Have you ever tried to observe or study something that you could not see?

 Scientists are faced with this problem each and every day. Geologists must attempt to describe things such as the interior of Earth or the way in which mountains and deserts came to be. Astronomers attempt to map the outer reaches of the Universe. Chemists and physicists across the centuries have been concerned with the smallest particles that make up each and every object in our natural world.

 Of course, that is the atom. How do scientists study things they can’t see? They make models. A scientific model is a tool constructed by the scientist based on all the known experimental evidence about a particular thing, such as an atom.  As time goes by and more experiments are performed, models evolve and change to account for new understanding. In this lesson, you will begin to learn about how the model of the atom was initially developed and how it has changed over time into what we now have come to accept as the modern model of the atom.

 Lesson Objectives  Explain the law of conservation of mass, the law of definite proportions, and the law of multiple proportions.  Describe John Dalton’s atomic theory.

 The understanding of the atom and atomic structure has evolved greatly over time. From early Greek philosophy to modern scientific experimentation, our understanding of this most basic unit of matter has changed frequently. This chapter will look at the beginnings of that evolution, focusing on the atomic model of John Dalton and the scientific laws upon which his model was based.

Early Atomic Models  Democritus was a Greek philosopher who lived from 460 B.C. until 370 B.C. He was among the first known individuals to suggest the idea of the atom as the basic unit of matter. The word atom comes from the Greek word atomos, meaning “indivisible.”

 Democritus was indeed ahead of his time, but his ideas were not useful at that time in describing chemical behavior because there was no experimental evidence to support it. His approach was a philosophical one rather than a truly scientific one. Many centuries would pass before the notion of atoms was merged with modern scientific experimentation and thought.

Beginnings of Atomic Theory  By the late 1700s, chemists had accepted the definition of an element as a substance that cannot be broken down into simpler substances by ordinary chemical means. It was also clear that elements combine with one another to form more complex substances called compounds. The chemical and physical properties of these compounds are different than the properties of the elements from which they were formed.

 There was some disagreement, however, about whether elements always combined in exactly the same ratio when forming a particular compound.  In the 1790s, a greater emphasis began to be placed on the quantitative analysis of chemical reactions. Accurate and reproducible measurements of the masses of reacting elements and the compounds that they form led to the formulation of several basic laws.

 One of these is called the law of conservation of mass, which states: that during a chemical reaction, the total mass of the products must be equal to the total mass of the reactants. In other words, mass cannot be created or destroyed during a chemical reaction, but it must always be conserved.

 The discovery that mass is always conserved in chemical reactions was soon followed by the law of definite proportions, which states that a given chemical compound always contains the same elements in the exact same proportions by mass.

 As an example, any sample of pure water contains 11.19% hydrogen and 88.81% oxygen by mass. It does not matter where the sample of water came from or how it was prepared. Its composition, like that of every other compound, is fixed.  Experiments also began to demonstrate that the same pairs of certain elements could combine to form more than one compound.

 Consider the elements carbon and oxygen. Combined in one way, they form the familiar compound called carbon dioxide.  In every sample of carbon dioxide, there is 32.0 g of oxygen present for every 12.0 g of carbon.  By dividing 32.0 by 12.0, this simplifies to an oxygen/carbon mass ratio of 2.66 to 1.  Another compound that forms from the combination of carbon and oxygen is called carbon monoxide.  Every sample of carbon monoxide contains 16.0 g of oxygen for every 12.0 g of carbon, which simplifies to an oxygen/carbon mass ratio of 1.33 to 1.

 In other words, a given mass of carbon needs to combine with exactly twice as much oxygen to make carbon dioxide as it would to produce carbon monoxide. The figure below illustrates the law of multiple proportions. Whenever the same two elements form more than one compound, the different masses of one element that combine with the same mass of the other element are in the ratio of small whole numbers.  Carbon monoxide, on the left, contains g  of oxygen for every 1 g of carbon.  Carbon dioxide, on the right, contains g of oxygen for every gram of carbon. The ratio of oxygen in these two compounds is 1:2, which is a ratio of small whole numbers.

 Sample Problem: Calculating Mass Ratios  Copper reacts with chlorine to form two compounds. Compound A contains 4.08 g of copper for every 2.28 g of chlorine. Compound B contains 7.53 g of copper for every 8.40 g of chlorine. What is the lowest whole number mass ratio of copper that combines with a given mass of chlorine?

 Step 1: List the known quantities and plan the problem.  Known  Compound A = 4.08 g Cu and 2.28 g Cl  Compound B = 7.53 g Cu and 8.40 g Cl  Apply the law of multiple proportions to the two compounds. For each compound, find the grams of copper that combine with 1.00 g of chlorine by dividing the mass of copper by the mass of chlorine. Then, find the ratio of the masses of copper in the two compounds by dividing the larger value by the smaller value.

 Step 3: Think about your result.  The ratio is a small whole-number ratio. For a given mass of chlorine, compound A contains twice the mass of copper as compound B.

Dalton’s Atomic Theory  In 1808, an English chemist and schoolteacher named John Dalton ( ) formulated an atomic theory based on the law of conservation of mass, the law of definite proportions, and the law of multiple proportions. His theory can be summarized in the following statements.  All matter is composed of extremely small particles called atoms.  Atoms of the same element are identical in terms of size, mass, and other properties. Atoms of one element are different from the atoms of any other element.  Atoms of different elements can chemically combine with one another in simple whole- number ratios to form chemical compounds.  Chemical reactions can be described as the separation, combination, or rearrangement of atoms. Atoms of one element cannot be changed into atoms of a different element as a result of a chemical reaction.

Dalton’s atomic theory states that chemical reactions are due to rearrangements of atoms. On the left, hydrogen molecules are shown as two H atoms, while oxygen molecules are shown as two O atoms. The chemical reaction forms water as the atoms rearrange. There are identical numbers of hydrogen and oxygen atoms before and after the reaction.

Sizes of Atoms  The graphite in your pencil is composed of the element carbon. Imagine taking a small piece of carbon and grinding it until it is a fine dust. Each speck of carbon would still have all of the physical and chemical properties of carbon.  Now imagine that you could somehow keep dividing the speck of carbon into smaller and smaller pieces. Eventually, you would reach a point where your carbon sample is as small as it could possibly be. This final particle is called an atom, which is defined as the smallest particle of an element that retains the properties of that element.

 Atoms, as you know, are extremely small. In fact, the graphite in an ordinary pencil contains about 5 × atoms of carbon. This is an almost incomprehensibly large number. The population of the entire Earth is about 7 × 10 9 people, meaning that there are about 7 × times as many carbon atoms in your pencil as there are people on Earth! For this to be true, atoms must be extremely small.  Can we see atoms? It’s not easy, but a modern instrument called a scanning tunneling microscope allows scientists to visualize the atom.  Images of individual gold atoms can be seen on the surface of a smooth sheet of gold metal using scanning tunneling microscopy.

Lesson Summary  The Greek philosopher Democritus believed that matter is composed of indivisible and indestructible building blocks, which he called atoms.  John Dalton transformed the ideas of Democritus into a scientific atomic theory, which began to explain the law of conservation of mass, the law of definite proportions, and the law of multiple proportions.  Matter cannot be created or destroyed in chemical reactions.  Elements react to form compounds in fixed proportions by mass.  Elements combine with one another in simple whole-number ratios.  An atom is the smallest unit of an element that retains the properties of that element.