The Building Blocks of Matter

Gap-Fill Exercise
Content © 2006 Melanie Cecere; Authors of Activity: Melanie Cecere & Tami Maloney; All rights reserved. No commercial, for-profit use of this material is allowed. E-mail comments and questions to Tami Maloney.

Fill in all the gaps, then press "Check" to check your answers. Use the "Hint" button to get a free letter if an answer is giving you trouble. You can also click on the "[?]" button to get a clue. Note that you will lose points if you ask for hints or clues!
About 2500 years ago a Greek philosopher named came up with the idea that all matter is composed of tiny, indivisible particles that he called "atoms" (which means "" in Greek). This is sometimes called the "discontinuous theory". Democritus believed that there were 4 elements-, , , . Not much happened in the field of atomic theory for more than 2000 years after that.
In the late 1700's, the concept of elements was universally accepted. Scientists differed on whether or not elements always combined in the same ratios for a particular compound. The transformation of a substance or substances into one or more new substances is known as a . In the 1790's increased study of chemical reactions and improved measuring techniques resulted in the discovery of several basic laws:

- The law of conservation of . Antoine was the first to explain correctly the theory of burning (1770's) based on careful measurements. He theorized that mass is neither nor during ordinary chemical or physical reactions.

- The law of proportions. Joseph developed a theory that said a chemical compound contains the same elements in exactly the same proportions by mass regardless of the size of the sample or source of the compound. For example, in NaCl, the proportion is always 39.34% Na and 60.6% Cl.

- The law of proportions. (John , 1803) If two or more different compounds are composed of the same two elements, then the ratio of the mass of the second element combined with a certain mass of the first element is always a ratio of small numbers. Some elements combine to form only one compound (NaCl). But some, like H and O, form more than one compound.

b. Dalton's Atomic Theory. In 1808, John Dalton combined these laws into an explanation based on atoms. He reasoned that elements were composed of and the only whole numbers of atoms can combine to form . Summed up:
- All is composed of extremely small atoms.
- Atoms of a given element are in size, mass, and other properties.
- Atoms cannot be subdivided, created, or .
- Atoms of different elements combine in simple ratios to form chemical compounds.
- In , atoms are combined, separated, or rearranged.

c. Modern Atomic Theory. Some of Dalton's ideas didn't hold up. In modern theory, atoms are NOT . Atoms can be changed from one element to another (not in chemical reactions, but in reactions). Atoms of the same element are not all exactly alike. Isotopes exist. But these two major concepts of Dalton did hold up:

- All matter is composed of .
- Atoms of any one element differ in properties from atoms of another element

The Structure of the Atom. An is the smallest particle of an element that retains the chemical properties of that element. Atoms have two regions. The is located near the center of the atom and contains one or more protons and may contain one or more neutrons. Protons are charged particles; are particles that are neutral. travel around the outside of the nucleus and have a negative charge.

a. Discovery of the electron. Electrons were discovered by J.J. in 1897 (almost 100 years after Dalton). In the 1870s, an English physicist named William studied the behavior of gases in a type of vacuum tube (cathode ray tube or "Crookes tube"). A cathode ray tube is a glass tube with air pumped out with high voltage to two (cathode is negative, anode is positive). The ray travels from the negative to the side of the tube. The wall opposite the cathode developed yellow/green fluorescence. An object placed in the middle of the tube casts a , and a paddle wheel placed on a rail between the electrodes rolled along the rails from the cathode to the . This showed that the ray had sufficient to set a wheel in motion. Crookes said some type of radiation or particles traveling across the tube were casting the shadow. He called them cathode rays. Crookes discovered that a magnetic field would deflect them the same way a wire carrying an electric current (which is known to have a negative charge) does. Additionally, the rays were deflected away form a charged object. These observations led observers to believe that the ray was composed of charged particles. Thompson later named these particles .
Later, the mass of the was found to be 9.109 x 10-31 kg. That's almost one two-thousandths the mass of the smallest known atom. This led to two more inferences about the atom. (1) Because atoms are electrically neutral, the atom must contain a charge to balance the negative electrons. (2) Because electrons have so much less mass than an atom, atoms must contain other particles that account for most of their .

b. Discovery of the nucleus. In 1911, Ernest (along with associates Hans Geiger and Ernest Marsden) conducted experiments that eventually led to the discovery of the . He bombarded foil with fast moving, positively charged particles (alpha particles. An alpha particle is a nucleus of a helium atom, 2 protons and 2 neutrons.) He surrounded the foil with a screen coated with ZnS. Each time an alpha particle hit the ZnS coating, a flash of was produced at the point of contact. By observing the flashes, he could see whether the a particles were deflected from their straight path. Most of the particles went straight through as if there wasn't any foil. This led him to the conclusion was that the atom is mostly . However, to his amazement, 1 in 8000 particles bounced back. Some were deflected slightly and some deflected at large . This led him to the conclusion that the atom must contain a small, dense, positively charged central portion or . The positively charged nucleus repelled the alpha particles that came near. The nucleus is so that not many "hits" were scored. Most went right through the "empty" portion. The atom must be mostly empty space between a positively charged nucleus (in which most of the mass is located) and the electrons that defined the volume of the atom.
- Shortcomings of Rutherford's Model: A large positive nucleus should electrons. Rutherford said that the motion of the electrons would keep them from falling in toward the nucleus (like the moon around the earth). BUT, classical "Newtonian" physics states that a charged particle traveling in a curved path radiates energy and the electrons would continually radiate energy and would fall into the nucleus (like artificial satellites do-they lose energy through friction with the earth's atmosphere and fall to lower orbits). BUT, atoms are stable and do not collapse, SO Neils Bohr concluded that classical physics could not apply to in an atom.
- The Model. In 1913, the Danish physicist Niels Bohr proposed improvements to the Rutherford model. In Bohr's model, there are certain definite in which an electron can travel around the nucleus without radiating energy. Each of these orbits is a circular orbit at a fixed distance from the nucleus. An electron in a given orbit has a certain amount of energy. The greater the distance of the electron from the nucleus, the greater the of that shell. The possible electron orbits became known as energy . The only way an electron can lose energy is by dropping from one energy level to a one. The electron then emits a of radiation that corresponds to the difference in energy levels. As long as they remain in orbit, they do not lose energy. In the state, all electrons in an atom are in the lowest energy level available. In the state, the electrons are at a higher energy level than the ones they normally occupy. The Bohr model is NOT the current model of the atom. The modern model is known as the " model". More on that next chapter!

