Before we're done with this program we will have learned about the laws of chemistry and how they led an opinionated schoolmaster named John Dalton to formulate a new theory of atoms based on the quantitative laws. We will learn about the law of combining volumes and the ensuing controversy which was finally resolved after Dalton's death by Avagadro's hypothesis with an elegance we have come to expect.
At the end of the program we will turn our attention to chemical symbols and equations.
1.1.1. Name and describe the meaning of the three fundamental laws of chemistry?
1.1.2. Briefly describe the controversy surrounding Berthollet and Proust.
1.1.3. Compare and contrast the Law of Constant Proportions and the Law of Multiple Proportions.
1.2.1. Why did the Greek philosophers reject the original atomic theory of Democritus?
1.2.2. Discuss the relationship between the laws of chemistry and Dalton's atomic theory.
1.2.3. What is the difference between an atom and an element? Between an atom and a molecule?
1.2.4. Discuss the law of combining volumes and the controversy surrounding its acceptance.
1.2.5. What is Avagadro's hypothesis and how did it reconcile the law of combining volumes with Dalton's atomic theory?
1.2.6. How does the modern concept of atoms differ from that of the ancient Greeks?
1.2.7. How do the laws of chemistry allow us to determine the relative atomic weights of the elements?
1.2.8. What is a gram molecular weight?
1.2.9. Use the periodic table to find the chemical symbols for the following elements: gold, silver, lead, potassium, sodium, silicon
1.2.10. How does the law of conservation of mass help in balancing chemical equations?
1.2.11. Why can we change the coefficients but not the subscripts in a chemical formula.
1.2.12. How much hydrogen will be required to completely react with 280 grams of nitrogen gas. How much ammonia will this reaction produce?
1.1.1. combinations of elements create the rich variety of substances which compose the universe and life
1.1.2. combining properties of elements required the assumption of discrete particles of matter
1.2.1. discovered that general rules apply
1.2.2. discovered that weights of chemical are important
1.2.3. showed how to recognize elements
1.2.4. showed that mass was not lost or gained
2.1.1. Total weight of products equals total weight of reactants
2.1.2. Mass is neither created nor destroyed in chemical reactions
2.1.3. 1000 grams of lead combines with 77 grams of oxygen to form 1077 grams of yellow lead oxide
2.1.4. 1000 g Pb + 77 g O --> 1077 g PbO
2.2.1. Proust vs. Berthollet
126.96.36.199. each announced contradictory laws around 1800
188.8.131.52. Proust: elements combine to form compounds in a fixed proportion by weight
184.108.40.206. Berthollet: elements combine to form compounds in variable ratios
220.127.116.11. Berthollet heated copper and tin (separate experiments) to form what seemed to him to be a continuous series of compounds of varying composition
18.104.22.168.1. cited solutions, alloys and glasses as similar models
22.214.171.124.2. Proust was able to show that Berthollet was analyzing impure compounds
2.2.2. Relative amounts of elements is always the same in a given compound
2.2.3. 1000 g Pb + 77 g O --> 1077 g PbO
126.96.36.199. excess of either lead or oxygen simply will not combine
188.8.131.52. will remain in original form
2.2.4. 2000 g Pb + 154 g O --> 2154 g PBO
2.2.5. Pb/O = 1000/1077 = 2000/2154 = 93%
2.2.6. composition of yellow lead oxide is always 93% lead and 7% oxygen
If two elements combine to form two or more different compounds, then a simple ratio exists between the two weights of one element that can combine with a fixed weight of the other
2.3.1. Proust demonstrated that Berthollet was seeing various mixtures of two separate compounds of copper and two of tin
184.108.40.206. heating in closed container with insufficient air or in open container for insufficient time
2.3.2. Law was established by Dalton around 1800
2.3.3. two elements may combine to form more than one compound. Each has a unique fixed percentage of the combining elements
2.3.4. the ratio of the weights of two elements which combine with the same element is always an integer
2.3.5. 1000 grams of lead will also react with 154 grams of oxygen to form 1154 grams of brown lead oxide
220.127.116.11. 154/77 = 2
18.104.22.168. lead will react with oxygen in two ways
22.214.171.124. one way uses exactly twice as much oxygen for a given amount of lead
2.3.6. other examples
126.96.36.199. 16 g of oxygen combines with 12 g of carbon to form carbon monoxide, 32 g of oxygen combines with 12 g of carbon to form carbon dioxide. 32/16 = 2
188.8.131.52. black oxide of copper: 16 g oxygen combines with 63 g copper, red oxide of copper 16 g of oxygen combines with 126 g of copper. 126/63 = 2
184.108.40.206. compounds of nitrogen and oxygen
4.5.1. All matter is made of atoms
