Text References

Speilberg & Anderson, none
Booth & Bloom 421-445; 451-456

Coming Up

Before we are done with this program we will have summarized the advances in chemistry in the nineteenth century culminating in the Periodic Table of the Elements. We will have seen how the structure of the periodic table reflects the quantum nature of electron patterns surrounding atoms. We will have seen why chemical bonds form, and how this allows for the formation of several types of chemical bonds, explains the unusual properties of water, and explains the properties of electrolytes, acids, bases, and salts.

Questions	Objectives	1. Introduction	2. Nineteenth Century Chemistry
3. The Periodic Table	4. Electron Patterns	5. Chemical Bonding	6. Types of Chemical Bonds
7. Electrolytes, Acids, Bases, & Salts	8. Important Chemical Compounds	9. Unusual Properties of Water	10. Summary & Conclusions

Questions

1. List some of the advances in Chemistry and Physics in the nineteenth century
2. Describe and distinguish between chemical and physical properties of the elements.
3. What is the meaning of the terms "triads" and "octets" of elements?
4. Why was the associations of octets of elements initially ignored by the scientific community?
5. What made Mendeleev's periodic table better than Meyer's?
6. Why do we say that the periodic table "represents a significant leap in understanding of matter?"
7. What are the noble gases, and why did they remain undiscovered for such a long time?
8. What is the difference between the substances H and H2?
9. Compare and contrast the composition and properties of the following pairs of compounds:

1. water - ammonia
2. sodium chloride - potassium nitrate
3. methane - alcohol
4. methane - acetylene
5. alcohol - sugar
6. sulfuric acid - hydrochloric acid
7. sugar - sodium chloride

10. In what ways is water an unusual substance?
11. Define the following properties of water and note the importance of each.

1. liquidity
2. solvent ability
3. temperature - density relationships
4. specific heat
5. latent heat
6. surface tension

12. Define the term electrolyte and comment on the types of substances which have those properties.
13. Define the following: atomic number; atomic mass; ion
14. Discuss the statement: "Gain or loss of particular number of electrons is characteristic of a given atom."
15. What are isotopes?
16. Why are the numbers 2, 8, 18 "magic" numbers with respect to atoms?
17. Compare and contrast electron orbits around an atomic nucleus with planetary orbits around the sun.
18. Compare the properties of metals and nonmetals and their location in the periodic table.
19. What is meant by the terms "representative elements" and "transition elements"?
20. Why does it matter whether valence electrons in atoms are paired or unpaired?
21. What does it mean to say an atom "tends to achieve the octet?"
22. How can the periodic table be used to predict the reactions of representative elements?
23. What is ionic bonding and what types of elements form this kind of bond?
24. What is covalent bonding and how does it satisfy the octet principle?
25. Using the periodic table draw the Lewis pictures (electron dot diagrams) of the following atoms: H, C, N, O, Cl, Ar, S, Si
26. Compare the Lewis pictures of representative elements in the same column of the periodic table.
27. How is a covalent bond represented in a Lewis picture of a molecule?
28. Use Lewis pictures to describe the chemical reaction between hydrogen and oxygen to form water.
29. Describe the properties of a polar covalent bond.
30. Describe the electron structure of metallic bonds.
31. Use the theory of covalent bonds to explain the following:

1. inert gases
2. nonmetallic diatomic gases
3. crystals
4. transition metals vs soft metals

32. Use the theory of chemical bonding to explain the following:

1. solubility of polar substances in water
2. solubility of ionic substances in water
3. electrical conductivity of electrolyte solutions

33. What are weak electrolytes?
34. What is the characteristic ion in acid solutions? In basic solutions?
35. Distinguish between strong acids and weak acids.
36. Describe what happens when ammonia is dissolved in water and when carbon dioxide is dissolved in water>
37. Discuss the statement, Acid plus base yields salt plus water.
38. What is pH and how is it measured?
39. What factors determine the pH of a particular solution?
40. Name, describe, and explain the unusual properties of water.

