CALCULATE

Stoichiometry

A. CHEMICAL BASIC LAW

        Chemistry is a part of natural science that studies matter that includes the composition, nature, and change of matter and energy that accompanies material changes. Careful research on reagents and reaction products has spawned basic chemical laws that show a quantitative or so-called stoichiometric relationship. Stoichiometry comes from the Greek word stoicheon which means element and metrain which means to measure. In other words, stoichiometry is a chemical calculation involving the quantitative relation of substances involved in the reaction. These basic chemical laws are the law of mass conservation, the law of fixed comparison, the law of volume comparison, and the law of multiple comparisons. The basic laws of chemistry are our foundation in studying and developing the next chemistry.

1. LAW MASS LAW (LAVOISIER LAW)

        At the beginning of the eighteenth century, chemists in their attempts to learn the heat and burning found a very strange thing. For example, if wood is burned, it will produce ash residue (solid) which is much lighter than the original wood. However, if the metal is burned in the free air, it will produce heavier oxide than the original metal. To answer the strangeness, chemists developed a method of experimentation carefully by using chemical balances in measuring the volume or mass of gases, liquids and solids that occur in chemical reactions. Therefore, the mass of the reactants and the reaction product can be measured carefully. The results of these experiments present the facts to the observer and demand them into the formulation of the fundamental law that describes the nature of chemistry. The basic law obtained is known as the law of conservation of mass, which is as follows.

'' Mass can not be created nor destroyed in any material change. ''

                  

        The basic legal facts of mass conservation have been proven in 1756 by the Russian scientist M.V. Lomonosov. Perhaps because of language problems, his work is not well known in Western Europe. Separately in 1783, a great French chemist Antoine Lavoisier did the same thing using a chemical balance to show that the sum of the mass of chemical reactions was equal to the amount of reactant mass.

Lavoisier conducts experiments by heating the mrerkuri in a sealed flask filled with air. After a few days, a red substance 

is formed that is mercury (II) oxide. The gas in the mass tube is reduced and can no longer be Supporting the burning (the

 candle in the tube does not fire anymore) and the animal will die if put into it. It shows that the oxygen gas in the tube is gone.

 It is now known that the remaining gas is nitrogen, while the oxygen from the air in the tube has run out with mercury. Furthermore, 

Lavoisier takes the mercury oxide and heats it up so it breaks down again. Then he weighed the mercury and the resulting gas. It 

turns out that the combined mass is equal to the mass of mercury (II) oxide used initially. Finally after several experiments and the 

results are the same, Lavoisier states the law of conservation of mass that is as follows.

   

 '' In every chemical reaction, the mass of substances before and after the reaction is always the same. ''

           Lavoisier was the first to observe that chemical reactions are analogous to algebraic equations.

Example:

       S (s) + O2 (g) → SO2 (g)

1 mol of S reacts with 1 mole O2 to form 1 mole of SO2. 32 grams of S reacts with 32 grams of O2 forming 64 grams of SO2. The total mass of the reactants is equal to the mass of the resulting product.

      H2 (g) + ½ O2 (g) → H2O (l)

1 mole of H2 reacts with ½ moles of O2 forming 1 mole of H2O. 2 grams of H2 reacts with 16 grams of O2 forming 18 grams of H2O. The total mass of the reactants is equal to the mass of the product formed.

2. PROUST LAW OR COMPARATIVE LAW STATUTE

        In 1799 the French chemist Joseph Proust, through various experiments found a provision that was validated by Proust's law, as follows. "The mass ratio of the constituent elements is always fixed, even if it is made in a different way" At that time Proust discovered that copper carbonate, both from natural and synthetic sources in the laboratory, had a fixed arrangement.

        To determine the composition of a compound, we can describe an example of the compound we have weighed, then the compounds are described into their elements. Each compound-forming element we weigh, it was obtained a certain comparison. If it is repeated, it will get the same comparison. Another method can also be done, namely by weighing the mass of compounds formed from the compounds of elements that each element is known mass. Of the many experiments on the composition of elements in the compound, always produce the following statement.

