Science 122 Program 25 Atomic Theory

©1998 RCBrill. All rights reserved

Laws of Chemistry &Atomic Theory
Program 25
Lesson 4.3

Text References

Speilberg & Anderson, none

Booth & Bloom, 234-245

Coming Up

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. Introduction

2. Laws of Chemistry

3. John Dalton

4. Dalton's Atomic Theory

5. Law of Combining Volumes

6. Avagadro's Law

7. Relative Atomic Weights

8. Chemical Symbols & Equations

9. Naming Compounds

10. Summary


1.1. Laws of Chemistry

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. Atomic Theory

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. State five laws of chemistry and explain them using examples.
1.2. Correlate the hypotheses of Dalton's atomic theory with the laws of chemistry
1.3. State the law of combining volumes and cite some examples
1.4. Describe the controversy surrounding the law of combining volumes
1.5. Use Avagadro's law to reconcile atomic theory and the law of combining volumes
1.6. Show how the laws of chemistry can be used to determine the relative weights of atoms
1.7. Demonstrate knowledge of the rules for naming simple chemical compounds
1.8. Demonstrate understanding of the symbolic representation of chemical elements and compounds
1.9. Use the law of of conservation of mass to balance simple chemical equations

1. Introduction

1.1. world would be dull and lifeless if composed only of uncombined elements

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. Lavoisier set the stage for determining laws of chemistry

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

1.3. Other chemists worked to discover other quantitative laws
1.4. Laws led to atomic theory

2. Laws of Chemistry

2.1. Conservation of Mass

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. Definite Proportions

2.2.1. Proust vs. Berthollet each announced contradictory laws around 1800 Proust: elements combine to form compounds in a fixed proportion by weight Berthollet: elements combine to form compounds in variable ratios Berthollet heated copper and tin (separate experiments) to form what seemed to him to be a continuous series of compounds of varying composition cited solutions, alloys and glasses as similar models 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 excess of either lead or oxygen simply will not combine 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

2.3. Multiple Proportions

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 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 154/77 = 2 lead will react with oxygen in two ways one way uses exactly twice as much oxygen for a given amount of lead

2.3.6. other examples 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 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 compounds of nitrogen and oxygen

3. John Dalton (1766-1844)

3.1. Quaker schoolmaster
3.2. regularly attended Manchester Philosophical and Literary Society meetings
3.3. heard many scientists discuss controversial laws of chemistry
3.4. controversial because there was no explanation for the laws
3.5. heard concrete evidence from many experiments involving combining proportions of atoms
3.6. Dalton saw that laws of chemistry were best explained by assuming the existence of atoms with certain properties
3.7. New System of Chemical Philosophy (1808)
3.8. explained laws of chemistry with revived atomic theory

4. Dalton's Atomic Theory (1808)

4.1. elements are collections of only one type of atom
4.2. compounds are made of molecules which are combinations of atoms
4.3. atoms of a given element are all the same, elements of different elements are different
4.4. weight of atoms distinguished one from another
4.5. Details of Theory

4.5.1. All matter is made of atoms revived Democritus' idea atoms are indivisible atoms are the building blocks of substances united concept of atom and element

4.5.2. All atoms of the same element have the same weight have the same weight, size, behavior Atoms of different elements have different weight, size, behavior 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 all atoms present before the reaction are also present after the reaction weight of reactants equals weight of products law of conservation of mass

4.5.4. When elements combine to form compounds each group of atoms is identical law of definite proportions all molecules of a compound are identical the ratio of atoms in a given compound is constant each compound has a fixed chemical "formula"

4.5.5. Different groupings of the same atoms are possible law of multiple proportions the same elements might combine in more than one way

4.5.6. Dalton's Symbols


5. Law of Combining Volumes (Gay-Lussac's Law)

5.1. volumes of reactants are always in a simple ratio

5.1.1. volumes of gases are also important in chemical reactions volumes standardized to STP (standard temperature and pressure) Gay-Lussac also established law of proportionality of volume and temperature for gases 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