Objectives

1.1. Summarize the major advances of the nineteenth century in chemistry and physics
1.2. Describe the concept of periodicity of properties
1.3. Describe the structure of the periodic table of the elements
1.4. Summarize the relationship between the periodic table, electron structure, and chemical properties, of the elements
1.5. Describe the physical nature of chemical bonds using the electronic shell structure of the atom as a model
1.6. Describe the various types of chemical and physical bonds in matter
1.7. Describe the unusual properties of water and explain them with the hydrogen bonding model
1.8. Define electrolyte and distinguish between strong and weak electrolytes at the molecular level
1.9. Describe the properties of ten important substances and ten categories of chemical compounds

1. Introduction

1.1. Two of mankind's greatest intellectual achievements are the classification of the elements and the theory of the chemical bond.
1.2. They represent our modern understanding of the chemical nature of matter and allow us to manipulate it in ways undreamed of by the alchemists.
1.3. They also represent the merger and mutual confirmation of the Newtonian and atomic paradigms. The concepts of energy and momentum can be applied to orbiting electrons, attracted to a nucleus by an inverse square force which has the same form as gravity, but which operates on the quality of charge rather that upon mass.
1.4. Admittedly new principles are involved, those of the strange world of the quantum of radiation. As it turns out, the quantum principle is the missing link between the Newtonian world of forces and energy, and the atomic world of chemical reactions.

2. Nineteenth Century Chemistry

2.1. Overview

2.1.1. 19th century saw rapid advances in all sciences

2.1.1.1. from discovery of electrolysis of water to atomic theory to the periodic table in chemistry
2.1.1.2. conservation of energy, electromagnetic theory, atomic particles, kinetic theory in physics
2.1.1.3. beginning of century: caloric theory, battery invented, atomic theory and laws of chemistry
2.1.1.4. mid century: kinetic theory, conservation of energy. electromagnetism, physical chemistry {reaction rates, energy, electrochemistry, organic chemistry} (also evolution)
2.1.1.5. end of century: periodic table, discovery of electron, radioactivity, X-rays

2.2. Rare Earths

2.2.1. "earth" meant "insoluble oxide which did not compose when heated"
2.2.2. earths were not uncommon, and many of the new elements discovered by electrolysis were extracted from their oxides
2.2.3. 1794, Johan Gadolin was shown an "earth" from a quarry in Ytterby, Sweden
2.2.4. properties unlike known "earths"
2.2.5. Became known as "rare earth" when properties were described
2.2.6. chemists suspected it contained new a element
2.2.7. 1803 three chemists discovered Cerium, named after the asteroid Ceres, discovered the previous year
2.2.8. 1839 Mosander studied Gadolin's rare earths

2.2.8.1. concluded they contained a mixture of new elements, all with similar properties
2.2.8.2. isolated lanthanum (Gk. hidden)
2.2.8.3. spectroscopy and periodic table helped to discover other rare earths

2.2.9. Properties are more alike than any other elements
2.2.10. Why do these substance all have such similar properties?

2.3. Properties of the Elements

2.3.1. Lavosier's list of elements (1789) was merely a listing

2.3.1.1. did not include description of properties
2.3.1.2. many new elements discovered by mid century, properties reported

2.3.2. physical properties

2.3.2.1. hardness, shininess, malleability, density, conductivity

2.3.3. chemical properties

2.3.3.1. method of preparation, solubility in acids, reaction with other elements

2.3.4. Atomic Weight

2.3.4.1. methods of determining the relative weight of atoms were perfected
2.3.4.2. most substances were composed of atoms of fractional weights
2.3.4.3. chlorine has atomic weight of 35.5
2.3.4.4. mass spectrograph eventually showed naturally occurring isotopes

3. The Periodic Table

3.1. Patterns of Properties

3.1.1. Triads: 1817 (Johann Dobereiner) noted that elements could be grouped in threes according to chemical properties in two ways

3.1.1.1. atomic weights were practically the same (iron, cobalt, nickel)
3.1.1.2. middle atomic weight was about average of the other two (chlorine, bromine, iodine)
3.1.1.3. first attempt to group by properties

3.1.2. Octets: 1864 (John Newlands) grouped in order of atomic weights

3.1.2.1. every eighth element had similar properties

"the eight element, starting from a given one, is a kind of repetition of the first, like the eighth note in a octave of music."

