''Kings Chemistry Survival Guide''.pdf

KINGS CHEMISTRY

SURVIVAL GUIDE

A guide for the hobbyist, enthusiast, or amateur for the preparation of common, and un-common

laboratory chemicals

EDITION 1

By Jared B. Ledgard

KINGS CHEMISTRY SURVIVAL GUIDE: EDITION 1 ®

Writers of scientific and technology literature

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illegal. Please adhere to these copyrights.

The author, writer, and publisher take no responsibility for the actions of anyone as a result of this manual. People

who use this manual to make or prepare lab chemicals, or related compositions in anyway take full responsibility

for their actions. Any injuries, deaths, or property damage caused or produced by the actions of person or persons

using this manual are not the result or responsibility of the author, writer, or publisher. Furthermore, any laws or

legal issues broken, violated, or disturbed in anyway by or as a result of person or persons using this manual are not

the responsibility of the author, writer, or publisher. Any attempt to sue or bring about any form of legal action

against the author, writer, or publisher as a result of injury, death, violation of law, or property damage caused by

the negative intentions of a person or persons who used information in this manual is a direct violation of freedom

of speech laws, and information right-to-know laws.

Information contained in this manual was compiled, formatted, and translated from a variety of chemical abstracts,

documents, and journals, all of which are therefore public record and hereby bound to freedom of speech and

information protection laws as discussed in the US constitution under the information-right-to-know acts. The

information contained in this manual was edited, and rewritten to fit a form readable by the common man as well as

scientist. The information is not the sole responsibility of the author, writer, or publisher. Any injuries, deaths, law

violations, or property damage associated with any of the procedures detailed in this manual are not the result, nor

responsibility of the author, writer, or publisher. Every procedure discussed in this manual has been successfully

carried out with safe, reliable, and effective results. Any attempt to sue or bring about any form of legal action

against the author, writer, or publisher as a result of a person or persons negligence, stupidity, or gross

incompetence is a direct violation of freedom of speech laws, and information right-to-know laws.

This manual is intended for educational purposes only, and the author, writer, and publisher are not aware of any

danger, or illegal acts this manual may or may not pose to people or property if used by person or persons with

negative intentions. The author, writer, and publisher have no intent, nor desire to aid or provide potentially

dangerous information to persons with desires to injure, kill, violate laws, or cause property damage. The

information contained in this manual is for reference purposes only, and the author, writer, and publisher made this

manual possible to inform, enlighten, and educate persons interested or curious in the art of laboratory chemistry.

This manual was created by the author, writer, and publisher to deliver knowledge and truth. Any attempts to sue or

bring about law suits against the author, writer, or publisher for any reason associated with this manual is a direct

violation of knowledge and truth, and is therefore, a violation of the US constitution.

SECTION 1: Introduction to Chemistry

A quick lesson in chemistry

Part 1: Introduction to chemistry

This book has been written to teach the art of general chemistry sciences to the reader. To do this, you should take a quick, yet vital

lesson in chemistry. First of all, the world of chemistry is a fascinating world filled with a huge variety of chemicals, chemical

reactions, formulas, laboratory apparatus, and an arsenal of equipment. All these elements are combined and used thoroughly to bring

about chemical change of matter from one form to the next. In this book, the form of change that we will deal with mostly, is the

formation of compounds that are regarded as general laboratory reagents.

The world of general chemistry is absolutely huge, and in essence, deals with virtually ten’s of thousands of chemical compounds.

Regardless how many possible chemicals there might be, most see chemicals as something evil or something that is a troublesome or

bothersome contaminant on our foods, households, and everyday possessions; however, in factuality, chemistry and the chemicals

involved are responsible for our modern civilization, and without them, we would all be in big trouble. The art of chemistry is as old

as life itself, and as old as our universe.

For most of you, the procedures in this book will not make sense at first, or will appear to be complicated; as a result, many of the

procedures in this book may seem foreign, or unfamiliar—if this is the case, then at this exact moment, you are in the right place. Bye

the time you have read this book, these “foreign” procedures will no longer be foreign to you, but in the meantime, lets get started on

the world of chemistry.

