The Structure of Atoms

Outline

• Fundamental Particles
  • Dalton, resummarized
  • Stoichiometry
  • Atoms: Points to Remember
  • A Short History of Atomic Masses
  • Isotopes
• Practice with Concepts
• Discussion Questions
• Optional Website Reading

Fundamental Particles

Recognizing that atoms were themselves divisible and composed of even more fundamental particles was a challenge for chemists working in the nineteenth century, when they were focussed on how atoms of different elements interacted, rather than on the interior structure of the atom.

Dalton's Laws

Dalton's original observations on how elements combine are so well attested that we now call them laws. Don't be mislead by the term: a law doesn't mean that the theory it describes is absolutely proven. No theory can be proven in the strictest sense (although any theory can be disproven ). We continue to use the atomic model because it accounts for many different observations, and has been useful in predicting the outcome of a number of experimental tests.

The conservation of mass is a basic principle of the universe. It means that (excepting whatever initial incident got us here) no mass is now being created or destroyed. In an ordinary chemical reaction, there is no difference between the mass of the reactants and the mass of the products. In a nuclear reaction, any mass loss is compensated by a release of energy equal to the mass loss times the speed of light squared (Einstein's famous E=mc²).

The law of constant composition means that a compound is always formed out of the same components. Water is always made of two parts of hydrogen to one part oxygen. Changing one of the components changes the compound into something else.

The law of multiple proportions means that elements combine in whole-number ratios (because the atoms are particles which are combining). This makes it much easier to determine the composition of complex molecules.

Stoichiometry

Stoichiometry comes from the Greek stoichoien (element) and metron (measure). It is used in different subjects to identify the elements of that subject. In chemistry it refers to the application of the laws of definite proportions and of the conservation of matter and energy to chemical reactions and processes. One of the main activities in chemical stoichiometry is balancing equations of chemical reactions, and determining how much mass of a compound was involved, or even how many atoms or molecules were involved in a particular reaction.

We'll come back to this term and its practical application, but before we can talk about measuring quantities in chemical reactions, we have to establish some units for atoms, and those units need to be based on the physical characteristics on atoms and elements.

Atoms: Points to remember

Here are the important things to remember about the characteristics of atoms:

• The number of protons (the atomic number) in an atom determines the type of element it is. Change the number of protons, and you change the kind of atom it is.
• Protons and neutrons have almost the same mass. An atomic mass unit (amu) is equal to the mass of a single proton. Protons and neutrons exist at the center (nucleus) of the atom.
• Protons and electrons have the same amount of electrical charge, but protons are positively charged, while electrons are negitively charged.
• In a neutrally-charged atom, the number of electrons is the same as the number of protons. Electrons can be added to some atoms (making them negatively charged) and taken from most (leaving them positively charged); this changes only the charge and not the kind of atom. It doesn't effectively change the mass of the atom.
• Atomic mass (mass number) for an individual atom is calculated by adding the number of protons and neutrons (non-charged particles in the nucleus) together; electrons are ignored because their mass is so small (1/1800 of a proton's mass).
• Electrons exist in special areas (called orbits, but don't confuse these with the nice paths the planets follow around the sun) around the nucleus of the atom. Two electrons can be in one orbit at a time, each with a different spin direction. To remember how many electrons go in which level, generate this table:

  Level	Types	Total # of orbits	Total # of electrons
  1	S (spherical)	1	2
  2	S, P	4	8
  3	S, P, D	9	18
  4	S, P, D, F	16	32

  The number of orbits is the square of the level; the number of electrons is double the number of orbits. We'll get to the different types of orbits in chapter 8.
• Electrons move between orbits by absorbing or emitting packets (quanta) of energy in the form of light. The amount of energy determines the color of light absorbed or emitted. All atoms of a given element and charge have the same spectra, or set of colors, which they give off when heated and then cooled.
• The distribution of electrons around the atom determines how it forms molecules with other atoms (compounds are molecules made up of two or more different elements). By comparing the chemical properties of various elements, the Russian chemist Dmitri Mendeleev in 1869 was able to create the very useful periodic table. All the elements in a column of the table share similar abilities to form molecules.
• The total number of electrons in the outermost level of an atom determines its position in the periodic table. Notice that Period 1 has 2 elements, Period 2 has 8, Period 3 has 8, then Period 4 has 18 (level 4 s and p orbitals fill, then level 3 d fills up). We'll study all this information in great detail: the key to remember now is that the electron distribution in the atom determines its chemical reactivity properties.
• Most of the elements on the periodic table are classified as metals (and they are all on the left side of the table), and are shiny, malleable solids at room temperature, capable of conducting heat and electric current easily. Nonmetals (all on the far right side of the table) are solids or gases at room temperature, conduct heat and metal poorly, and (in the case of solids) are usually rigid or brittle. Eight elements are classified as metalloids or semi-metals, and share characteristics of both metals and non-metals. The noble gas atoms (elements in the farthest column on the right) do not form molecules; they have completely full electron orbitals. Atoms of all other elements are normally found in molecules.
• Atoms tend to acquire or release electrons so that they have the configuration of a noble gas.