3.1.2.2. called it the law of octaves.
3.1.2.3. laughed at by other scientists because it seemed too musical

3.1.2.3.1. asked whether he had ever tried classifying the elements in the order of the initial letters of their names
3.1.2.3.2. ie. eight note octave
3.1.2.3.3. later exonerated, awarded the Davy prize

3.1.3. Dmitri Menedleev (1834 - 1907) & Lothar Meyer (1830 - 1895)

3.1.3.1. simultaneously and independently invented periodic tables similar to today
3.1.3.2. Represents a significant leap in understanding of matter

3.1.3.2.1. a basis for understanding chemical reactions and properties
3.1.3.2.2. strongly supported atomic theory
3.1.3.2.3. suggested elements composed of smaller building blocks
3.1.3.2.4. suggested smaller patterned structure

3.2. Menedleev's was based on chemical properties

3.2.1. Mendeleev

3.2.1.1. born in Siberia,14th child of a teacher
3.2.1.2. father became blind, cared for by mother who managed a glass factory
3.2.1.3. 1848 traveled across Asia to enter him in University of Moscow
3.2.1.4. denied entry because he was Siberian
3.2.1.5. left for St. Petersburg, in 1850 admitted to training school for teachers
3.2.1.6. mother died the same year
3.2.1.7. became professor at St. Petersburg
3.2.1.8. resigned in 1890 over a dispute with authorities over his "too liberal views"

3.2.2. suggested undiscovered elements

3.2.2.1. where elements didn't fit correctly or atomic weights seemed wrong
3.2.2.2. left blank spaces in the table
3.2.2.3. predicted properties of missing elements
3.2.2.4. eka boron, eka aluminum, eka silicon
3.2.2.5. three soon discovered, properties nearly as predicted

3.2.3. compared with modern table

3.2.3.1. Mendeleev knew of only sixty elements, status of some uncertain
3.2.3.2. Mendeleev's chart arranged by atomic weight vs. atomic number

3.3. Meyer's was based on physical properties, esp. density

3.3.1. did not make predictions about the gaps

3.4. The Modern Periodic Table

Here's a link to an interactive online periodic table

3.4.1. Metals and Nonmetals

3.4.1.1. stair step line divides elements into two main categories
3.4.1.2. metals have similar properties, nonmetals have similar properties

3.4.1.2.1. metals

3.4.1.2.1.1. solid at room temperature (except mercury)

3.4.1.2.1.2. shiny

3.4.1.2.1.3. hard

3.4.1.2.1.4. malleable

3.4.1.2.1.5. good conductors of heat and electricity

3.4.1.2.2. nonmetals

3.4.1.2.2.1. may be solid, liquid, or gas

3.4.1.2.2.2. hard and brittle

3.4.1.2.2.3. generally not good conductors (except graphite form of carbon)

3.4.1.2.3. elements near division line may have some properties of both

3.4.1.2.3.1. ie carbon and silicon

3.4.2. Rows and Columns
3.4.3. Groups

3.4.3.1. A & B

3.5. The Noble Gases

3.5.1. Mendeleev's chart was missing a whole column of gases

3.5.1.1. called noble gases, inert gases, rare gases, helium group
3.5.1.2. undiscovered because they do not react to form compounds
3.5.1.3. all gases, not abundant, hid from chemical observation for two centuries

3.5.2. First clue ignored by scientific community: Cavendish

3.5.2.1. removed nitrogen from air, then oxygen, still had a tiny bubble of gas

3.5.2.1.1. about 1/120 (0.8%) of volume of original air
3.5.2.1.2. most experimenters would have discounted as "experimental error"
3.5.2.1.3. experiment repeated in 1894 (by Rayleigh and Ramsay)

3.5.2.1.3.1. used spectroscope to recognize as new element

3.5.2.1.3.2. named argon, Greek for inert

3.5.3. Discovery of Helium

3.5.3.1. Pierre Janssen 1868 observed solar eclipse with spectroscope
3.5.3.2. found bright yellow line in solar spectrum, not attributed to any known element
3.5.3.3. British astronomer (Lockyer) suggested element was not found on Earth

3.5.3.3.1. named it helium from Greek Helios
3.5.3.3.2. Ramsey and Rayleigh recognized a whole family

3.5.3.3.2.1. isolated helium on Earth

3.5.3.3.2.2. also found krypton, neon, xenon

3.5.3.3.3. Radon discovered by Curies as daughter of Radium decay

4. Electron Patterns

4.1. Introduction

4.1.1. atomic theory spurred the progress of chemistry

4.1.1.1. from trial and error of alchemists to systematic discovery of formulations
4.1.1.2. many generalizations of creation of compounds: what works and what doesn't