The world of chemistry involves every single aspect, corner, and micro drop of everything that is matter. Our solar system and the

entire universe all function on a chemical level—In essence, chemistry is everything. The universe and everything in it is composed of

atoms and molecules, and within this massive space, there exists tens of millions of chemical compounds—either known or unknown.

The compounds that are known make up only 5% of the naturally occurring compounds, leaving a massive 95% of them being

synthetic (prepared in the lab)—95% of all chemicals are synthetic. Note: synthetic does not denote anything that is less superior to

natural. Synthetic means creating natural in an un-natural way.

Chemistry has been divided into three fields over the last 100 years to better organize and format the system. The three major branches

of chemistry include: Inorganic chemistry, Organic chemistry, and Biochemistry. In short, inorganic chemistry deals with ionic

compounds, which make up the chemical compounds that do not contain active carbon. Organic chemistry is the largest branch of

chemistry and it deals with covalent compounds, which make-up our everyday items like plastics, drugs, dyes, pesticides, insecticides,

resins, fibers, and explosives. Organic means “carbon bearing” which means any compound that bears carbon is classified as organic.

Gasoline, turpentine, and candle wax are specific examples of organic compounds. Last but not least, biochemistry studies the field of

enzymes, organisms, plants, and animals and their active chemical processes. Genetics research studies the DNA and RNA of living

things and is a sublevel of biochemistry. DNA and RNA is composed of organic compounds all linked and actively working together.

Biochemistry deals heavily with peptides, amino acids, carbohydrates, ect., ect., all of which play a major role in natural process such

as cells, metabolism, and the like.

1. Chemical bonding: Oxidation states

First things first, you need to understand the nature of elements, and their oxidation states (number of bonds). Every single element is

capable of forming chemical bonds with other elements (with the exception of a few “noble gases”). The oxidation states are what

determines how many bonds a particular element can form, and to what other elements. When elements combine, they form chemical

compounds. All of the atoms within a chemical compound show specific oxidation states. Oxidation states are not really states, but

definitions of bonding, which are dictated by each individual element. Each element can form any where from either 0 to 7 bonds.

These numbers represent the number of bonds the element can form (look at a modern periodic table, such that included in the “Merck

Index”—the oxidations states are written in the upper left corner of each element). These numbers clearly indicate the number of

bonds each element is capable of forming.

As most people are aware, periodic tables include rows and columns filled with elements. The elements within any given column have

similar properties and characteristics along with similar oxidation states. For example, the elements of column 5A on the periodic

table include nitrogen, phosphorus, arsenic, antimony, and bismuth. All these elements have similar oxidation states and properties.

Phosphorus for example, can form compounds with three bonds or five bonds (indicated by the numbers +3, –3, and +5). Phosphorus,

like arsenic and antimony have oxidation states of +3, –3, and +5. Phosphorus can form either +3 or +5 oxidation states when it bonds

to elements with higher electro negativities (also listed on some periodic tables), and –3 oxidation states with elements that have lower

electro negativities. Each element has different electronegative energies. Metals for example, have electro negativities ranging from

0.60 to 1.9. Non-metals have electro negativities ranging from 1.9 to 4.0. In essence, elements that are metals combine with the

elements called non-metals forming positive oxidation states, with the so-called non-metals forming negative oxidation states.

In a specific example, when phosphorus reacts with non-metals it forms +3 and +5 oxidation states because its electronegative energy

is less then the other non-metals, but when it bonds to metals, its oxidation state is –3 because its own electro negative energy is

greater then most metals.

Either way, when two elements combine for example, the element with the greater electronegative energy forms negative oxidation

states, and the element with the lower electronegative energy forms positive oxidation states. In another example, chlorine and

bromine both have greater electronegative energies, so when they combine with phosphorus, the phosphorus forms +3 and +5

oxidation states (see the illustration below). When elements combine they form compounds, which are called molecules.

