Humboldt State University ® Department of Chemistry

Richard A. Paselk

Chem 109 - General Chemistry - Spring 2015

Lecture Notes 5: 30 January

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Chapter 2: Atoms Molecules & Ions

Chemistry has a very long history, however progress was limited until quantitative measurements became the accepted norm for lab work in the 17th and 18th centuries. A fundamental observation was (and is) that mass is unchanged in chemical processes. This observation is summarized in the:

Law of Conservation of Mass:

Mass is neither created nor destroyed during a chemical change. (Strictly speaking there is no measurable change.)

For example, if we burn gasoline (octane) in air we will get carbon dioxide and water:

2 C8H18 + 25 O2 right arrow 16 CO2 + 18 H2O

If we were to weigh (determine the mass) of the octane and oxygen vs. the carbon dioxide and water we would find them to be identical - the masses are the same on both sides of the equation (that's why its called a chemical equation, the two sides are "equal"). Looked at another way, if you count the atoms, the numbers of each kind of atom on each side are identical - so we can also say that atoms are conserved in chemical processes.

This is really the fundamental assumption of chemistry, and thus the measurement of mass is the fundamental process underlying much of chemical work.

Once conservation of mass was accepted it was noted that combinations of various elements to form compounds nearly always result in the formation of compounds with constant proportions by mass. This set of observations is summarized in "Proust's Law," now known as the

Law of Definite Proportion:

a given compound always contains the same proportion by mass of its constituent elements.

But of course there is more than one way to combine many elements to give compounds. For these compounds a new set of observations applies:

The Law of Multiple Proportions:

When two elements combine to form a series of compounds, the ratio of the mass of the second element which will combine with 1 gram of the first will always be reducible to a small whole number. (Similarly, with multi-element compounds {other than macromolecules}, the ratios of the elements also reduce to small whole numbers.)

Dalton's Atomic Theory

These various laws imply that matter is made up of small discreet units, which we call atoms. The earliest "successful" theory of atoms is that of John Dalton (1803). Dalton's atomic theory states:

  1. All matter is composed of ultimately small particles, called atoms.
  2. Atoms are permanent and indivisible - they can neither be created nor destroyed.
  3. Elements are characterized by their atoms. All atoms of a given element are identical in all respects. Atoms of different elements have different properties.
  4. Chemical change consists of a combination, separation, or rearrangement of atoms.
  5. Chemical compounds are composed of atoms of two or more elements in fixed ratios.

All of these statements are close to reality, and nearly describe chemical behavior. But there are exceptions. Thus atoms can be created and destroyed via nuclear processes. They consist of different forms called isotopes. Atoms are not the smallest particles, etc.

Classical Characterization and Descriptions of Atoms

During the latter half of the 19th century we began to actually characterize what these "atoms" might be like. There were two main lines of experimental work:

When two electrodes are placed in an evacuated tube and a high voltage placed across them a stream of negatively charged particles flows between the electrodes (from the negative cathode to the positive anode). Since they come off of the cathode they are called cathode rays, and the evacuated tube a cathode ray tube (the ancestor of todays TV or computer monitor cathode ray tube or "CRT"). These rays can be diverted with magnets or charged plates outside the tube. In fact by careful manipulation 19th century physicists were able to determine the ratio of the mass to the charge of the cathode rays. (The cathode ray is also known as the electron.)

In the early 20th century Milliken determined the charge on the electron by watching the rate of fall of tiny oil drops between two charged plates. The droplets had been ionized (charged) by exposure to x-rays.

image of millikin oil drop apparatus

Image of Millikin's refined apparatus for the oil drop experiment (Millikin, 1917)

public domain image via Wikipedia

He noted that the droplets behaved like they had a multiple of some smallest charge on them, and determined this charge to be that of a single electron. Thus Milliken determined the charge on a single electron by making the assumption that the stepwise charge on the oil drops was due to single electron differences = -1.6 x 10-19 Coulomb. With Thompson's determination of the charge/mass ratio for cathode rays (electrons), the mass of the electron could be calculated as = 9.11 x 10-28g.