4.1.2. many unanswered question at beginning of 20th century

4.1.2.1. why do certain elements have similar properties (periodic table?)
4.1.2.2. why do some elements occur as diatomic pairs, others as single atoms, others as solids of uncertain formula?
4.1.2.3. why does hydrogen atoms combine with oxygen atoms in 2:1 ratio
4.1.2.4. why are some compounds soft with low melting temperatures, others hard with high melting?
4.1.2.5. Why do some substance dissolve in water and others do not?
4.1.2.6. Why are some substances electrolytes while others are not?

4.1.3. answers came from physics of the atom which discovered pieces of atoms

4.1.3.1. cathode rays
4.1.3.2. radioactivity
4.1.3.3. Millikan oil drop experiment
4.1.3.4. Rutherford's gold foil and alpha bullets
4.1.3.5. photoelectric effect
4.1.3.6. matter waves and quantum mechanics
4.1.3.7. spectroscopy
4.1.3.8. mass spectrograph showed existence of isotopes
4.1.3.9. discovery of the neutron as nucleon

4.1.4. chemists were quick to adopt the nuclear model to explain their discoveries and make new predictions

4.2. Nuclear Model Reviewed

4.2.1. atoms consist of dense nucleus of protons and neutrons surrounded by cloud of negative charge

4.2.1.1. proton and neutron are similar in mass but neutron has no charge
4.2.1.2. proton and electron are opposite in charge, but proton is 1800 times more massive

4.2.2. electrons occupy certain energy levels around nucleus, define outer boundaries of atom

4.2.2.1. prevent nuclei of atoms from touching

4.3. Patterns in Electron Structure

4.3.1. patterns

4.3.1.1. atomic numbers of noble gases differ by certain amounts

4.3.1.1.1. He = 2
4.3.1.1.2. Ne = 10
4.3.1.1.3. Ar = 18
4.3.1.1.4. Kr = 36
4.3.1.1.5. Xe = 54
4.3.1.1.6. Rn = 86

4.3.1.2. magic numbers (2, 8, 18, 32)
4.3.1.3. answer lies in quantum mechanics
4.3.1.4. each shell and subshell can contain a fixed number of electrons
4.3.1.5. filling of shells and subshells takes place in and ordered and structured way
4.3.1.6. shown clearly in periodic table

4.3.1.6.1.

4.3.1.7. are these just Pythagorean coincidences, or do they mean something?

4.3.2. planetary model of electrons is incomplete, electrons in orbits differ from satellites:

4.3.2.1. can be only in certain energy levels called shells, not in between
4.3.2.2. each electron shell has a limited number of quantum substates
4.3.2.3. no two electrons can occupy the same quantum state

4.3.2.3.1. there are definite maximums of electrons in a given energy level
4.3.2.3.2. defined by mathematical relationships

4.3.2.4. electron orbits cannot be known precisely

4.3.2.4.1. electron "orbits" are really fuzzy regions of probable locations called orbitals

4.3.3. electrons are in different energy levels, called shells

4.3.3.1. energy is sum of kinetic and potential energies, momentum is both angular and spin
4.3.3.2. energy and momentum are quantitized
4.3.3.3. principal quantum number is the "shell" number, like the row in a stadium
4.3.3.4. number of electrons in each energy level is limited by properties of electron
4.3.3.5. shells consist of "subshells", the number of which get greater as the principle quantum number increases
4.3.3.6. maximum number of electrons per shell = 2n²

4.3.4. most stable arrangement is 8 electrons in outermost subshell

4.3.4.1. four electron pairs

4.4. Representative Elements and Transition Elements

4.4.1. 2 left columns and 6 right columns are representative elements

4.4.1.1. have A group numbers on periodic table
4.4.1.2. number of electrons in outermost (valence) shell is same as A group number
4.4.1.3. governs chemical properties

4.4.2. middle columns are transition elements

4.4.2.1. have B group numbers on periodic table
4.4.2.2. number of outer electrons varies from one to three
4.4.2.3. complications due to sublevels