"A pure compound is always composed of fixed elements with a fixed mass ratio."

Example:

S (s) + O2 (g) → SO2 (g)

The ratio of mass S to mass of O2 to form SO2 is 32 grams S to 32 grams of O2 or 1: 1. This means that every gram of S just reacts with one gram of O2 forming 2 grams of SO2. If 50 grams of S is required, 50 graM O2 is required to form 100 gra-SO2.

H2 (g) + ½ O2 (g) → H2O (l)

The ratio of mass of H2 to mass of O2 to form H2O is 2 gram H2 to 16 gram O2 or 1 2 8. This means, Every one gram of H2 is precisely bersiKsi with 8 gram O2 form 9 gram H2O. If provided 24 grams of O2, it takes 1 gram of H2 to form 27 grams of H2O.

3. COMPARATIVE LAW

John Dalton's interest in studying two elements that can form more than one compound turns out to produce a conclusion called the law of multiple comparison:

'' If two elements can form more than one compound, then the ratio of the mass of a single element, which is intertwined with another element of a certain mass is a simple integer ''.

For example, copper with oxygen, carbon with oxygen, sulfur with oxyeen, and phosphorus with chlorine. The mass ratio of the two elements is as follows.

1 Copper and oxygen form two copper oxide compounds.

Copper oxide copper oxygen copper: oxygen

           I 88.8% 11.2% 1: 0.126

           I 79.9% 20.1% 1: 0.252

2 Carbon and oxygen can form two compounds

Carbon + oxygen → Carbon monoxide (I)

Carbon + oxygen → Carbon diocide (II)

Carbon oxygen carbon compound: oxygen

I 42.8% 57.2% 1: 1.33

II 27.3% 72.7% 1: 2,67

3 Sulfur (sulfur) with oxigan can form two oxygen compounds, namely sulfur oxide (I) and sulfur trioxide (II)

Sulfur oxygen sulfur compounds: oxygen

I 50% 50% 1: 1

            

II 40% 60% 1: 1,5

        Up to now this law is still acceptable, but needs to be corrected on simple numbers. If the comparison is a simple number (1, 2, 3, 4, 5) means the compound formula is also simple, such as H2O, CO2, and H2SO4. However, now found compounds with large numbers, such as sucrose and arachidonic acid.

4. COMPARATIVE LAW VOLUME

        The relationship between the volumes of the gases in a chemical reaction was investigated by Joseph Louis Gay-Lussac in 1905. In that study it was found that at a constant temperature and pressure, every single volume of oxygen gas would react with two volumes of hydrogen gas yielding two volumes Water vapor, thus the ratio between the volume of hydrogen, the volume of oxygen and the volume of water vapor in sequence is 2: 1: 2. Another example: one volume of hydrogen gas will react with one volume of chlorine gas yielding two volumes of hydrogen chloride gas; The ratio of hydrogen volume, chlorine volume and volume of hydrogen chloride sequence is 1: 1: 2. In the reaction between the nitrogen gas and the hydrogen gas forming the ammonium gas, the volume ratio of the three gases is 1: 3: 2 (N2: H2: NH3).

Based on the above description, it can be concluded that:

"At the same temperature and pressure, the ratio of the reactant gas volume to the gas volume of the reaction product is a simple integer (equal to the ratio of the reaction coefficient)"

Example:

N2 (g) + 3 H2 (g) → 2 NH3 (g)

The gas volume ratio is equal to the ratio of the reaction coefficient. This means that every 1 mL of N2 gas exactly reacts with 3 mL of H2 gas to form 2 mL of NH3 gas. Thus, to obtain 50 L of NH 3 gas, it takes 25 L of N2 gas and 75 L of H2 gas.