Goldstein used a modified Crook's tube with the cathode in the middle with a hole in the center and found positive rays, which he called "canal rays." These particles had positive charges in multiples of +1.6 x 10-19 C, but variable charge to mass ratios! Thus matter consists of electrons and positively charged particles to make neutral matter.

Thompson proposed an atom based on this information (c. 1890) called the "plum pudding model" in which the atom is a positively charged mass with negative electrons distributed through it like raisins in a pudding.

However, this model was destroyed by the next great bit of experimental evidence. Rutherford used a source of alpha-rays and aimed them at a thin foil of metal (1911). By the plum pudding model he expected that these particles would be slowed or deflected. He was very surprised to find that nearly all of the particles passed through the foil unobstructed, while some were deflected a lot, even being essentially reflected back to the source!

So what does this mean for the description of the atom? It must consist of mostly empty space, with the positively charged portion confined to a very small space, a nucleus, in the center.

Chemical Nomenclature

Nomenclature is covered in section 2.8 (pp 60-70) in your textbook-you should be able to do the examples and exercises in the suggested problems. Note also the Discussion Module, Sapling and these notes.

First, let's look at the the elements that you should learn the names of, as listed on the web:

HSU Chemistry Elements Names

The common ions and acids and bases are summarized on the handout and the web:

HSU Chemistry Table of Ions & HSU Chemistry Table of Acids

Note that formulae are more or less written with the elements ordered by electronegativity (elements on the right side precede those on the left).

Covalent vs. Ionic compounds:

This distinction will be important in some aspects of naming chemical compounds.

IUPAC vs traditional names

There are two common naming systems:

The IUPAC/Stock system

This is the modern, systematic scheme developed by the International Union of Pure and Applied Chemists.

Recognize these traditional names for metal ions:

FYI - Nomenclature examples etc.

Polyatomic ions

In these ions a group of atoms are covalently bound to each other and functions as a single charged particle. You should memorize the following names and be able to write formulas for the compounds and vice versa:

  • ammonium ion:
  • cyanide ion:
  • carbonate ion:
  • nitrate ion:
  • nitrite ion:
  • phosphate ion:
  • sulfite ion:
  • oxalate ion:
  • hydroxide ion:
  • acetate ion:
  • bicarbonate ion:
  • sulfate ion:

Compounds

Salts (Ionic Compounds) and Bases

Ions of opposite charge may combine to form neutral compounds. Thus the ions must combine in ratios such that the charges cancel. When the negative ion is hydroxide, the compound is considered a base. Examples:

  • potassium chloride:
  • sodium hydrogen sulfate (sodium bisulfate):
  • aluminum hydroxide:
  • iron(III) carbonate:
  • iron(III) hydrogen carbonate (iron(III) bicarbonate):
  • ammonium phosphate:
  • copper(I) oxalate:
  • calcium hydroxide:

Acids

Acids are compounds which give hydrogen ions (protons) in solution. There are two common inorganic acid types in terms of nomenclature:

  • Hydro-( )-ic acids: If hydrogen combines with a nonmetallic element the resulting acid is named by adding the prefix hydro- and replacing -ide by -ic. Examples:
    • HBr:
    • HCl:
  • Oxo acids: When non-metallic elements react with oxygen the resulting products often react with water or form ions which can react with protons to from acids. These oxo acids are named by replacing the -ate suffix with -ic acid or -ite suffix with -ous acid. Examples:
    • carbonic acid:
    • sulfuric acid:
    • acetic acid:
    • nitric:
    • Chloroacids (Hydrochlorous acid - Perchloric acid)

The HSU Chemistry Table of Common Acids is a useful summary of the acids you should be familiar with.

Special names

These compounds don't follow the rules, but have been in common use so long they keep their traditional names. Examples:

  • Water:
  • methane:
  • ammonia:
  • hydrogen peroxide:

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© R A Paselk

Last modified 30 January 2015