4.4.2.3.1. we will largely ignore transition elements

4.4.3. once number of outer electrons is known, chemical properties are easily understandable

5. Chemical Bonding

5.1. caused by electrical forces within and between atoms and by nature of electrons

5.1.1. magic numbers: eight electrons in four pairs per valence shell, except for first shell
5.1.2. inert gases all have magic number of electrons

5.1.2.1. called octet rule (exception is hydrogen and helium)

5.1.2.1.1. first "shell" can have a maximum of two electrons
5.1.2.1.2. neutral inert gases have stable electron configuration

5.1.2.1.2.1. not totally inert, can be made to react under special conditions

5.1.3. other atoms can get a magic number of electrons by gaining or losing to become ions

5.2. three or less electrons can be gained or lost from valence shell

5.2.1. more than three is possible but not energetically favorable

5.3. pairs of electrons are favored over single electrons

5.3.1. pairs form only when there are more than four
5.3.2. single electrons may be transferred to other atoms or shared between atoms to form chemical bonds

6. Types of Chemical Bonds

6.1. ionic bonds

6.1.1. electron transfer from one atom to another
6.1.2. between metals and nonmetals
6.1.3. all but group IV

6.2. covalent bonds

6.2.1. sharing of electrons to complete the octet
6.2.2. unpaired electrons are shared between atoms
6.2.3. Lewis pictures

Lewis pictures are a way of representing chemical bonds by using the outermost or valence electrons:

6.2.4. between atoms of similar properties
6.2.5. sharing is unequal to cause polar molecules

6.3. metallic bonds

6.3.1. delocalized electrons form "sea" in metallic crystal lattice
6.3.2. strong metals such as iron also have covalent bonds

6.4. physical bonds

6.4.1. weak bonds due to unbalanced electrical forces
6.4.2. van der Waals bonds
6.4.3. hydrogen bonds

7. Electrolytes, Acids, Bases, & Salts

7.1. Water and Electrical Conduction

7.1.1. pure water does not conduct electricity
7.1.2. certain substance increase conductivity when dissolved

7.1.2.1. sugar and alcohol don't help, but salts do

7.1.3. electrolyte = substance which increases electrical conductivity of water

7.1.3.1. dry salts don't conduct electricity either

7.2. Acids

7.2.1. turn litmus red
7.2.2. are corrosive to most metals
7.2.3. result from aqueous solutions of oxides of nonmetals
7.2.4. have sour taste
7.2.5. react with bases to form salts
7.2.6. produce hydrogen ions in aqueous solution

7.2.6.1. hydrogen ion may attach to water molecule to form hydronium ion

7.3. Strong vs. Weak Acids

7.3.1. strong acids are substances which completely yield their hydrogen ions to the solution

7.3.1.1. solution reactions proceed to completion, so nearly all hydrogen is released as ions
7.3.1.2. examples:

7.3.2. weak acids are weak electrolytes which yield some of their hydrogen as ions

7.3.2.1. solution does not proceed to completions, some hydrogen remains bonded to molecule

7.3.2.1.1. example:

7.3.2.1.2. double arrows means that reaction reaches equilibrium with some fraction of hydrogen still bonded

7.4. Bases (or alkalis)

7.4.1. turn litmus blue
7.4.2. have bitter taste
7.4.3. have slippery feel
7.4.4. result from aqueous solution of oxides of metals

7.4.4.1. metals nearer left side of periodic table form strongest bases

7.4.5. react with acids to form salts
7.4.6. produce hydroxide ions in aqueous solution

7.5. Strong vs. weak Bases

7.5.1. strong bases are ionic compounds which completely yield their hydroxide ions to the solution

7.5.1.1. solution reactions proceed to completion, so nearly all hydrogen is released as ions
7.5.1.2. examples:

7.5.2. weak bases are covalent compounds (weak electrolytes) which accept hydrogen ions from water molecules.