Elements such as lithium, sodium, and potassium form only one bond, because they have only a +1 oxidation state, and because their

electronegative energies are quite low (ranging from 1.0 to 0.6). A more complex array of oxidation states is demonstrated in the

element nitrogen (a key element found in all amphetamines). It’s capable of forming +1, +2, +3, +4, +5, –1, –2, and –3 oxidation

states (see the illustration below). Another crucial element, carbon, is capable of forming +2, +4, and –4 oxidation states, and the all

important oxygen, forms only a –2 oxidation state. Hydrogen can form +1 and –1 oxidation states. Remember the elements helium,

neon, and argon (called the noble gases) form no oxidation states. Note: The oxidation states of each element (and column of elements

on the periodic table) have been determined by trial and error over some 200 years of chemical research and study.

2. Ionic compounds and ionic bonds

Ionic compounds are composed of elements bonded together that have marked differences in electro negativities. Ionic compounds

make up the bulk of “inorganic compounds”, and are composed primarily of metals bonded to non-metals. In ionic compounds, the

oxidation states of each element follows the same rules governed by the number of bonds each element can form. In the case of ionic

compounds, the positive and negative numbers represented by the number of bonds each element can form, is more detailed and also

represents a charge attributed to each element. For example, when phosphorus bonds to chlorine, it forms +3 or +5 oxidation states,

and the chlorine forms a single –1 oxidation state; however in this example, because the electronegative difference between the

phosphorus and the chlorine is not very significant, the resulting phosphorus trichloride or pentachloride is not considered fully to be

ionic. However, in the case of sodium chloride, a +1 sodium ion is bonded to a –1 chlorine atom, with each positive and negative

mark defined as a charge. Compounds that have their oxidation states defined as actual charges are considered to be ionic. As a

reminder, remember that oxidation states (the numbers) define the number of bonds an element can form, nerve mind the positive or

negative marks each number has. In ionic compounds the molecules are made up of positive and negatively charged atoms

corresponding to their oxidation state number (the number of bonds each element can form, i.e., the oxidation state number defines the

number of bonds each element can form, but not their electrical charge in all molecules—just in ionic molecules.

The electrical charge of each element within an ionic molecule is different then the element’s electronegative energy. Note:

Electronegative energy determines whether the element forms positive or negative oxidation states. Electrical charge is determined

after the atoms combine, and is represented by the positive or negative oxidation state independently from the actual number of bonds

each element can form.

As previously stated, chlorine is more electronegative then sodium, so when they combine the chlorine forms a –1 oxidation state

(notice on a periodic table that chlorine has an oxidation state of +1, –1, +5, and +7; and sodium has an oxidation state of +1). Some

periodic tables give the electronegative energy of each element, and using such a periodic table, you will notice that the electro

negativity of chlorine is remarkably higher then that of chlorine. Because the difference between electronegative energies is so great,

the chlorine becomes negatively charged, and the sodium becomes positively charged. These charged atoms attract each other, and

hence form a bond based on their electrical attractions (like two magnets)—this is the basis of “ionic” bonds.

Oxidation states also determine the number of electrons that can be captured. As previously discussed, ionic compounds like sodium

chloride form their bonds based on electrical attractions. These attractions are determined by the number of electrons a particular atom

captures. When chlorine combines (reacts) with sodium it forms a –1 oxidation state. Again, because the difference in electronegative

energies is so great, the chlorine grabs or captures one of the sodium’s electrons. This capturing causes the chlorine to become

negatively charged. As a result, the sodium atom becomes positively charged. Atoms become negatively charged when they capture

electrons, and become positively charged when they loose electrons. This capturing and loosing of electrons is the scientific

foundation to ionic bonding and ionic compounds.

Currently there are about 200,000 ionic compounds known to man (most of them being synthetic). The most common ionic compound

is table salt or sodium chloride. Some common examples of ionic compounds include potassium permanganate, sodium azide, sodium

nitrate, potassium chloride, sodium fluoride, potassium chlorate, and zinc sulfate. Ionic compounds make up the majority of the earth,

solar system, and the universe.

