Humboldt State University ® Department of Chemistry

Richard A. Paselk

Chem 109 - General Chemistry - Spring 2013

Lecture Notes 4: 30 January


Matter, cont.

Matter has both physical properties and chemical properties. These are properties which do not depend on the quantity of substance and therefore they can be used to identify a substance (sometimes referred to as intensive properties).

States of Matter. Matter can exist in three states under earth-surface conditions:

A fourth state of matter commonly occurs under special conditions: a plasma. A plasma is an ionized fluid - can be contained by magnetic fields.


Density is defined as the mass of a given volume of a substance: Density = mass/volume. Note that this weeks laboratory exercise give practice in Density, significant figures etc.

Let's try some density problems. First note that the units of density are g/cm3 or

Known: Density = mass/volume, generally expressed as g/mL = g/cm3

Solve: (35.987 g) / (20.0 mL) = 1.79935 g/mL

note that the units are those of density so we are confident we set it up correctly.

How about sig figs? Use multiplication/division rules, so count: 3 for 20.0 and 5 for 35.987, therefore should have three sig figs:

1.79935 g/mL = 1.80 g/mL

Extra Example: Using a jewelers balance a student found that a coin weighed 2.34 carats in air. By weighing it again submerged in water she found it had a volume of 0.034 mL. What is its density? (1 carat = 200 mg, defined)*

Known: 1 carat = 200 mg (defined), density is g/mL

Solve: (2.34 carats)(200 mg/carat)(1 g/1,000 mg) / 0.034 mL = 13.764706 g/mL

How about sig figs? Both conversion factors are defined, so exact. Two measurements: 2.34 and 0.034 = 3.4 x 10-2. Thus the answer will have only two sig figs since using counting rule - least number of sig figs.

13.764706 g/mL = 14 g/mL


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 carbon 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.

Another fundamental concept is that of elements. Elements are substances which cannot be broken down further into simpler substances by chemical or (non-nuclear) physical means.

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. 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 a-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.


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© R A Paselk

Last modified 28 January 2013