220.127.116.11. revived Democritus' idea
18.104.22.168. atoms are indivisible
22.214.171.124. atoms are the building blocks of substances
126.96.36.199. united concept of atom and element
4.5.2. All atoms of the same element have the same weight
188.8.131.52. have the same weight, size, behavior
184.108.40.206. Atoms of different elements have different weight, size, behavior
220.127.116.11. if different atoms of the same element had different weight, how could you explain the law of constant proportions
4.5.3. Atoms are rearranged during chemical reactions, never created or destroyed
18.104.22.168. all atoms present before the reaction are also present after the reaction
22.214.171.124. weight of reactants equals weight of products
126.96.36.199. law of conservation of mass
4.5.4. When elements combine to form compounds each group of atoms is identical
188.8.131.52. law of definite proportions
184.108.40.206. all molecules of a compound are identical
220.127.116.11. the ratio of atoms in a given compound is constant
18.104.22.168. each compound has a fixed chemical "formula"
4.5.5. Different groupings of the same atoms are possible
22.214.171.124. law of multiple proportions
126.96.36.199. the same elements might combine in more than one way
5.1.1. volumes of gases are also important in chemical reactions
188.8.131.52. volumes standardized to STP (standard temperature and pressure)
184.108.40.206. Gay-Lussac also established law of proportionality of volume and temperature for gases
220.127.116.11. not important in solids and liquids because they have fixed volume
5.1.2. 1 volume of nitrogen and 1 volume of oxygen react to form 2 volumes of nitrogen oxide
5.1.3. 2 volumes of hydrogen and 1 volume of oxygen react to form 2 volumes of water vapor
5.5.1. N + O --> NO, ie. two volumes produce one volume
5.5.2. N + 3H --> NH3, ie four volumes produce one volume
5.5.3. argued that Gay-Lussac must be wrong
18.104.22.168. cast aspersions on Gay-Lussac's experimental integrity
22.214.171.124.1. Gay-Lussac was known as a careful experimenter, Dalton as a "coarse" experimenter
5.6.1. Dalton believed the atoms were surrounded by caloric, so how could 3 or 4 volumes be compressed into two?
5.6.2. There was no reliable model for the gaseous state of matter. Kinetic theory of gases was 40 years away
Amadeo Avagadro (1776 - 1857)
Italian of noble birth, practiced law but abandoned it to teach physics at the University of Turin due to his love of mathematics and physics.
6.1.1. Required a modification of Dalton's atomic theory.
6.2.1. particles = smallest unit of matter present in gas
6.2.2. may be molecules or atoms
6.3.1. NN + OO --> NO + NO
6.3.2. HH, HH + OO --> HOH + HOH
6.3.3. NN + HH, HH, HH --> NHHH + NHHH
6.5.1. Dalton's authority was a strong factor
6.5.2. static theory prevailed which pictured atoms or molecules of a gas in contact with each other like fluffy balls of wool packed loosely in a crate. Equal volumes of gases having equal numbers of molecules required widely separated molecules
6.5.3. cause of gas pressure was thought to be repulsion of one atom of gas by another of like kind
6.5.4. gas atoms were thought to be surrounded by thick shells of self repulsive caloric
6.5.5. another good example of a theory ahead of its time
6.5.6. see how the concept of caloric, a paradigm of the times, caused the rejection of a god explanation
6.6.1. the amount of a substance which weighs one gram atomic or gram molecular weight
6.6.2. Avagadro's number (6x1023) of anything
8.1.1. Matter is composed of atoms
8.1.2. Each type of atom has a certain weight and is identical to all other atoms of the same type
8.1.3. Atoms combine in certain proportions to form compounds
8.2.1. shorthand notation for names of elements
8.2.2. convenient for discussing chemicals and reactions
8.2.3. most are logical or mnemonic derivations of name of element
126.96.36.199. Aluminum: Al
188.8.131.52 Carbon: C
184.108.40.206 Magnesium: Mg
8.2.4. several are derived from Latin or Greek name
220.127.116.11. iron: Ferrum
18.104.22.168. sodium: Natrium potassium: Kallium
22.214.171.124. copper: Cuprum
126.96.36.199. lead: Plumbum (Pb)
8.2.5. certain conventions adopted
188.8.131.52. each symbol is capital letter followed by small letter
184.108.40.206.1. some early discovered elements keep only first letter
220.127.116.11. first letter of symbol is capitalized, but element name is not
18.104.22.168. Use the symbol with no subscript to designate the uncombined element
22.214.171.124.1. for solids, most liquids, inert gases
126.96.36.199.2. Al, Fe, Na, Hg, Ne
188.8.131.52.3. Diatomic Gases
184.108.40.206.3.1. H2 O2 Cl2N2 F2Br2
8.3.1. Compiles all known chemical elements
8.3.2. shows atomic number, atomic weight, atomic symbols
8.3.3. organization allows prediction of chemical reactions
8.4.1. compounds are substances composed of more than one atom