5.1.4. 1 volume of nitrogen and 3 volumes of hydrogen react to form 2 volumes ammonia

5.2. volume of products is always two regardless of volumes of reactants
5.3. Gay-Lussac did not explain the results
5.4. Joule said it was one of the most important discoveries made in physical science
5.5. Dalton's rules suggested contrary results

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 cast aspersions on Gay-Lussac's experimental integrity Gay-Lussac was known as a careful experimenter, Dalton as a "coarse" experimenter

5.6. Dalton had worked with the gases on a weight basis and had worked out the relative weights of the atoms, but Gay-Lussac was convinced that the results are unrelated to weights, but rather a characteristic of the gaseous state

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

6. Avagadro's Law

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. resolved dispute in simple, elegant manner

6.1.1. Required a modification of Dalton's atomic theory.

6.2. equal volumes of gases at equal temperature and pressure contain the same number of particles

6.2.1. particles = smallest unit of matter present in gas
6.2.2. may be molecules or atoms

6.3. some elemental gases are diatomic N2, O2, Cl2, F2 (consist of collection of molecules rather than individual 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

6.4. equal numbers of molecules in products and reactants
6.5. idea not accepted until resurrected in 1858

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. The Mole

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

7. Relative Atomic Weights

7.1. correct formulas and combining volumes allowed calculation of relative weights of atoms
7.2. weight oxygen : weight hydrogen = 8 : 1, volume hydrogen : volume oxygen = 2 : 1
7.3. can only be true if each oxygen atoms weights 16 times as much as each hydrogen atom. The sequence of illustrations below shows the logic.





7.4. atomic weight scale later refined and modified to use carbon-12 as base weight
7.5. The relative weight of any atom that combines with either oxygen or hydrogen can be determined
7.6. Most atoms will combine with either hydrogen or oxygen or both

8. Chemical Symbols and Equations

8.1. Atomic Theory

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. Atomic Symbols

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 Aluminum: Al Carbon: C Magnesium: Mg

8.2.4. several are derived from Latin or Greek name iron: Ferrum sodium: Natrium potassium: Kallium copper: Cuprum lead: Plumbum (Pb)

8.2.5. certain conventions adopted each symbol is capital letter followed by small letter some early discovered elements keep only first letter first letter of symbol is capitalized, but element name is not Use the symbol with no subscript to designate the uncombined element for solids, most liquids, inert gases Al, Fe, Na, Hg, Ne Diatomic Gases H2 O2 Cl2N2 F2Br2

8.3. Periodic Table

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. Compounds

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
8.4.5. examples is a molecule of water containing two atoms of hydrogen and one atom of oxygen is two molecules of water, each containing two atoms of hydrogen and one atom of oxygen

8.5. Chemical Reactions

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. Chemical Equations

8.6.1. shorthand for describing chemical reactions
8.6.2. iron reacts with sulfur to form iron sulfide Fe + S --> FeS

8.6.3. compound names are simple if only two elements are involved naming is mostly consistent (oxide, sulfide, chloride, nitride) 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. Balancing equations

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
8.7.5. practice

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? each "recipe" uses 4 g of hydrogen 16 g is 4 "recipes" 4 "recipes" of oxygen is 4 x 32 g or 128 g 4 "recipes" of water is 4 x 36 g or 144 g.

9. Naming Compounds

9.1. general statements

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. Compounds composed of two elements

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. Certain groups of atoms have a tendency to stay grouped

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. Two elements which form more than one compound

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. Prefixes such as mono-, di-, tri-, tetra-, penta- given to second element

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

9.6. Other more complicated systems exist, but this is enough for most simple compounds.

10. Summary and Conclusions

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.

10.1. Atomic and Newtonian Paradigms

10.1.1. Can the two paradigms be reconciled?
10.1.2. Atoms are Newtonian particles
10.1.3. Atoms have mass, so do they obey laws of mechanics?
10.1.4. How is heat connected with the motion of particles?
10.1.5. How does heat transfer take place between molecules?

10.2. HEAT as a form of energy is the link, as we will see in the next lesson, Kinetic Theory