Chapter 1.5: Evidence for Atoms

It is important to note that from the time that the first ideas of atoms arose, and for thousands of years thereafter, there was not one iota of evidence for the particulate nature of matter or the physical existence of atoms. The idea of atoms was purely a product of imagination, and while there was vigorous debate about the nature of matter, this debate could not be settled scientifically until there was objective evidence one way or another.

  

1.1 Atoms
1.2 Realities
1.3 History
1.4 Elements
1.5 Evidence
1.6 Parts
1.7 Iinteractions
1.8 He and H2
     

So the question arises, how did scientists in the nineteenth century eventually produce clear evidence for the existence of atoms? We have already said atoms are much too small to be seen by any direct method. So what would lead scientists to the unavoidable conclusion that matter is composed of discrete atoms? In fact, often a huge intuitive leap must be made to explain the results of scientific observations.

For example, the story about Isaac Newton (1643-1727) and the falling apple captures this truism, namely the remarkable assumption that the movement of the earth around the sun, the movement of the moon around the earth, and the falling of an apple to earth are all due to a common factor, the force of gravity, which acts at a distance and obeys an inverse square relationship (1/r2, where ”r” is the distance between two objects). This seems like a pretty weird and rather over-blown assumption; how does this “action at a distance” work? Yet, followed scientifically it appears to be quite powerful and remarkably accurate. The point is that Newton was able to make sense of the data - something that is in no way trivial. It requires a capacity for deep, original and complex thought. That said, it was not until Albert Einstein proposed his general theory of relativity (1915) that there was a coherent mechanistic explanation for gravitational forces.

The first scientific theory of atomic structure was proposed by John Dalton (1766 - 1844), a self taught Quaker [religious dissenters, that is non-Anglicans, were not allowed access to English universities at that time ] living in Manchester, England. In 1805 Dalton published his atomic theory to explain the observed law of multiple proportions. Rather surprisingly, Dalton never really explained what led him to propose his atomic theory, although he certainly used it to explain existing rules about how different elements combine. Among these rules (or Laws) was the observation that the total matter present in a system did not change when a chemical reaction occurred, although a reaction might lead to a change from a solid to a gas or vice versa. Inspection of his laboratory notebooks suggests that he first began to develop this atomic theory as he was experimenting on the nature of different gases and their solubility in water. In 1803 he wrote:

"Why does not water admit its bulk of every kind of gas alike? This question I have duly considered, and though I am not able to satisfy myself completely I am nearly persuaded that the circumstance depends on the weight and number of the ultimate particles of the several gases. (emphasis added)" .

This was the first indication that he was thinking about gases in terms of particles (atoms/molecules, that is combinations of atoms). In case you missed this extraordinary deduction (and it is very easy to do), Dalton made the leap to the idea of atoms with different weights from his observations that different gases dissolved in water in different amounts and the law of multiple proportions.

Dalton’s atomic theory (1805) had a number of important points:

  • Elements are composed of small indivisible, indestructible particles called atoms
  • All atoms of an element are identical and have the same mass and properties
  • Atoms of a given element are different from atoms of other elements
  • Compounds are formed by combinations of atoms of two or more elements
  • Chemical reactions are due to the rearrangements of atoms, atoms (matter) are neither created nor destroyed during a reaction.


Based upon these tenets he was able to explain many of the observations that had been made up to that time, by himself and others, about how matter behaves and reacts. More modern atomic theories have made some modifications, for example to include the existence of atomic isotopes (that is: atoms with different numbers of neutrons, but the same number of protons and electrons) and the conversion of energy into matter and vice versa, but Dalton’s core ideas remain valid.

The law of multiple proportions

The law of multiple proportions is an empirical law, that is a law based on observation rather than theoretical logic. It states that when two elements (for example carbon and oxygen) combine to form more than one type of compound, such as carbon monoxide and carbon dioxide, the ratio of the mass of oxygen in carbon monoxide is always in some whole number to the mass of oxygen in carbon dioxide.

