Conservation of Mass and Thermal Decomposition

Here's a fundamental rule that never breaks: no atoms are lost or made during chemical reactions. This means the mass of your starting materials (reactants) always equals the mass of what you end up with (products).

But sometimes your scales might suggest otherwise! If mass appears to increase, it's because a gas is joining the reaction - like when magnesium burns and oxygen from the air gets added to form magnesium oxide. If mass seems to decrease, it's because a gas escapes - think of heating limestone where carbon dioxide gas floats away, leaving you with less solid material.

Thermal decomposition is simply using heat to break substances apart. When you heat metal carbonates (like limestone), they always split into a metal oxide and carbon dioxide gas - a pattern that's dead useful for predicting reaction products.

Quick Check: If your experiment shows unexpected mass changes, ask yourself: "What gas might be entering or leaving this reaction?"

Balancing Equations and Atomic Calculations

Balancing chemical equations is like solving a puzzle - you need equal numbers of each type of atom on both sides. Start by counting atoms of each element, pick the most unbalanced one, and add coefficients (the numbers in front) until everything matches up.

State symbols tell you the physical form: (s) for solid, (l) for liquid, (g) for gas, and (aq) for dissolved in water. These matter because they explain why mass might seem to change in reactions.

Relative atomic mass calculations help you find average atomic masses when isotopes are involved. The formula weighs each isotope by how common it is. Meanwhile, percentage composition tells you how much of each element is in a compound - divide the element's total mass by the compound's total mass, then multiply by 100.

Atom economy measures how wasteful a reaction is. If you only get one product, you've achieved 100% atom economy - everything you started with becomes useful product.

Pro Tip: Always write state symbols in equations - they're often the key to explaining unexpected results!

Moles and Avogadro's Constant

The mole is chemistry's way of counting particles - it's like saying "a dozen" but for atoms and molecules. One mole contains exactly 6.02 × 10²³ particles, which is Avogadro's constant.

Getting your units right is crucial. Always work in grams for mass calculations. If you're given kilograms, multiply by 1000 to convert to grams. If you've got milligrams, divide by 1000.

You can calculate the mass of a single atom by dividing the relative atomic mass by Avogadro's constant. This gives you the tiny mass of just one atom in grams - a mind-bogglingly small number that shows just how light individual atoms really are.

Think of moles as chemistry's universal translator - they let you convert between the number of particles, mass, and volume of gases. Once you master moles, quantitative chemistry becomes much more straightforward.

Memory Hook: Remember "MAV" - Mass, Amount (moles), Volume. These three quantities are connected in most chemistry calculations!

Empirical Formulas and Percentage Yield

Finding empirical formulas from experimental data follows a clear pattern. Take the mass of each element, divide by their atomic masses to get moles, then divide all results by the smallest number to find the simplest whole number ratio.

Percentage yield compares what you actually got in an experiment to what the theory predicted you should get. Real reactions rarely give 100% yield because they're messy - some reaction mixture gets lost, side reactions occur, or the main reaction doesn't go to completion.

Understanding why yields aren't perfect helps you improve experimental technique. Reversible reactions can't reach completion because products keep turning back into reactants. Some chemicals might react in unexpected ways, creating unwanted products instead.

When you're separating your desired product from the reaction mixture, some inevitably gets left behind. This is why industrial chemists spend so much time optimising separation techniques - every percentage point of yield recovered saves money.

Reality Check: If someone claims 100% yield in a complex reaction, be suspicious - real chemistry is rarely that perfect!

Limiting Reactants and Solution Concentration

Limiting reactants control how much product you can make - like having 10 burger buns but only 8 patties means you can only make 8 burgers. Calculate moles of each reactant, check the balanced equation's ratio, and see which runs out first.

The step-by-step method is foolproof: calculate moles, write the balanced equation to find molar ratios, then compare what you have with what you need. Whichever reactant you don't have enough of becomes your limiting factor.

Concentration can be expressed in mol/dm³ or g/dm³. The triangle method works brilliantly here - concentration equals moles divided by volume, with easy rearrangements for finding missing values.

Titration calculations rely on concordant results - readings within 0.10 cm³ of each other. Average these reliable results and ignore any obvious outliers that suggest experimental errors.

Lab Tip: Always identify your limiting reactant before predicting how much product you'll make - it's the chemical equivalent of checking you've got enough ingredients before cooking!

Gas Volumes and Experimental Uncertainty

Gas volumes follow predictable patterns at room temperature and pressure. One mole of any gas occupies 24 dm³ (or 24,000 cm³). This makes gas calculations surprisingly straightforward once you remember the conversion factors.

Whether you're working in dm³ or cm³ depends on your data, but the principle stays the same. The neat thing about gases is that volume directly relates to moles - double the moles, double the volume.

Uncertainty in measurements affects all experimental work. Calculate it by finding the range (biggest value minus smallest value) and dividing by 2. The more you measure, the smaller your percentage uncertainty becomes.

Reducing uncertainty improves data quality. Using larger quantities, more precise equipment, or taking multiple readings all help. Understanding uncertainty helps you evaluate whether differences between results are meaningful or just experimental noise.

Practical Point: When measuring gases, small uncertainties in volume can lead to big errors in calculated moles - always consider your experimental limitations!