity requires volume in liters. Convert with before calculating concentration.
- Applying to reactions is wrong when moles change because the equation only describes dilution of the same solute. For chemical reactions, use balanced equation mole ratios.
- Using Celsius temperature in gas law calculations is wrong because requires absolute temperature. Convert with .
Practice Questions
- 1 For , how many moles of form from if excess aluminum is present?
- 2 What volume of solution contains of ?
- 3 A gas sample has , , and . How many moles are present using ?
- 4 Why do real gases deviate more from ideal behavior at high pressure and low temperature, and which assumptions of the ideal gas model fail?
Understanding General Chemistry Stoichiometry and Solutions
Stoichiometry works because a balanced chemical equation is a conservation statement. Atoms are not created or destroyed during an ordinary reaction, so the coefficients show the fixed particle scale of the reaction. Those coefficients describe numbers of molecules, formula units, or moles.
They do not describe grams. This distinction prevents a common error. A student must first change a measured mass into an amount in moles before using the reaction relationship.
Unit labels are useful checkpoints during every step. If grams are supposed to disappear, they must cancel before the final answer. If a unit remains that does not fit the requested quantity, the setup needs correction.
Limiting reactant problems model a real mixture in which one substance runs out first. The other reactant is left over and is called the excess reactant. A reliable method is to treat each starting reactant as if it were the one controlling the reaction, then calculate the product possible from each.
The smaller product amount identifies the reactant that stops the reaction. This idea matters in laboratory work because adding an excess of one cheaper or safer reagent can help use up a valuable reagent.
The measured yield can be lower than expected when material sticks to glassware, a reaction is incomplete, side reactions occur, or product is lost during filtering and transfer. A yield above one hundred percent often signals wet product or contamination rather than an unusually successful reaction.
Solution calculations connect an amount of dissolved substance to a measured volume of liquid. The concentration tells how much solute is present in each liter of the final solution, not how much solvent was first placed in the container. In careful preparation, a solid is dissolved, transferred to a volumetric flask, then water is added until the liquid reaches the calibration mark.
Reading the bottom of the meniscus at eye level helps avoid volume error. Concentration matters because reaction rates, electrical conductivity, acidity, and biological effects often depend on the number of dissolved particles. Ionic compounds can separate into ions in water, so one mole of a salt may create more than one mole of dissolved particles.
During dilution, only adding solvent preserves the solute amount. A reaction, evaporation, or spill means the simple dilution relationship no longer applies.
Gas problems require attention to temperature scale and pressure conditions. Kelvin is used because it begins at absolute zero, where the simple particle motion model reaches its lower limit. Gas pressure comes from collisions between particles and the walls of a container.
Higher temperature means faster motion and stronger or more frequent collisions. The ideal gas model assumes particles have negligible size and no attractions. That works well for many gases at ordinary conditions.
At low temperature, particles move slowly enough for attractions to matter. At high pressure, their own volume becomes important.
Real gas corrections account for these effects. When a gas is collected over water, the measured pressure includes water vapor, so the water vapor contribution must be removed before finding the pressure of the dry gas.