Mass O in carbon monoxide / Mass carbon monoxide = an integer x Mass O in carbon dioxide / Mass carbon dioxide

This law makes complete sense in terms of atomic theory which assumes that each molecule is composed of a whole (positive integer) number of atoms, and each atom of a particular element is identical to every other atom of that element, that is, an atom of oxygen in carbon monoxide is the same as an atom of oxygen in carbon dioxide, or an atom of oxygen in any other imaginable molecule. Atoms cannot be divided into parts, that is, there is no such thing as a half or a quarter atom; they also do not have a memory of where they have been. An atom of oxygen that was once in the brain of a dinosaur behaves no differently than an atom of oxygen that has been in the oceans for the last 300 million years. The atomic formula for carbon monoxide is CO and the formula for carbon dioxide is CO2. You wouldn’t dream (we hope) of writing carbon dioxide as C0.5O or water (H2O) as HO0.5. Such formulae would make no sense in modern atomic theory.

Connecting the real world and the molecular world

One problem chemists have is that we deal with things that happen at a scale that we cannot see (and is very difficult to imagine). As a result we have to develop a number of skills and tools to connect the molecular world with the world we can see. For example: a large part of the later sections of the book is taken up with developing ways to help you visualize molecular level structures and events. But, while thinking about the molecular level is important, it is also necessary to be able to connect those molecular changes with what we can see and measure in the world we live in. For example we might visualize a chemical reaction as a molecule of one reactant interacting with another molecule to give a product. We generally write reactions down in the form of equations, for example:
2H2(g) + O2(g) → 2 H2O(g)


This equation can mean many things - the simplest being that two molecules of hydrogen react with one molecule of oxygen, resulting in two molecules of water. However in the world that we live in we can not measure out materials in terms of atoms or molecules, we have to use mass or volume. In order to relate such measurements to the reaction equation, we need a way to translate between the number of molecules and mass. The unit used for this purpose is the mole. The mole is simply a number, 6.022 x 1023. It is a big number to be sure – but just a number. So why this number? The reason is that it enables us to convert directly between the mass of atoms and molecules, measured in atomic mass units (amu), and the mass of the element or compound in grams.

One atom of carbon-12 is defined as having a mass of exactly 12 amu, so 1 mole of carbon-12 has a mass of 12 grams and 12 grams of carbon-12 contains 6.022 x 1023 carbon atoms.

All the other elements have masses that are defined relative to this mass. So for example, the equation above could mean all these things:  


Once you know this relationship and the atomic mass of all the atoms in your reaction, you can calculate the mass in grams of every reactant or product from a given mass or reactant (or product). In the accompanying workbook there are a number of activities that will allow you to work through some of these kinds of calculations.

A note on the conservation of matter:

A key component of Dalton's atomic theory was the assumption of the conservation of matter, that is that matter can not be created nor destroyed, but can change between different states. The most common example is the transformation of water from ice (solid) to liquid to vapor (gas). People are often confused about the conservation of matter, one because it is not completely obvious that when a cube of ice evaporates matter is not, in fact, lost. This is something that requires careful experimental observations to confirm - it is certainly not self-evident. Another factor which may contribute to confusion is that in the modern world, most people have been exposed to Einstein's famous equation
e (energy) = m (mass, which is a measure the amount of matter) x c2 (the speed of light, squared).

Based on this equation, you might reasonably assume that matter and energy are freely interconvertable, but the conditions under which matter converts into energy or energy into matter are not so common, and when they occur in the “normal world” they involve extremely small mass changes. When plants absorb light, they do not convert it into matter, but use the energy to rearrange atoms and molecules, a topic we will return to later.

 

All living organisms use some kind of energy to make changes, but energy is not directly converted into matter. In fact it is the interconversion of matter into energy that is ultimately responsible for the light given off by the sun and the energy released by nuclear power plants (and atomic bombs). In Dalton's day, the possibility of the interconversion of matter and energy was not known, and from our perspective as chemists is not something we need to consider. But so that we do not confuse you further, energy and matter are forms of the same basic stuff and the total amount of energy + matter in the universe is a constant (or so it appears to modern astrophysicists).

 

1.1 Atoms
1.2 Realities
1.3 History
1.4 Elements
1.5 Evidence
1.6 Parts
1.7 Iinteractions
1.8 He and H2

Question to answer:

  • In what ways is Dalton’s atomic theory different from the ideas of the Greek philosophers?
  • Which tenets of Dalton’s theory still hold up today?
  • Design an experiment to investigate whether there is a change in mass when water changes phase.
  • What data would you collect? How would you analyze it?

Questions for ponder:

  • How did Dalton conclude that there were no half-atoms?
  • Which parts of Dalton's theory were unfounded speculation and which parts based on direct observation?
27-Jun-2012