CO (g) + H2O (g) → CO2 (g) + H2 (g)

        The gas volume ratio is equal to the ratio of the reaction coefficient. This means that every 1 mL of CO gas reacts exactly with 1 mL of H2O gas to form 1 mL of CO2 gas and 1 mL of H2 gas. Thus, as much as 4 L of CO gas requires 4 L of H2O gas to form 4 L of CO2 gas and 4 L of H2 gas.

B. THE DALTON ATOM THEORY

        Learning about atomic theory is very important because the atom is the constituent material in the universe. By understanding the atom we can study how one atom interacts with others, know the properties of the atom, etc. so that we can use the universe for the benefit of mankind.

        The name "atom" comes from the Greek word "atomos" introduced by Democritus which means it can not be subdivided or the smallest part of matter that can not be divided again. The concept of an atom which is an indivisible composer of matter was first introduced by Greek and Indian philosophers.

        More modern concepts of atoms emerged in the 17th and 18th centuries where chemistry began to develop. Scientists began using weighing techniques to get more precise measurements and use physics to support the development of atomic theory.

        John Dalton a British teacher used the concept of atoms to explain why elements always react with simple round numbers (hereinafter more commonly known as multiple comparison laws) and why gases are more soluble in water than others. Dalton composes his atomic theory based on the laws of conservation of mass and fixed law of comparison. Where the atomic concept is as follows:

·         Each element is composed of tiny particles called as atoms.The atoms of the same element are identical and the atoms of the element are not different in some basic ways.

·         The chemical compound is formed from a combination of atoms. A compound always has the same number of atoms and types of atoms.

·         Chemical reactions involve the reorganization of atoms that change how they bind but the atoms involved do not change during the chemical reaction.

This Dalton atom model is commonly referred to as a model of billiard balls atoms where different billiard balls are symbols of different elemental atoms

C. AVOGADRO LAW

        Avogadro's Law (Hypotes Avogadro, or Avogadro's Principle) is a gas law named after the Italian scientist Amedeo Avogadro, who in 1811 proposed the hypothesis that:

Gases of the same volume, at the same temperature and pressure, have the same number of particles.

        That is, the number of molecules or atoms in a gas volume does not depend on the size or mass of the gas molecule. For example, 1 liter of hydrogen and nitrogen gas will contain the same number of molecules, as long as the temperature and pressure are the same. This aspect can be expressed mathematically,

Where:

V is the volume of gas.

N is the number of moles in the gas.

K is a matching constant.

The most important consequence of Avogadro's law is that the ideal gas constant has the same value for all gases. That is, constants

Where:

P is the gas pressure

T is temperature

        Has the same value for all gases, independent of the size or mass of the gas molecule. The Avogadro hypothesis is evidenced by the theory of gas kinetics.

        One mole of ideal gas has a 22.4 liters volume under standard conditions (STP), and this number is often called the ideal gas molar volume. Real gases (non-ideal) have different values.

Example: On the formation of H2O molecules

2L H2 (g) + 1L O2 (g) ® 2L H2O (g)

2 molecule H2 1 molecule O2 2 molecule H2O

Note:

If the volume and number of molecules of one substance is known, then the volume and number of molecules of other substances can be determined by using the equation:

And Information :

V = volume of molecule (L)

X = number of particles (molecules)

D. MASS ATOM AND MASS MOLECULES RELATIVES

        The atom is a very small particle so that the atomic mass is also too small when expressed in grams. Therefore, the chemists created a way to measure the mass of an atom, that is, by the relative atomic mass. The relative atomic mass (Ar) is the ratio of the average mass of an atom by one-twelve times the mass of one carbon-12 atom.

The smallest unit of a substance can also be a molecule. The molecule is composed of two or more atoms held together by chemical bonds. The relative molecular mass (Mr) is the average mass ratio of a molecule with one-twelve times the mass of one carbon-12 atom.

