image Atoms

Georgia Perimeter College

Objectives

  1. Demonstrate an understanding of the general features of the modern model of the atom, including protons, neutrons and electrons (their locations within the atom, their relative charges and masses).
  2. Demonstrate an understanding of the terms atomic number, mass number, atomic mass and (and the differences between these terms).
  3. Write the nuclide notation of an element and relate it to the mass number, atomic number, number of protons, neutrons and electron.
  4. Demonstrate an understanding of the term isotope.
This section addresses, in whole or in part, the following Georgia GPS standard(s):
  • S8P1. Students will examine the scientific view of the nature of matter.
    a. Distinguish between atoms and molecules.
    b. Describe the difference between pure substances (elements and compounds) and mixtures.

This section addresses, in whole or in part, the following Benchmarks for Scientific Literacy:
  • All matter is made up of atoms, which are far too small to see directly through a microscope. The atoms of any element are alike but are different from atoms of other elements. Atoms may stick together in well-defined molecules or may be packed together in large arrays. Different arrangements of atoms into groups compose all substances.
  • Scientific ideas about elements were borrowed from some Greek philosophers of 2,000 years earlier, who believed that everything was made from four basic substances: air, earth, fire, and water. It was the combinations of these "elements" in different proportions that gave other substances their observable properties. The Greeks were wrong about those four, but now over 100 different elements have been identified, some rare and some plentiful, out of which everything is made. Because most elements tend to combine with others, few elements are found in their pure form.

This section addresses, in whole or in part, the following National Science Education Standards:
  • A substance has characteristic properties, such as density, a boiling point, and solubility, all of which are independent of the amount of the sample. A mixture of substances often can be separated into the original substances using one or more of the characteristic properties.
  • Chemical elements do not break down during normal laboratory reactions involving such treatments as heating, exposure to electric current, or reaction with acids. There are more than 100 known elements that combine in a multitude of ways to produce compounds, which account for the living and nonliving substances that we encounter.
  • Matter is made of minute particles called atoms, and atoms are composed of even smaller components. These components have measurable properties, such as mass and electrical charge. Each atom has a positively charged nucleus surrounded by negatively charged electrons. The electric force between the nucleus and electrons holds the atom together.
  • The atom's nucleus is composed of protons and neutrons, which are much more massive than electrons. When an element has atoms that differ in the number of neutrons, these atoms are called different isotopes of the element.

Atoms

All of the matter around you is made of atoms.
Atoms are the smallest particle of matter with constant properties.

All atoms are comprised of three types of subatomic particles:

  1. Protons
  2. Electrons
  3. Neutrons

All protons are exactly the same.
All neutrons are exactly the same.
All electrons are exactly the same.

Protons have a positive charge (+).
Electrons have a negative charge (-).
Neutrons do not have any charge. They are electrically neutral.

The nucleus of atom contains neutrons and protons.

Electrons orbit the nucleus.
Electrons are outside of the nucleus in electron shells that have different shapes at different distances from the nucleus.

In a stable atom, the number of electrons = the number of protons.

Mass of Subatomic Particles

Almost all the mass of an atom is concentrated in the nucleus.

Protons and neutrons have almost exactly the same mass.
The mass of a proton or neutron is 1.66 E -24 grams (1.66 x 10-24 g) or one atomic mass unit (AMU).

Electrons have a mass that is much less than that of the other particles (only about 1/1835 the mass of a proton). The mass of an electron is 9.05 E -28 grams (or 9.05 x 10-28 g). This number is a billionth of a billionth of a billionth of a gram.

Note: It is not possible for anyone or any machine that uses light to see a proton. The wavelength of light is too large to be able to detect anything as small as a proton.

The atom is mostly empty space.
Ernest Rutherford shot subatomic particles at a very thin piece of gold.
Most of the particles went straight through the gold. It was like shooting a rifle into a thin line of trees.
Some of the particles bounced off, some stuck inside, but the major portion of them passed through the gold foil.
By Rutherford's calculations, the nucleus in an atom is like a B-B in a boxcar.

Elements

There are only a few more than one hundred elements. Of those, only 83 are not naturally radioactive, and of those, only 50 or so are common.

Know the common elements well enough so that if you read or hear about one of them, you instantly know what it is.

Learn how to spell the names of the elements. Learn the symbols. Some of the symbols have one letter, some have two, but each element symbol has one and only one upper case letter in it.

Common Elements

You should know the name and symbol for the following elements. If you see the name, you should know the symbol. If you see the symbol, you should know the name. For some elements, there are other names, sometimes Latin, from which the element symbol was derived, which are listed in parentheses.