8.4.2. ratios of atoms are fixed by chemical properties of the elements
8.4.3. ratios are determined in the laboratory by chemical analysis
8.4.4. atomic symbols are used to show kinds and numbers of atoms
220.127.116.11. is a molecule of water containing two atoms of hydrogen and one atom of oxygen
18.104.22.168. is two molecules of water, each containing two atoms of hydrogen and one atom of oxygen
8.5.1. rearrangements of atoms to form new compounds
8.5.2. mass is conserved
8.5.3. example: hydrogen gas plus oxygen gas produces water
8.6.1. shorthand for describing chemical reactions
8.6.2. iron reacts with sulfur to form iron sulfide
22.214.171.124. Fe + S --> FeS
8.6.3. compound names are simple if only two elements are involved
126.96.36.199. naming is mostly consistent (oxide, sulfide, chloride, nitride)
188.8.131.52. common names given preference ( water vs. dihydrogen oxide)
8.6.4. (aq)aqueous solution, (s) solid, (g) gas added to atomic symbol indicate state of substance
8.7.1. includes amounts of each substance
8.7.2. number of atoms must balance on both sides of reaction
8.7.3. cannot change subscripts because they represent the fixed chemical formula
8.7.4. can change coefficients since they represent the number of molecules
8.8. Amounts of Substance Involved in Reactions
8.8.1. Chemical formulas show the numbers of atoms
8.8.2. We count atoms by weighing them
8.8.3. Molecular weight is the weight of all atoms in a molecule
8.8.4. Gram molecular weight (GMW) is the weight in grams which equals the molecular weight
8.8.5. A gram molecular weight of any substance contains the same number of atoms or molecules
8.8.6. a GMW of water weights 18 grams and contains the same number of molecules as a GMW of iron (56 grams)
8.8.7. 4 g of hydrogen plus 32 g of oxygen yields 36 g of water
8.8.8. Given 16 grams of hydrogen, how much oxygen will it combine with and how much water will be formed?
184.108.40.206. each "recipe" uses 4 g of hydrogen
220.127.116.11. 16 g is 4 "recipes"
18.104.22.168. 4 "recipes" of oxygen is 4 x 32 g or 128 g
22.214.171.124. 4 "recipes" of water is 4 x 36 g or 144 g.
9.1.1. mostly systematic, name tells composition
9.1.2. a few rules will help to understand most substances
9.1.3. organic substance are more complicated
9.2.1. has the suffix "-ide" for the second element
9.2.2. metals are written first, nonmetals second
9.2.3. hydrogen chloride
9.2.4. sodium chloride
9.2.5. carbon dioxide
9.2.6. potassium iodide
9.3.1. SO4 is sulfate, SO3 is sulfite
9.3.2. NO3 is nitrate, NO2 is nitrite
9.3.3. CO3 is carbonate
9.3.4. PO4 is phosphate
9.3.5. ClO3 is chlorate
9.3.6. OH is hydroxide
9.3.7. NH4 is ammonium
9.4.1. suffixes "-ic" and "-ous" attached to first element
9.4.2. "-ous" given to element with largest proportion
9.4.3. mercurous oxide Hg2O, mercuric oxide HgO
9.4.4. ferrous oxide Fe2O3, ferric oxide FeO
9.5.1. CO2 is carbon dioxide, CO is carbon monoxide
9.5.2. SO3 is sulfur trioxide
9.5.3. CCl4 is carbon tetrachloride
In this lesson we have seen how quantitative measures of chemicals in reactions led to basic laws, know today as The Laws of Conservation of Mass ,Definite Proportions, and Multiple Proportions . Dalton's atomic theory provides a rational and ingenuous explanation for the observed laws. Although Dalton's original version contained some errors, our modern atomic theory recognizes a few other basic laws which make the theory consistent. One of these is the Law of Combining Volumes, stated by Joseph Guy- Lussac, which notes that the volumes of products and reactants of reacting gases always occur in small integer ratios. Avagadro added the idea that equal volumes of gases contain equal numbers of molecules. Avagadro noted that certain elemental gases are diatomic.
With these laws chemists of the nineteenth century were able to calculate atomic weights very accurately by measuring the weight ratios of elements which reacted with either oxygen or hydrogen.
Further advances in chemical knowledge occurred when the modern system of chemical symbols was combined with a shorthand for describing chemical reactions. These equations provide an easy way to predict the weights of substances undergoing chemical reactions. These equations help us to understand many things about chemical reactions and properties, including the ability to predict the amounts of chemicals used in various reactions. They also help us to understand chemical changes in the world around us.
The large number of chemical substances which are know to us requires a system of naming. For simple compounds the naming is straightforward, but it works well even for the most complicated substances.
The confirmation of the atomic theory led to more questions about the nature of matter, specifically the relationship between the physical properties of atoms. Since matter is made of atoms, then atoms themselves should obey Newton's laws.