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CHEM 1411 General Course Guide - Lab

Early History of Chemistry

This video by AACT traces the history of chemistry from the discovery of fire, through the various metal ages, up to the great philosophers.

400 B.C. = Greeks proposed all matter was composed of 4 fundamental substances: water, earth, fire, air. They considered questions like:

  • Whether matter is continuous (infinitely divisible into smaller pieces), or
  • if matter is composed of small, indivisible particles.
    • Demokritos (or Democritus) of Abdera (c. 460 -c. 370 B.C.) and Leucippos believed in the latter and used the term atomos (atoms) to describe these ultimate particles.

The Greeks had no way of testing out their ideas, and thus no definitive conclusion could be reached about the divisibility of matter.

The next 2000 yrs. = Alchemists were believed to be mystics and fakes but were actually serious scientists who discovered several elements and learned to prepare the mineral acids. Alchemy is a pseudoscience also considered medieval chemical science.

Sixteenth century = The foundations of modern chemistry

  • Georg Bauer (1494-1555): A German scholar and "father of mineralogy" developed systematic metallurgy (extraction of metals from ores).
  • Paracelsus (Philippus Theophrastus Bombastus von Hohenheim [1493-1541]): A Swiss alchemist/physician developed the medicinal application of minerals.

Seventeenth and eighteenth centuries

  • Robert Boyle (1627-1691): The first "chemist" to perform quantitative experiments. He measured the relationship between pressure and volume of air. He published a book (The Skeptical Chymist) and then physics and chemistry were born. Boyle had the idea about chemical elements and stated that a substance was an element unless it could be broken down into two or more simpler substances. As elements became more accepted, the Greek system of four elements died.
  • Georg Stahl (1660-1734): A German chemist suggested that a substance burning in a closed container will stop burning because the air in the container becomes saturated with "phlogiston". This was later found out to be oxygen gas.
  • Joseph Priestley (1733-1084): An English clergyman and scientist discovered oxygen gas and was found to support vigorous combustion; thus, supposed to be low in phlogiston.


  • Oxygen was originally called "dephlogisticated air".
  • Oxygen is from the French oxygè​ne, meaning "generator of acid," because it was initially considered to be an integral part of all acids.
  • Oxygen gas was first observed by Karl W. Scheele (1742-1786), a Swedish chemist. His results were published after Priestley's and that is why Priestley is normally credited with the discovery of oxygen.

This video by CrashCourse explains how we went from alchemists to chemists, the law of conservation of mass, how we came to have a greater understanding of how chemical compounds work, and eventually a complete understanding of what atoms and molecules are.

Antoine Laviosier (1743-1794): A French chemist was finally able to explain the true nature of combustion.

  • In his experiments, he carefully weighed the reactants and products of different reactions and suggested mass is neither created nor destroyed.
  • Law of conservation of mass: Mass is neither created nor destroyed in a chemical reaction.
  • He discovered that combustion involved oxygen (he named), not phlogiston.
  • He discovered that life was supported by a process that also involved oxygen and was similar to combustion.
  • He wrote the first chemistry textbook, Elementary Treatis on Chemistry, but the French Revolution broke out and he was executed due to association with collecting taxes for the government (seen as an enemy of the people in 1794).

Joseph Proust (1754-1826): A French chemist who showed that a given compound always contains exactly the same proportion of elements by mass. 

Copper carbonate is always 5.3 part copper to 4 parts oxygen to 1 part carbon (by mass).

  • Law of definite proportion (originally called "Proust's Law): A given compound always contains exactly the same proportion of elements by mass.

John Dalton (1766-1844): An English schoolteacher considered if elements were composed of tiny individual particles, a given compound should always contain the same combination of those atoms. This also explained why the same relative masses of elements were always found in a given compound.

  • He discovered another principle that convinced him more of the existence of atoms.
  Mass of Oxygen That Combines with 1 g of Carbon
Compound I 1.33 g
Compound II 2.66 g


Carbon and oxygen form two different compounds that contain different relative amounts of carbon and oxygen.

Compound II contains twice as much oxygen per gram of carbon as compound I. (This can be explained in terms of atoms).

Compound I could be CO and compound II could be CO2


  • Law of multiple proportions: When two elements form a series of compounds, the ratios of the masses of the second element that combine with 1 g of the first element can always be reduced to small whole numbers.

For more information and problems concerning the law of multiple proportions, view Organic Chemistry Tutor's Law of Definite Proportions Chemistry Practice Problems - Chemical Fundamental Laws video.