7.5.2.1. example:

7.5.2.2. another incomplete reaction, some ammonia remains simply dissolved in water

7.5.2.2.1. example:

7.5.2.2.2. double arrows means that reaction reaches equilibrium with some fraction of hydrogen still bonded

7.6. Neutralization Reactions

7.6.1. acid + base = water + salt

7.6.1.1. essential reaction between acid and base is

7.6.1.2. other ions remain in solution,

7.6.1.2.1. nonmetal from acid + metal from base = salt
7.6.1.2.2. may form solid ionic substance under certain conditions

7.7. The pH scale

7.7.1. pH means "powers of hydrogen"
7.7.2. pH is negative logarithm of hydrogen ion concentration in aqueous solutions

7.7.2.1. pure water contains a certain amount of hydrogen ions due to dissociation of water molecules

7.7.2.1.1. contains equal numbers of hydrogen and hydroxide ions so is neither acid nor basic
7.7.2.1.2. concentration is 1E-7 moles per liter, so pH of pure water is 7
7.7.2.1.3. pH can be changed whenever hydrogen ion concentration changes

7.7.2.1.3.1. adding hydrogen ions (ie. dissolving acids in water)

7.7.2.1.3.2. adding hydroxide ions (dissolving strong bases in water)

7.7.2.1.3.3. removing hydrogen ions (dissolving weak bases in water)

7.7.2.2. pH of aqueous solutions depends on strength and concentration of substance

7.7.2.2.1. concentration means how much of the substance is dissolved in water
7.7.2.2.2. strength means what fraction of its hydrogen or hydroxide is surrendered to solution

7.7.2.3. pH of some common substances

7.7.3. pH regulation (buffering) is important in the natural world

7.7.3.1. pH of ocean is maintained at nearly constant level
7.7.3.2. pH of blood is extremely important to bodily functions

7.7.3.2.1. must be maintained within narrow limits (7.35-7.45)

7.7.3.3. pH of stomach must be within narrow limits for proper digestion
7.7.3.4. most organisms have low tolerance for pH variation
7.7.3.5. most foods are slightly acidic

7.7.3.5.1. citrus fruits (citric acid), peppers (oxalic acid), tomatoes, Coca-Cola (phosphoric acid), etc.

7.7.3.6. most cleaning chemicals are basic

7.7.3.6.1. laundry detergents, soaps can damage materials which cannot tolerate high pH

7.7.3.6.1.1. wool, car paints, hair, skin are all pH sensitive

7.7.4. pH measurement

7.7.4.1. electronic sensors
7.7.4.2. pH sensitive dyes

8. Important Chemical Compounds

8.1. Categories of Compounds

8.1.1. Inorganic

8.1.1.1. Acids

8.1.1.1.1. Highly corrosive, especially with metals. Formed from reaction of nonmetal oxides and water. React with bases to form water and salt. Contain easily ionized hydrogen atom. Sour taste at low concentrations.

8.1.1.1.2. hydrochloric
8.1.1.1.3. nitric
8.1.1.1.4. sulfuric
8.1.1.1.5. acetic
8.1.1.1.6. citric

8.1.1.2. Bases

8.1.1.2.1. Slippery feel, bitter taste in low concentrations. Formed from reaction of metallic oxides and water. React with acids to form salt and water. Contain easily ionized hydroxide radical (OH).

8.1.1.2.2. sodium hydroxide
8.1.1.2.3. potassium hydroxide
8.1.1.2.4. ammonium hydroxide

8.1.1.3. Salts

8.1.1.3.1. Solid, soluble crystals with high melting temperatures. Ionic bonds between metal and nonmetal ions. Formed from reaction of acid and base.

8.1.1.3.2. sodium chloride
8.1.1.3.3. potassium nitrite
8.1.1.3.4. calcium chloride

8.1.1.4. Oxides

8.1.1.4.1. Highly stable chemical combinations of oxygen and another element.

8.1.1.4.2. water
8.1.1.4.3. carbon dioxide
8.1.1.4.4. iron oxide (rust)
8.1.1.4.5. silicon dioxide (quartz)

8.1.1.5. Hydrides

8.1.1.5.1. Chemical combinations of hydrogen and another element.

8.1.1.5.2. water
8.1.1.5.3. methane
8.1.1.5.4. ammonia

8.1.2. Organic

8.1.2.1. Hydrocarbons

8.1.2.1.1. Highly combustible chemical combinations of carbon and hydrogen.

8.1.2.1.2. methane
8.1.2.1.3. ethane
8.1.2.1.4. acetylene
8.1.2.1.5. propane
8.1.2.1.6. butane
8.1.2.1.7. octane
8.1.2.1.8. hexane