        Ar Y = the average mass of 1 molecule Y / (1/12 x mass 1 atom C-12) In the above formula used the atomic mass and the average molecular mass. Why use average atomic mass? Because elements in nature have several isotopes. For example, carbon in nature has 2 stable isotopes of C-12 (98.93%) and C-13 (1.07%). If the abundance and mass of each isotope are known, the relative atomic mass of an element can be calculated by the formula:

        Ar X = {(% isotope 1 x isotope mass 1) + (% isotope 2 x isotope mass 2) + ...} / 100 If it is known that the relative atomic mass of each constituent element of a molecule, its relative molecular mass is equal to the sum of the relative atomic masses of all the molecular compounds of the atom. A molecule having the formula AmBn means that in 1 molecule there are m atoms A and n atoms B. Thus the relative molecular mass of AmBn can be calculated as follows.

Mr. AmBn = m x Ar A + n x Ar B

E. CONCEPT MOL

        In reacting substances, many things we need to consider for example the form of substances in the form of gas, liquid and solid. It is quite difficult for us to react the substances in the three states of matter, in solid form by size in mass (grams), in liquid form by volume of liquids in which there is a solvent and there is a dissolved substance. Similarly, the gaseous material has a gas volume size.

        This condition requires chemists to provide a new unit that can reflect the amount of substances in various substances. Avogadro tried to introduce a new unit called mole. The definition for 1 (one) mole is the number of substances containing particles of 6.023 x 1023. This number is known as Avogadro Numbers denoted by the letter N.

        The above diagram shows the equation which states the relationship of the number of moles to the number of particles for atoms and molecules Considering the mass aspect of the substance, 1 mol of the substance is defined as the mass of the substance corresponding to its relative molecular mass (Mr) or its atomic mass (Ar).

        For 1 mole of Carbon substance it has mass corresponding to Carbon atomic mass, it is known from the periodic table that the mass of carbon atoms is 12 sma, so that the mass of the substance is also 12 grams. For that 1 mole of matter we can change into equation form:

Number of Moles (n)

Mass (m)

Volume Gas (V)

Number of Particles (X)

Mollification (M)

F. EMPIRICAL AND MOLECULUM FORMULAS

        The chemical formula of a substance can explain or state the relative amount of atoms present in the substance. The chemical formula is divided into molecular formulas and empirical formulas. The empirical formula is the simplest formula of a compound. This formula only expresses the ratio of the number of atoms present in the molecule.

The empirical formula of a compound can be determined if one knows:

- mass and Ar each element

- % mass and Ar each element

- comparison of mass and Ar of each element

The molecular formula of a substance describes the number of atoms of each element in a single molecule of that substance.

When the empirical formula is known and Mr. is also known then the molecular formula can be determined.

POLICY

Solubility (M)

ü Molecularity is a way of expressing the concentration (concentration) of the solution.

ü Declare the number of moles of solute in each liter of solution, or the amount of mmol of solute in each mL of solution.

ü Formulated:

ü For example: a 0.2 M NaCl solution means that in each liter of solution there is 0.2 mol (= 11.7 g) NaCl or in each mL of solution there is 0.2 mmol (= 11.7 mg) NaCl.

Dilution formula

V1.M1 = V2.M2

V1 = Volume before dilution (liter)

M1 = Molarity before dilution (M)

V2 = Volume after dilution (liter)

M2 = Molarity after dilution (M)

G. MOLALITY

        The molality mentions the mole ratio of the solute in kilograms of solvent. Molality is expressed between the number of moles of solute and the mass in kg of solvent. What is the symbol of the molality of matter? Molality is symbolized by m

with

N = number of moles of solute ........................ (mole)

P = solvent mass ..................................... (kg)

M = molality ......................................... (mol Kg-1)

H. FRACTION MOL

        The mole fraction is a unit of concentration that expresses the ratio of the number of moles of one of the solution components (the number of moles of solvent or the number of moles of solute) to the total number of moles of the solution. The mole fraction is symbolized by X. For example in a solution containing only two components, namely substance B as a solute and A as a solvent, then the fraction of mole A is symbolized XA and XB for the mole fraction of the solute.

or

With XA = solvent mole fraction

XB = mole fraction of solute

NA = number of moles of solvent

NB = the number of moles of solute

The amount of solvent mole fraction with solute is equal to 1.