Dalton's Atomic Theory

In 1808, Dalton published A New System of Chemical Philosophy and with it, his theory of atoms.

Dalton's Atomic Theory:

  1. Each element is made up of tiny particles called atoms.
  2. The atoms of a given element are identical; the atoms of different elements are different in some fundamental way(s).
  3. Chemical compounds are formed when atoms of different elements combine with each other. A given compound always has the same relative numbers and types of atoms.
  4. Chemical reactions involve reorganization of the atoms - changes in the way they are bound together. The atoms themselves are not changed in a chemical reaction.

In Dalton's time, the formula for water was unknown, but that it was composed of oxygen and hydrogen (8 g of oxygen present for 1 g of hydrogen).

Dalton's thought process:

  • If water was OH, then an oxygen atom would have 8 times the mass of a hydrogen atom. If water was H2O, then that would mean that each oxygen atom is 16 times as massive as each atom of hydrogen.
  • Since the formula was unknown, Dalton could not specify the relative masses of oxygen and hydrogen.
  • Dalton's assumption: Nature would be as simple as possible; thus making the formula for water OH. Hydrogen was given a mass of 1 and oxygen a mass of 8.

Dalton prepared the first table consisting of atomic masses (sometimes called atomic weights) that was later proved to be incorrect because of his assumptions on formulas of certain compounds.

Joseph Gay-Lussac (1778-1850):  A French chemist measured (under the same conditions of temperature and pressure) the volumes of gases that reacted with each other.

Amadeo Avagadro (1776-1856): An Italian chemist interpreted Gay-Lussac's results and proposed that at the same temperature and pressure, equal volumes of different gases contain the same number of particles (this assumption is called Avogadro's hypothesis).

  • In this type of condition, the volume of gas is determined by the number of molecules present and not by the size of the individual particles.

(a) One volume of chlorine gas reacted with one volume of hydrogen gas to produce two volumes of hydrogen chloride gas, and one volume of oxygen gas reacted with two volumes of hydrogen gas to produce two volumes of water vapor. (b) A summary of Avogadro’s hypothesis, which interpreted Gay-Lussac’s results in terms of atoms. Note that the simplest way for two molecules of hydrogen chloride to be produced is if hydrogen and chlorine each consist of molecules that contain two atoms of the element.

If this assumption is correct, then 2 molecules (a collection of atoms) of hydrogen react with 1 molecule of oxygen to produce 2 molecules of water.

These observations would then be assuming that gaseous hydrogen, oxygen, and chlorine are composed of diatomic (two-atom) molecules: H2O2, and Cl2.

This observation suggests that the formula for water is not OH like Dalton thought, but H2O.

Avogadro's interpretations were not accepted by most chemists and in the nineteenth century, measurements of masses for different elements that combined to form compounds were taken.

This resulted in a list of relative atomic masses.
Jöns Jakob Berzelius (1779-1848): A Swedish chemist contributed to this list by discovering cerium, selenium, silicon, and thorium. He developed modern symbols for the elements used in writing the formulas for the compounds.


There are seven diatomic molecules.
H2N2F2​, O2I2, Cl2Br2
An easy way to remember them is:
Have NFear OIce Cold Beverages

OR by the seven that is drawn on the periodic table + hydrogen



For more information on Dalton's Atomic Theory, visit Khan's Academy.

Early Experiments to Characterize the Atom

J. J. Thomson (1856-1940): An English physicist who studied electrical discharges in partially evacuated tubes called cathode-ray tubes.
Cathode rays: The "rays" emanating from the negative electrode (cathode) in a partially evacuated tube.

Since this ray was produced at the negative electrode and repelled by the negative pole, Thomson hypothesized that the ray was a stream of negatively charged particles, which are now called electrons.

In the left illustration, you are able to see the deflection of the cathode rays by an applied electric field as well as the charge-to-mass ratio of an electron that was determined by Thomson. e represents the charge on the electron in coulombs (C) and represents the electron mass in grams.

Thomson was trying to understand the structure of the atom with the experiments that he was conducting. He hypothesized that all atoms must contain electrons. He knew atoms were known to be electrically neutral and began to consider that an atom consisted of a diffuse cloud of positive charge with the negative electrons embedded randomly in it.

The illustration to the right is called the plum pudding model because the electrons are the raisins dispersed in a pudding (positively charged cloud) just like the English dessert. 

