The History of the Atom During all of history, there have been many different ideas about what an atom is (most of them incorrect). In the next unit we’ll discuss how atomic theory developed historically and figure out what the modern atom actually looks like.

Early Greek Models Leucippus, Democritus, Plato – 5th century BC. The early Greek philosopher Leucippus was the first guy who believed that matter was made of tiny, indivisible particles called atoms. His student Democritus usually gets credit for this discovery, though.

The Atomic Model Matter is made of empty space through which atoms move. Atoms are tiny, indestructible objects. Different atoms have different sizes and shapes (Plato’s addition: These shapes are geometric solids) Changes in matter are due to changes in the combinations of atoms. This theory, along with a non-atomic theory (favored by Aristotle) reigned for over two millennia.

Some random observations (18th century) For a very long time, nobody did any scientific experimentation, so the idea of the atom was left in the realm of pure thought. Enlightenment philosophy Taught that reason was the only true basis for knowing, so experimentation restarted.

Law of Conservation of Mass (Lavoisier): The weights of the products in a chemical reaction will be the same as the weights of the reactants. This was a revolutionary idea! If you burn wood the ashes weigh less than the wood originally did. Lavoisier was able to figure this out because he did a lot of work with gases and realized that they actually have mass.

Law of definite composition (Proust): A chemical compound will always have the same composition, regardless of how it was made. This was revolutionary because impurities made it so this didn’t always seem true.

John Dalton’s atomic theory (1808): All matter is made of small, indestructible particles called atoms. False: Atoms can be broken. A better statement that’s still true: “Matter cannot be broken by chemical means.”

Atoms of the same element have identical properties. False: The presence of isotopes disproves this.

Atoms of different elements have different properties. True

Atoms obey the law of conservation of mass The weight of what we make is equal to the weight of what we started with. True

Atoms obey the law of multiple proportions (this is known as Dalton’s law): If two elements can make more than one chemical compound, the ratios of the masses of one of the elements that combine with a given amount of the second element will be a small, whole number.

An example: Hydrogen and oxygen make two chemical compounds. In one compound, 2 grams of hydrogen combine with 16 grams of oxygen. In another compound, 4 grams of hydrogen combine with 32 grams of oxygen. The ratio of oxygen in one compound to another is 32:16 = 2:1.

Why is this important? It leads to the far more important understanding that atoms always combine in whole number ratios to form chemical compounds (i.e. H2O, never H2.1O0.8). This law is TRUE!

Dalton’s laws may seem fairly obvious to us, as do many of the correct principles from these other atomic theories. Let’s see why it’s not so obvious:

History of the Atom – Modern Theories of Atomic Structure

Thomson (1897): “Plum Pudding Model” Cathode Ray Experiment – showed electrons exist The model: Atoms are big balls of positive charge with negative charges (called electrons) embedded in it – imagine a 3-D chocolate chip cookie. Because there were no “anode rays”, the positively charged particles were seen as too big to move from one place to another. First characterization of the electron.

Thomson’s Plum Pudding Model Electrons (negative Charge) Positively Charged

Rutherford (1911): “Gold Foil Experiment” How he explained these observations: Positively charged alpha particles would pass through the gold foil most of the time without being deflected. The alpha particles would only be deflected when they passed near the small positively charged nucleus of the atom (because like charges repel one another)

Rutherford’s model of the atom: Nuclear model: The positive charge is all concentrated in the middle of the atom, while the negatively charged electrons float around throughout the rest of the atom’s empty space.

Our current model of the atom Says that the atom contains three types of subatomic particles that have the following properties: Proton (p+) Location- nucleus Charge +1 Mass - 1 amu (1.67 x 10-27 kg) Neutron (nº) Location - nucleus Charge - 0 Mass - 1 amu Electron (e-) Location - orbitals surrounding nucleus Charge -1 Mass ~ 0 (1/1836 amu, or 9.11 x 10-31 kg)

Writing Atomic Symbols: Atomic symbols have the general format: Where: A is the atomic mass (sometimes called the “mass number”), which is equal to protons + neutrons. Units of atomic mass are “amu.” Z is the atomic number. This is equal to the number of protons in all atoms and the number of electrons in neutral atoms. (Explain how the protons = electrons, making the charge cancel.) The atomic number determines the element. X is the atomic symbol. If you match the number of protons to the atomic number on the periodic table, you can figure out what the symbol is supposed to be. X A Z

Protons, neutrons, and electrons practice sheet The atomic masses on the periodic table won’t always match