c. Composition of the nucleus. Further experiments showed that the nucleus was made up of still smaller particles called . Rutherford realized that protons could not account for the entire of the nucleus. He predicted the existence of a neutral particle that would account for the missing mass. In 1932, James confirmed the existence of the . So what holds the nucleus together, you ask? forces: short-range proton-neutron, proton-proton, and neutron-neutron forces hold the nuclear particles together. For example, when two protons are extremely close together there is a strong attraction between them. More than 100 protons can exist close together in a nucleus. In summary--

d. The size of the atom. The region occupied by electrons is referred to as the electron . The distance between the center of the nucleus and outer limit of the cloud is the atomic . Atomic radii range from 40 - 270 picometers (pm.) (1 pm = 10-12 m or 10-10 cm. 1 cm : 100 km :: 100 pm : 1cm) Nuclei have a radii of about 0.001 pm. Density of nuclei -- 2 x 108 metric tons/cm3.

Counting Atoms
a. Atomic number: Atoms of the same element have the same number of . The atomic number is the # of in the nucleus of each atom of that element. The atomic # identifies an element. The number is the total number of protons and neutrons in the nucleus. are atoms of the same element that have different masses, but the same chemical behavior. is the general term for any isotope of any element. There are two ways to express which nuclide you're talking about. notation uses the atomic symbol and mass (eg, C-12). Nuclear notation uses the atomic symbol mass number, and number.

b. Relative Atomic Mass. One atom of O-6 has a mass of 2.657 x 10-23 g. It's more convenient to use relative atomic masses. C-12 is the standard. C-12 was arbitrarily assigned a mass of 12. One atomic mass unit () is exactly 1/12 the mass of a C-12 atom or 1.660540 x 10-27 kg. Relative weights are directly proportional to actual weights.

c. Average Atomic Masses. The atomic mass is the weighted average of the atomic masses of the naturally occurring isotopes of an element. For example, Cu has an average atomic mass of 63.55. In nature, 69.17% is Cu-63 (mass of 62.929598 amu), and 30.83% is Cu-65 (mass of 64.927793 amu).

(.6917)(62.929598) + (.3083)(64.927793) = 63.55 amu

d. The Mole: The is a unit in chemistry that is used to describe a certain number of particles. Just as a "dozen" items is shorthand for 12 items, the mole is shorthand for number of items. Atoms and molecules are quantified in , and the mole serves as a bridge connecting all the different quantities that you'll come across in chemical calculations. A mole is defined as the quantity of a substance that has a in grams numerically equal to its weight or atomic weight. The symbol for mole is "".
The average mass of carbon is 12.01 amu. This is a ridiculously tiny number of grams. It is too small to handle normally. The mass of a substance is its average atomic mass in grams from the Periodic Table. The molar mass of carbon, for example, is numerically equal to the average atomic . This means that 1 mole of carbon is equal to grams of carbon. The molar mass is the weight in grams in one . It is the atomic weight (or molecular weight) of a substance in . The unit for molar mass is . The mole can be used as a factor. First, find the molar of each of the atoms in the substance in question. Then calculate molecular mass for one of the substance. Using the mole as a conversion factor, you can convert moles to grams, grams to moles, moles to # of molecules, # of molecules to moles, etc.

Nuclear Chemistry

a. Radioactive decay. The nucleus of many isotopes is unstable. occurs when an isotope emits particles of protons and neutrons in order to stabilize its nucleus. There are two primary types of decay.
- decay occurs when an alpha particle (2 neutrons and 2 protons) is emitted. Alpha decay reduces the atomic number by and the atomic mass by . You end up with a totally different element. Large, heavy elements tend to undergo alpha decay. Alpha particles are relatively large (radioactively speaking) and it can be by something as thin as a piece of notebook paper. They can't even penetrate your outer layer of !
- decay is the emission of an from the nucleus (yes, the nucleus-a proton can transform magically into a neutron by giving off an electron). Because the number of protons has changed, you end up with a totally different element. Usually, when a nucleus undergoes beta decay, the atomic number increases by and the atomic mass remains the same. To identify the new element resulting from beta decay, simply balance the atomic numbers and masses. Beta particles are much smaller and higher energy than alpha particles. However, most solid objects can stop particles. They can penetrate skin, but not heavy clothing.
- radiation is best described as bundles of energy resembling light energy or x-rays. They move at the speed of light, but they don't have or electrical properties. Gamma radiation does not cause a change of nature of an element (atomic number, mass or charge) but it accounts for the energy changes that accompany radioactivity. Because the number of remains the same, gamma radiation from an element does not result in a different . Since gamma radiation is very high energy, they can pass through many objects. It takes a thick heavy shielding, such as , is necessary to shield gamma radiation.

b. Half-life. The stability of radioactive substances can be expressed in . This is the time it takes for of a sample of a radioactive element to change to another element. It can range from a billionth of a second to millions of . Since radium has a half-life of 1590 years, if you have 1 gram today, .5 gram will be left years from now. After another 1590 years, .25 grams will be left.