This video by Tyler DeWitt discusses the discovery of the electron by the cathode-ray tube experiment.

Robert Millikan (1868-1953): An American physicist conducted experiments that allowed him to determine the magnitude of the electron charge by involving charge oil drops.

The illustration on the left shows the schematic representation of the apparatus that was used to determine the charge of the electron. The fall of the charged oil droplets can be halted by the adjustment of the voltage across the two plates. The voltage and the mass of the oil droplet can then be used to calculate the charge.

Millikan calculated the mass of an electron is:
9.11 x 10-31 kg

For more information regarding Robert Millikan and the oil drop experiment, view Tyler DeWitt's video Charge of an Electron.

Henri Becquerel (1852-1908): A French scientist who, in 1896, accidentally discovered that a piece of mineral containing uranium could produce its image on a photographic plate in the absence of light. He credited this phenomenon to the spontaneous emission of radiation by the uranium, which he called radioactivity.

There are three types of radiation emission:

  • Gamma (γ) rays: High-energy "light"
  • Beta (β) particles: High-speed electron
  • Alpha (α) particles: A 2+ charge which is twice the charge of an electron & with the opposite sign
    • Mass of an α particle is 7300 times that of an electron

Ernest Rutherford (1871-1937): A New Zealand scientist who named the α and β particles and the γ ray. He coined the term half-life to describe an important attribute of radioactive elements. 

Rutherford was testing the accuracy of Thomson's plum pudding model which involved directing α particles at a thin sheet of metal foil. His hypothesis was that if Thomson's model was accurate, then the α particles should go through the thin foil with very minor deflections. 

This is an illustration of the experiment that Rutherford was using to prove the accuracy of Thomson's plum pudding model. There is a box that contains the source of α particles and shoots a beam towards a thin metal foil. A screen has been placed around the foil with a slight opening for the beam in order to detect the scattered α particles.


Instead of the particles passing through with only a minor interference in their path, the results showed that most of the α particles passed through the foil, but many of them were deflecting at large angles and reflecting (but not hitting the detector).

Due to these results, Rutherford knew that the plum pudding model was incorrect and the large deflections were caused by a positively charged concentrated center containing most of the atom's mass.

Nuclear atom: An atom with a dense center of positive charge (the nucleus) with electrons moving around the nucleus at a distance that is large relative to the nuclear radius.

This video by Tyler DeWitt explains more about Rutherford's Gold Foil Experiment.

Modern View of Atomic Structure

The atom consists of:

  • A nucleus (with a diameter of about 10-13 cm); and
  • Electrons which move about the nucleus at an average distance of about 10-8 cm from it
The Mass and Charge of the Electron, Proton, and Neutron
*The magnitude of the charge of the electron and proton is 1.60 x 10-19 C
Particle Mass Charge*
Electron 9.109 x 10-31 kg 1-
Proton 1.673 x 10-27 kg 1+
Neutron 1.675 x 10-27 kg None



The nucleus contains:

     Protons: A positively charged particle equal in                   magnitude to the electron's negative charge

     Neutrons: Virtually has the same mass as a proton,           but does not have a charge.

The nucleus is small in size compared with the overall size of the atom and it has an extremely high density.

Electrons make up most of the atomic volume of the atom and depending on the number and arrangement of them, it affects the atom's ability to interact with other atoms to form molecules. Due to this, atoms of different elements (which have different numbers of protons and electrons) show different chemical behavior thus resulting in the reason why atoms have different chemical properties.

The atomic number for sodium is 11 which means that it also has 11 protons, and since atoms have no net charge (they are neutral) it means that it also has 11 electrons.

  • Atomic number (number of protons): is written as a subscript

Atomic number = # of protons = # of electrons

  • Mass number (the total number of protons and neutrons) is written as a superscript

Mass number = # of protons + # of neutrons

Mass number → 23                                       
                             Na <-- Element Symbol
Atomic number → 11                                           

However, each sodium atom has neutrons in its nucleus, and different sodium atoms exist that have different numbers of neutrons.

  • Isotopes: Atoms with the same number of protons but different numbers of neutrons.

Since the chemistry of an atom is due to its electrons, isotopes show almost identical chemical properties and in nature, most elements contain a mixture of isotopes.

Isotopes of Hydrogen

Atomic Number 1 1 1
Number of Protons 1 1 1
Number of Neutrons 0 1 2

This video by Tyler DeWitt explains what isotopes are and how to write the atomic number and mass number in isotope notation.