A Short History of Atomic Masses

Let's look at the real problems chemists ran into in trying to determine compound proportions. Assume that we don't have access to the periodic table, and we haven't yet discovered that we can figure out the number of atoms by balancing the valence electrons involved.

In the late 1700s, a number of chemical researches led to the realization that atoms combine to form compounds according the the law of constant proportionality, which became of the three major hypotheses of Dalton's atomic theory. Hydrogen is the lightest element, and combines with oxygen in the mass proportion 1:8 to form water. Other compounds can also be analyzed in terms of their relative masses. Note that this doesn't tell us how many atoms are involved in the compound. The early chemists, John Dalton among them, believed that water was made of one hydrogen and one oxygen atom.

How can we determine how many atoms are involved? In the early 1800s, Avogadro suggested that at the same pressure, temperature, and volume, different gasses will have the same number of atoms or molecules. Since (keeping temperature and pressure constant) a volume of hydrogen gas combines with a volume of chlorine gas to produce two volumes of hydrogen chloride, this would mean either that the hydrogen and chlorine atoms split, or that (as Avogadro proposed) hydrogen and chlorine existed as 2-atom molecules in their gaseous form. Most chemists followed Dalton's assumption [wrong, as it turns out] that like atoms repelled one another and could not form molecules like H₂. They rejected Avogadro's suggestion for over 50 years. By 1861, however, there was enough evidence to support it, and chemists begin to use Avogadro's hypothesis use it to determine the proportions of compounds by atom count as well as by weight.

Experiments like sparking a mixture of hydrogen gas and oxygen gas to produce water led to the determination that while the mass proportions of water are 1 mass of hydrogen to 8 masses of oxygen, the atomic proportions are 2 hydrogen atoms to 1 oxygen atom. With this as a scale, chemists were able to determine a atomic weight of 16 "hydrogen masses" for oxygen.

The atomic weights for non-gasses were determined using other techniques. In the early 1800s, research in metals yielded some peculiar information about the way metals absorb heat energy. Specific heat is the amount of heat, in calories, required to raise the temperature of 1 gram of a substance 1 degree centigrade. [By definition, a calorie is the amount of heat required to raise the temperature of 1 gram of water from 20 degrees centigrade to 21 degrees centrade at sea level, so specific heats are always done in comparison to water.] The specific heat of each element and compound is one of the physical characteristics of that compound. In 1819, Pierre Dulong and Alexis Petit discovered that the atomic weight of an element multiplied by the specific heat of the element was a constant of approximately 6.3. This is an empirical relationship: while Dulong and Petit didn't understand the reason for the relationship, they were able to apply it to many, but not all, metals successfully.

Isotopes

So far, we haven't used absolute masses at all, only relative masses or ratios. Determining a standard for atomic masses, even relatively, though, proved to be more difficult than anticipated. Originally, chemists adopted oxygen as the standard, with an atomic weight of 16. It combines readily with many elements, and more easily isolated (and more safely stored!) than hydrogen. As techniques were refined, however, it became clear that using the O-16 scale was unacceptable: hydrogen was calculated to have a mass of less than 1 atomic weight, and other elements had fractional rather than whole number atomic weights. At the atomic level, this would mean fractional atoms were combining with each other, in violation of the theory that atoms were the simplest components of matter.

It wasn't until the 1930s, after the discoveries that the "unsplittable" atom was actually composed of electrons, protons and neutrons, that chemists were able to determine why they found fractional atomic weights. Most elements have atoms with different numbers of neutrons in the nucleus; these are istopes of the element. Some isotopes are unstable and lose neutron/proton pairs from the nucleus over time in bursts of radiation, so they are called radioactive isotopes. In any given sample of oxygen, there would be some O-16 atoms, composed of 8 protons and 8 neutrons, and some O-18 atoms, composed of 8 protons and 10 neutrons. The mass standard was actually too high.

Once the dominance of carbon in organic molcules was realized, chemists and biologists agreed to reset the standard for atomic masses. An atomic mass unit is now defined as 1/12 of the isotope C-12 atom.

The periodic table gives atomic weights in terms of a natural sample of the element. If you look on the periodic table, you will see the atomic weight of carbon is given as 12.01; this is because in nature, carbon occurs in C-12, C-13, and radioactive C-14 isotopes.

Practice with the Concepts

Putting all the discoveries together allows us to look at different compounds and determine their composition by weight and by number of atoms.

Discussion Questions

• Why do we need to convert from mass to atoms or moles? Why not just count atoms?

Optional Readings

• Explore the Isotopes for different elements at Brookhaven National Laboratory's Table of Isotopes site.

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