8.1.2.2. Alcohols

8.1.2.2.1. Hydrocarbons with hydroxyl (OH) radical replacing one or more hydrogen atoms.

8.1.2.2.2. ethyl alcohol (ethanol)
8.1.2.2.3. methyl alcohol (methanol)
8.1.2.2.4. isopropyl alcohol (isopropanol)
8.1.2.2.5. glycerine (gylcerol)

8.1.2.3. Carbohydrates

8.1.2.3.1. Carbon, hydrogen, oxygen in various forms and structures. Generally have (CH20)n. Includes sugars, starches,
8.1.2.3.2. .
8.1.2.3.3. sucrose
8.1.2.3.4. glucose
8.1.2.3.5. fructose
8.1.2.3.6. starch
8.1.2.3.7. cellulose

8.1.2.4. Lipids

8.1.2.4.1. Fats and oils. Triglyceride esters of fatty acids.

8.1.2.5. Proteins

8.1.2.5.1. Carbon, hydrogen, oxygen, nitrogen, (sulfur). Polymers of amino acids held together by peptide bonds. Tens of thousands of know varieties. Used for structure (skin, bones, hair, nails), immune response, enzymes, hormones, etc.

8.2. Ten Important Compounds

There are too many compounds to know about all of them, but certain compounds can be used as models, others can be related to them.

Properties of compounds are explained by theory of chemical bonds. Knowing what needs to be explained will make bonding theory more necessary and useful.

8.2.1. Water

8.2.1.1. most important compound on Earth separate section later

8.2.2. Sodium Chloride

8.2.2.1. salt is general name for a class of compounds but also specifically means "table salt"
8.2.2.2. analysis shows it to be NaCl (40% Na, 60%Cl by weight)

8.2.2.2.1. will explore 1:1 ration of Na:Cl later

8.2.2.3. properties

8.2.2.3.1. brittle, white, crystalline solid which is moderately soluble in water

8.2.2.3.1.1. 365 gm/liter at room temperature (20 C)

8.2.2.3.1.2. only slightly more so in warm than in cold

8.2.2.3.2. melting point is 800 C, much higher than sugar which look similar

8.2.2.3.3. present in large amounts in seawater, often found as evaporite
8.2.2.3.4. made in laboratory by mixing sodium hydroxide and hydrochloric acid

8.2.3. Hydrochloric Acid

8.2.3.1. industrial name is muriatic acid

8.2.3.1.1. one of Lavoisier's elements

8.2.3.2. solution of gaseous hydrogen chloride HCl in water
8.2.3.3. gas and aqueous solution is highly corrosive

8.2.3.3.1. dissolves many metals, but not all

8.2.3.4. acidic properties

8.2.3.4.1. sour taste, corrosiveness, turns litmus red
8.2.3.4.2. shared with other acid substance

8.2.3.5. can be made directly from combustion of hydrogen in chlorine
8.2.3.6. industrial HCl prepared from NaCl and sulfuric acid

8.2.4. Ammonia

8.2.4.1. gas with formula NH₃
8.2.4.2. characteristic odor (smells like ammonia)

8.2.4.2.1. used in smelling salts

8.2.4.3. very soluble in water

8.2.4.3.1. makes ammonium hydroxide solution

8.2.4.4. has alkaline or basic properties

8.2.4.4.1. feels slippery
8.2.4.4.2. bitter taste
8.2.4.4.3. do not corrode most metals

8.2.4.5. reacts with hydrogen chloride to form ammonium chloride

8.2.4.5.1. in aqueous solution also forms extra water
8.2.4.5.2. acid + base --> salt + water

8.2.5. Methane (CH₄)

8.2.5.1. "natural gas" used as fuel

8.2.5.1.1. combustion is reaction with oxygen
8.2.5.1.2. forms water and carbon dioxide

8.2.5.2. member of large family called hydrocarbons
8.2.5.3. colorless, odorless
8.2.5.4. insoluble in water
8.2.5.5. neither acid nor base, doesn't react with either

8.2.6. Acetylene

8.2.6.1. also hydrocarbon gas, similar but different from methane

8.2.6.1.1. unsaturated hydrocarbon
8.2.6.1.2. neither acid nor base, but soluble in water