3. Covalent compounds and covalent bonds

Covalent compounds make up the bulk of chemical compounds known to man, but they only a make-up a small percentage of the

chemical compounds found on earth and earthly like planets, and virtually most solar systems. As previously stated, there are about

200,000 ionic compounds known to man, with a potential of another 100,000 left undiscovered throughout the universe; however,

covalent compounds number in the millions. For example, currently there are 16,000,000 covalent compounds known to man (as of

2003). The possible number of covalent compounds is practically endless, as the combination of these compounds is virtually infinite.

Covalent compounds contain covalently bonded carbon atoms. The term “organic” means ‘carbon bearing covalent substance’.

Covalent compounds all contain specific carbon atoms, which make-up the foundation or infrastructure of all organic compounds. A

covalent compound such as hexane for example, is composed of covalently bonded carbon atoms all bonded together to form a

chain—this chain represents the backbone or infrastructure of the molecule. The carbon atoms that make up these backbones or

infrastructures, are themselves bonded directly to other atoms such as hydrogen, oxygen, nitrogen, sulfur, phosphorus, arsenic, ect.,

ect. Such examples of covalent compounds (organic compounds) include: ethyl alcohol, isopropyl nitrate, aspirin, acetaminophen,

cocaine, and octane.

Covalent bonds are much different then ionic bonds, as they share electrons rather then “capture” them. Remember that ionic bonds

are formed when two or more elements with distinctive differences in electro negativities react with one another—whereby the greater

electronegative element captures an electron (or more) from the less electronegative element(s). Covalent bonds, however, are formed

when two or more elements combine and the electrons are shared (paired) rather then captured. In order for a covalent bond to form,

the electronegative differences between the elements cannot be very significant, meaning their differences are much less then those

encountered with ionic bonds.

Covalent bonds cover a whole echelon of reactions, many of which can be very complex and/or require special conditions depending

on the chemicals and reaction conditions, and usually require multiple reactions and steps to achieve desired products. In other words,

ionic compounds tend to be rather simplified compounds with easy formulas, whereas organic compounds can be huge molecules,

which require many steps for their preparation. These multiple steps are the basis for organic chemistry, as it deals with a whole

multitude of reactions and functional groups—most of these reactions and functional groups will not be discussed in this book (as it

would take about 100,000+ pages), but what functional group reactions that will be discussed are the amino functional groups

commonly found in amphetamines and derivatives.

In general, covalent bonds are less stable then ionic bonds. Most ionic compounds are stable solids with relatively high melting points

(ranging from 200 to 2400 Celsius). Many ionic compounds can be heated to very high temperatures without any significant

decomposition, such examples include: aluminum oxide, iron oxide, sodium chloride, and magnesium chloride. Most organic

compounds decompose when heated to temperatures above 300 to 500 Celsius. The high melting points of ionic compounds are due

primarily to crystal structure, and the result of strong electrical attractions between the elements and the molecules—these attractions

can lead to super strong crystal lattices, as seen in some compounds like aluminum oxide (emeralds), and other ionic oxides (gems and

sapphires). There is one mere example of an organic compound that should be demonstrated here; diamonds are composed of

covalently bonded carbons atoms, with the molecules forming super strong crystal lattices.

Other then this isolated example, most covalent compounds are solids or liquids with relatively low melting points and boiling points.

This is the result of weaker electrical attractions between the molecules. In covalent compounds the weaker attractions exist primarily

because the covalent molecules lack ionic charges, and are thereby not attracted or repelled to each other very much. Because of the

lack of electrical attractions between covalent molecules, the boiling points of covalent molecules are the result of “intermolecular”

forces (the melting points will be discussed shortly). Intermolecular forces are forces that exist between elements within one molecule

upon different elements within another molecule. Such an example would be water, common hydrogen oxide. Water which is

composed of two hydrogens bonded to a single oxygen has a significant boiling point of 100 Celsius at sea level, although it is a

relatively small and light molecule. The reason water has such a high boiling point for its small size and weight, is due to

intermolecular force attractions between the central oxygen atom of one molecule upon the two hydrogens of another water molecule

(adjacent water molecule). The non-bonding type attractions (intermolecular forces) that water molecules have to each other is what

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