Helium He Lithium Li Hydrogen H Sodium (Natrium) Na
Boron BCarbon C Silicon Si Calcium (Lime) Ca
Beryllium Be Fluorine F Neon Ne Sulfur (Brimstone) S
Phosphorus PNitrogen N Aluminum Al Potassium (Kalium) K
Chlorine Cl Argon Ar Magnesium Mg Iron (Ferrum) Fe
Bromine BrOxygen O Manganese Mn Copper (Cuprum) Cu
Cobalt Co Nickel Ni Chromium Cr Lead (Plumbum) Pb
Zinc Zn Krypton Kr Rubidium Rb Silver (Argentum) Ag
Iodine I Platinum Pt Cadmium Cd Tin (Stannum) Sn
Cesium Cs Barium Ba Francium Fr Antimony(Stibium) Sb
Bismuth Bi Arsenic AsStrontium Sr Tungsten(Wolfram)W
Radon Rn Xenon Xe Polonium Po Gold (Aurum) Au
Radium Ra Uranium UMercury (Hydrargyrum) Hg

History of Development of the Model of the Atom

  1. Dalton's Atomic Theory (1808): This was the first atomic theory that was based on experiments.

    a. Each element is made up of tiny indivisible particles called atoms (* actually, atoms are divisible; they can be broken up into smaller particles -- some of these smaller particles are discussed below)

    b. Atoms of a given element are identical; atoms of different elements are fundamentally different. (There may be some differences between atoms of the same element; we will discuss these differences below under isotopes.)

    c. Chemical compounds are formed when atoms combine with each other. A given compound always has the same relative numbers and types of atoms.

    d. Chemical reactions involve the reorganization of the atoms, changes in the way they are bound together. The atoms themselves are not changed in a chemical reaction. (This is true except for nuclear reactions, but Dalton didn't know about nuclear reactions at the time.)

  2. Discovery of the electron: by J. J. Thompson, 1897; electron is a negatively charged particle with very little mass.
  3. Discovery of the proton: by Ernest Rutherford, 1911; a particle in the nucleus (center) of the atom, has a positive charge equal in magnitude (size) to the negative charge of the electron, but is nearly 2000 times the mass of the electron.
  4. Discovery of the neutron: by James Chadwick, 1932; a particle in the nucleus, has no charge and a mass nearly equal to that of the proton.

Atomic Number

The atomic number is the number of protons in the nucleus of an atom. It is listed on the periodic table for each element. No two elements have the same atomic number (or the same number of protons), so the atomic number identifies the element.

: (symbol: Z)

Mass Number

Mass number: (symbol: A) total number of protons and neutrons in the nucleus (not listed on the periodic table, since it varies)

Atoms of the same element have the same atomic number, but may have different mass numbers.

Shorthand symbol notation for a particular atom (also called nuclide symbol notation):

AE (E = element's symbol; A = mass number; Z = atomic number)
Z

For example:

23Na
11
represents a sodium atom which always has 11 protons and in this case has a mass number of 23.
(Note: This means that there are 12 neutrons. 23 - 11 = 12)

Mass number - atomic number = # neutrons

Practice question:
Write the symbol notation for the atom which has an atomic number of 9 and a mass number of 19.
How many protons and neutrons does this atom contain?
If this is a neutral atom, how many electrons does this atom contain?
What if this atom has a +2 charge -- how many electrons does it have?

ANSWERS

Isotopes

Two atoms of the same element which have different numbers of neutrons are called different isotopes of that element. An isotope is an atom whose nucleus contain the same number of protons but a different number of neutrons.

Atomic Mass (Atomic "Weight") vs. Mass Number

The atomic mass (or atomic "weight") listed on the periodic table is the weighted average of the masses of all naturally occurring isotopes of an element in atomic mass units (amu).

Definition: 12 amu (exactly) = the mass of one atom of a 12C isotope; all other atomic masses are defined relative to this.

"Weighted average" means that the masses of each of the isotopes are "weighted" differently in calculating the average mass (just as quizzes might be weighted differently from tests or final exams in calculating your average grade). Masses of isotopes that are more abundant are weighted more heavily in the calculation of the average atomic mass for an element than masses of isotopes that are less abundant.

Example: There are four naturally occurring isotopes of iron. Below are listed the natural percent abundances and the masses of each of these isotopes:

Isotope: % Abundance: Mass (amu):
54Fe
26

56Fe
26

57Fe
26

58Fe
26

5.82

 

91.66

 

2.19


0.33

53.94

 

55.935

 

56.935


57.933


The atomic mass listed on the Periodic table would be calculated using the "weighted average" method. So we wouldn't simply add up the 4 masses (last column) and divide by 4, because that wouldn't be taking into account the different numbers of isotopes present in a normal sample of iron. Adding up the 4 masses and dividing by 4 counts each of the masses equally (25 % of the average, in this case). By the "weighted average" method, the masses of isotopes that are more abundant are weighted more heavily in the calculation of the average than the less abundant isotope masses:

0.0582(53.94) + 0.9166(55.935) + 0.0219(56.935) + 0.0033(57.933) = 55.85 amu

The first isotope's mass of 53.94 amu is weighted as only 5.82% (0.0582) of the average mass, etc.

Don't worry about having to do these calculations; the results are already listed in the Periodic Table for you, but it's nice to understand where those values come from.


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Content provided by Dr. Margaret Venable, Dr. Michael Denniston, and and Dr. David Wilner.

Some of this information is used with permission of Chemtutor.

Page created by Pamela J.W. Gore
Georgia Perimeter College,
Clarkston, GA

Page created November 16, 2006