Comprehensive Chemistry Revision Notes

Atomic Structure and Subatomic Particles

Think of an atom as the smallest piece of an element that still acts like that element - it's literally the foundation of everything you can touch! Inside every atom, you'll find three main subatomic particles that work together like a perfectly balanced team.

The nucleus sits at the atom's centre, packed with protons (positive charge, mass = 1) and neutrons (no charge, mass = 1). Meanwhile, electrons (negative charge, almost no mass) whizz around the nucleus in shells. Here's the clever bit: atoms always have equal numbers of protons and electrons, making them electrically neutral overall.

Every atom gets two important numbers that tell its story. The atomic number shows how many protons it has (like an element's ID card), whilst the mass number counts all the protons and neutrons together. For example, potassium has 19 protons and a mass number of 39.

Quick Tip: Remember that nearly all an atom's mass comes from the nucleus - electrons are roughly 1,840 times lighter than protons!

Isotopes and Relative Atomic Mass

Here's where atoms get interesting - isotopes are like different versions of the same element! They have identical numbers of protons (same element) but different numbers of neutrons, giving them different masses. It's like having the same car model but with different engine sizes.

Hydrogen-1 is the most common isotope of hydrogen, whilst carbon has several isotopes including carbon-12. The name always tells you the mass number - dead simple! Since electrons determine chemical properties, isotopes of the same element react identically despite having different masses.

Relative atomic mass is basically a weighted average that accounts for all an element's isotopes and their natural abundance. You'll find these values on the periodic table with the symbol Ar. The formula is: (sum of isotope abundance × isotope mass number) ÷ (sum of all abundances).

Key Point: Mass numbers are always whole numbers, but relative atomic masses often aren't because they're averages of different isotopes!

The Periodic Table and Atomic Models

Mendeleev's periodic table was brilliant for its time - he arranged elements by increasing relative atomic mass and spotted repeating patterns (periodicity). He even left gaps for undiscovered elements and predicted their properties! The only hiccup was tellurium and iodine, which he had to swap based on their chemical properties rather than mass.

Today's modern periodic table arranges elements by atomic number (number of protons) rather than mass, solving Mendeleev's tellurium-iodine problem. Metals live on the left, non-metals on the right, with horizontal rows called periods and vertical columns called groups.

The journey from plum pudding model (1897) to nuclear atom was revolutionary. Rutherford's gold foil experiment (1905) showed that atoms are mostly empty space with a tiny, dense, positive nucleus. When alpha particles bounced back, it proved the nucleus was incredibly small but contained all the positive charge and most of the mass.

Bohr's model (1913) suggested electrons orbit in specific energy levels, explaining why different elements produce distinct flame colours. When electrons absorb energy, they jump to higher levels, then release light of definite frequencies when dropping back down.

Remember: The nucleus is about 1/10,000th the size of the whole atom - imagine a marble in a football stadium!