XA + XB = 1

I. NUMBERS OF OXIDATION

1. DEFINITION OF NUMBERS OF OXIDATIONS

The oxidation number is the formal charge of the atom in a molecule or in an ion allocated such that a lower electronegativity atom has a positive charge. Since the electrical charge is not different in terms of molecules composed of the same atom, the atomic oxidation number is the quota of net electrical charge divided by the number of atoms. In the case of ions or molecules containing different atoms, atoms with greater electronegativity can be considered anions and the smaller are considered cations. For example, nitrogen is oxidized 0 in N2; Oxygen oxidation -1 in O22-; In NO2 +4 nitrogen and oxygen -2; But in NH3 -3 nitrogen and hydrogen +1. Thus, the oxidation numbers may differ for the same atoms that are combined with different pairs and the atoms are said to have a formal charge equal in value to the oxidation number. Although the price of this formal payload does not reveal the actual charge, it is very convenient for calculating valence electrons and in handling redox reactions.

2. THE CONCEPT OF OXIDATION AND REDUCTION REACTIONS

        In everyday life often encountered chemical reactions that can be classified in oxidation reactions, reduction reactions and oxidation-reduction reactions (redox), such as combustion, kararatan, processing of metal from the seeds.

Based on its development, the concept of oxidation-reduction is explained from the following points:

A. Merging and Spending Oxygen

If a piece of iron is placed in the open air, over time the metal is rusty. Reaction of iron-stamping

Takes place as follows:

4Fe (s) + 3O2 (g) ------> 2Fe2O3

        In corrosive events, iron reacts with oxygen. We say iron is oxidized. The word "oxidation" literally means "oxygenation". Iron rust is an oxide of the formula Fe2O3, such as iron ore on the earth's crust, in the iron ore industry processed into pure iron according to the following reaction:

Fe2O3 (s) + 3CO (g) ------> 2Fe (s) + 3CO2 (g)

        In the manufacture of pure iron, there is the expenditure or reduction of oxygen from iron ore (Fe2O3). We say, Fe2O3 has been reduced. The word reduction literally means "subtraction". So: Oxidation is a merging event in the following reaction equation:

2Cu + O2 ----> 2CuO

2Fe + O2 ----> 2FeO

4Fe + 3O2 ----> 2Fe2O3

Reduction is the process of taking or removing oxygen from a substance.

2FeO + C ----> 2Fe + CO2

CuO + H2 ----> Cu + H2O

B. Electrons Release and Capture

        In the Fe oxidation state to Fe2O3, Fe atoms release the electrons into Fe3 + ions. So the notion of oxidation can be extended to the release of electrons. In contrast to the reduction of Fe2O3 to Fe, Fe3 + ions capture electrons into Fe atoms. Then the notion of reduction can also be expanded into electron capture events. With this broader understanding, the concept of oxidation and reduction is not limited to reactions involving oxygen alone.

Oxidation is the electron-releasing reaction.

Examples of oxidation reactions:

Na ----> Na + + e

Zn ----> Zn2 + + 2e

Fe2 + ----> Fe3 + + e

S2- ----> S + 2e

Reduction is an electron reception or capture reaction.

Contoh reaksi reduksi :
K⁺ + e ---- >K
Cu²⁺ + 2e ---->Cu
Co³⁺ + e----> Co²⁺
Cl₂ + 2e ---->2Cl^-

Examples of reduction reactions:

K ++ e ----> K

Cu2 + + 2e ----> Cu

Co3 + + e ----> Co2 +

Cl2 + 2e ----> 2Cl-

C. Oxidation-Reduction Based on Oxidation Numbers

Oxidation = Addition (rise) oxidation number

Reduction = Reduction (down) oxidation number

Oxidation number: a number indicating the ability of an atom to bind or disconnect electrons

Example:

Fe2O3 (s) + .... 3CO (g) → 2Fe (s) + .. 3CO2 (g)

+3........+2......0.......+4

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