8.2.6.2. C₂H₂, burns with smoky flame unless enriched with oxygen

8.2.6.2.1. produces hot flame, used in welding torches

8.2.6.3. made from water + calcium carbide

8.2.6.3.1. used in miner's lamps before flashlights

8.2.7. Sugar

8.2.7.1. table sugar is C₁₂H₂₂O₁₁

8.2.7.1.1. carbohydrate
8.2.7.1.2. many different kinds of sugars

8.2.7.2. white crystalline solid, similar in appearance to table salt

8.2.7.2.1. differs in taste, melting point, solubility
8.2.7.2.2. much more soluble in warm water than cold

8.2.7.3. neither acid nor base

8.2.7.3.1. doesn't react easily with either

8.2.7.3.1.1. sulfuric acid dehydrates (removes water, leaves carbon)

8.2.7.3.1.2. hydrolysis breaks into two other sugars, glucose and fructose

8.2.7.3.1.2.1. both are C6H12O6 but different arrangement of atoms

8.2.8. Ethanol (grain alcohol)

8.2.8.1. member of class of compounds called alcohols
8.2.8.2. present in beer, wine, liquors as C₂H₅OH
8.2.8.3. infinitely soluble in water

8.2.8.3.1. due to oxygen atom since hydrocarbons are not generally soluble
8.2.8.3.2. will react with some acids and bases in aqueous solution

8.2.9. Potassium Nitrate (KNO₃)

8.2.9.1. saltpeter occurs naturally in many locations
8.2.9.2. important in manufacture of gunpowder
8.2.9.3. related compound is sodium nitrate (soda niter)
8.2.9.4. in general category of salt compounds

8.2.9.4.1. made from reaction of potassium hydroxide with nitric acid

8.2.10. Sulfuric Acid (H₂SO₄)

8.2.10.1. colorless syrupy extremely corrosive liquid
8.2.10.2. infinitely soluble in water

8.2.10.2.1. solution is highly exothermic

8.2.10.2.1.1. chemical reaction rather than just simple solution

8.2.10.2.2. water may boil when added in drops

8.2.10.3. most important industrial chemical

8.2.10.3.1. 65 billion pounds produced yearly
8.2.10.3.2. common in laboratories
8.2.10.3.3. commonly used to produce many other chemicals

8.2.10.4. reacts with almost anything including hydrocarbons
8.2.10.5. typical acid otherwise

9. Unusual Properties of Water

9.1. common but unusual in many ways

9.1.1. properties are important to life and planet

9.2. water molecules stick to each other more than most other light molecules

9.2.1. most (methane, HCl, ammonia, H2S) are gases at room temperature, water is liquid
9.2.2. due to shape of molecule and hydrogen bonding
9.2.3 asymmetry of molecule makes the hydrogen end slightly negative and the oxygen end slightly positivle
9.2.4 makes water molecules unusually attracted to one another

9.3. unusual density/temperature relationship

9.3.1. water is most dense at 4 deg C
9.3.2. most substances get less dense with temperature
9.3.3. most solids are more dense than corresponding liquid

9.3.3.1. ice is less dense than water
9.3.3.2. prevents lakes from freezing from bottom up
9.3.3.3. expands and exerts pressure when it freezes

9.4. water has highest specific heat

9.4.1. absorbs heat with small change in temperature
9.4.2. useful as coolant
9.4.3. moderates climate

9.5. water has highest latent heat

9.5.1. much heat needed to melt ice or to boil water

9.6. water is best solvent (dissolves many substances)
9.7. water has high surface tension

9.7.1. like a "skin"
9.7.2. forms drops
9.7.3. hard to penetrate surfaces
9.7.4. forms meniscus in containers

10. Summary & Conclusions

Although there is a large amount of material included in this lesson, it is based on a few simple principles.. First we summarized some of the discoveries of chemistry in the 19th century.

We learned about the periodic properties of the elements and the first attempts to classify the elements based on similar properties that culminated in Mendeleev's table.

The patterns in the number of electrons can be related to the position of the element in the periodic table, especially for the representative elements. This pattern also can be used to understand chemical bonding, both ionic and covalent..

The remainder of the program focused on specific examples of varilous substances: substances that contain both ionic and covalent bonds such as electrolytes, acids, and bases; examples of types ofimportant chemical compounds; polar hydrogen bonding in water that gives it unusual properties.