using english to calculate

Stoichiometry

Stoichiometry is the basic chemical calculation that states quantitative relation of chemical formulas and chemical equations. In chemistry, stoichiometry (sometimes called stoichiometry of reaction to distinguish it from compositional stoichiometry) is a science that studies and quantifies quantitative relationships of reactants and products in chemical reactions (chemical equations). This word comes from the Greek stoikheion (element) and metriā (size). Stoichiometry is based on the basic laws of chemistry, namely the law of conservation of mass, the law of fixed comparison, and the law of multiple comparisons

Early Stoichiometric Stage

In early chemistry, the quantitative aspect of chemical change, ie stoichiometry of chemical reactions, did not receive much attention. Even when attention has been given, experimental techniques and tools do not produce the correct results.

One example involves the theory of flogstones. Flogistonis tried to explain the phenomenon of burning with the term "combustible substances". According to the flogitonists, combustion is the release of a substance can be etrbakar (from a burning substance). This substance is then called "flogiston". Based on this theory, they define combustion as a flogiston release of a combustible substance. Changing the mass of wood when burning fits well with this theory. However, changes in metal mass when calcined do not match this theory. Nevertheless flogistonis accept that both processes are essentially identical. Increasing the mass of calcined metals is a fact. Flogistonis attempts to explain this anomaly by stating that flogiston has a negative mass.

Philosopher from Flanders Jan Baptista van Helmont (1579-1644) conducted a famous "willow" experiment. He grows willow seeds after measuring the mass of the flowerpot and the soil. Since there is no change in the masses of flower pots and soil when the seeds grow, he assumes that the mass obtained only by the water entering the ore. He concludes that "the root of all matter is water". Based on the current view, the hypothesis and experiments are far from perfect, but the theory is a good example of the growing demeanor of quantitative chemical aspects. Helmont recognized the importance of stoichiometry, and clearly preceded his time.

In the late 18th century, German chemist Jeremias Benjamin Richter (1762-1807) discovered the equivalent concept (in terms of modern chemical equivalent chemistry) with a careful observation of the acid / base reaction, ie the quantitative relationship between acid and base in the neutralization reaction. Richter equivalents, or what are now called chemical equivalents, indicate a certain amount of matter in the reaction. An equivalent in neutralization relates to the relationship between a number of acids and a number of bases to neutralize them. Proper knowledge of equivalents is essential to produce good soap and gunpowder. So, this kind of knowledge is very important in practical terms.

At the same time Lavoisier establishes the law of conservation of mass, and provides the basic concepts equivalent to its accurate and creative experiments. Thus, the stoichiometry that handles quantitative aspects of chemical reactions becomes a basic chemical methodology. All fundamental laws of chemistry, from the law of conservation of mass, the law of comparison remained until the laws of gas reaction were all grounded stoichiometry. These fundamental laws form the basis of atomic theory, and are consistently explained by atomic theory. However, it is interesting to note that, the equivalent concept was used before the atomic theory was introduced.

A. Basic Law of Chemistry

1. Understanding Basic Law of Chemistry

The basic law of chemistry is the basic law governing the mechanism of the occurrence of a chemical reaction involving reactants and products.

2. Basic Chemical Laws

The basic law of chemistry consists of 5 basic laws such as mass conservation law, fixed comparison law, multiple comparison law, volume comparison law and avogadro law.

ü  The Basic Law of Conservation of Mass

As the name suggests, this law was discovered by a scientist named Antonie Lavoiser. In his law, he states that "The total mass of matter after reaction and before reaction is the same". The point of his statement is that "the sum of the mass of a substance acting as a reactant is equal to the sum of the mass of a substance acting as a product".

The point is → Reactant Mass = Mass Product

ü  Comparative law remains = law proust

"The ratio of the mass of the elements in each compound is fixed"

Example:

          A. In the compound NH3 ------ mass N: mass H = 1 Ar. N: 3 Ar. H

                                                                                       = 1 (14): 3 (1)

                                                                                      = 14: 3

          B. On SO3 compound ------ mass S: mass O = 1 Ar. S: 3 Ar. O

                                                                                        = 1 (32): 3 (16)

                                                                                              = 32: 48

                                                                                              = 2: 3

Advantages of Proust Law:

If known mass of a compound or mass of one element that make up the compound then the mass of other elements can be known.

ü  Law of multiple comparisons = dalton's law

"If two elements can form two or more compounds for the mass of one element equal to the number then the ratio of the mass of the second element will be proportional to the integer and the simple".

Example:

When the element of Nitrogen with compounded oxygen can be formed,

       NO where the mass N: O = 14: 16 = 7: 8

       NO2 where the mass N: O = 14: 32 = 7: 16

     For the same mass of Nitrogen the Oxygen mass ratio of the compound

NO: NO2 = 8: 16 = 1: 2

Besides that, another example that is Nitrogen and Oxygen can form six kinds of compound. The ratio of the weight of oxygen that reacts with one part of nitrogen is:

0.57: 1.14: 1.74: 2.28: 2.86: 3.42

1: 2: 3: 4: 5: 6

This comparison is an easy & round number, so it is in accordance with Comparative law.

ü  Law of the Gay / Lussac Volume / Legal Comparison

At the same pressure and temperature, the volume of the reacting gas and the gas volume of the reaction product are simple and integer ratios. This is the sound of the law triggered by our legendary scientist named Gay lussac. His theory was not without evidence or without research. To prove his theory, he conducted a simple experiment by reacting hydrogen gas with oxygen gas into a container, then into the container was given an electric flower flow so that oxygen gas and hydrogen gas can react. After the reaction is complete, water vapor is generated as a product and residual H2 and O2 gas are not reacted. After that, the resulting water vapor is directly separated from inside the container. The experiments were conducted repeatedly at fixed temperature and pressure and the measurements showed that the ratio of the volume of hydrogen and oxygen gas and water vapor was always 2: 1: 2.

2 H2 + O2 → 2H2O

Comparison of coefficient numbers = 2: 1: 2 (the coefficient numbers are integers and simple)

ü  Avogadro's Law

At the same temperature and pressure, a gas having the same volume also has the same number of particles, This is the statement of the oldest legendary scientist, Mr. Avogadro. The point of his statement is that the amount of particles of a gas is independent of the Mass Or Mr. possessed by the gas, while in the same volume, temperature and pressure, the amount of particles of a gas will always be the same. The analogy is this: 1 liter of nitrogen gas and 1 liter of chlorine gas have the same number of particles while under the same pressure and temperature. So when a gas is in the same temperature, volume and pressure the amount of particles from the gas will always be the same.

B. The Concept of Molar Mol and Mass (})

In SI systems, one mole is defined as the sum of the material composed of entities (atoms, molecules, or other particles) a sum of the atoms in 12 grams of carbon-12. The value of the number of atoms is 6.022 × 1023 called the Avogadro number, NA.

Dalton recognizes that it is important to determine the mass of each atom because its mass varies for each type of atom. Atoms are so small that it is impossible to determine the mass of one atom. So he focused on the relative mass values ​​and made the atomic mass table (figure 1.3) for the first time in human history. In the table, the lightest element mass, hydrogen set one as standard (H = 1). The atomic mass is a relative value, meaning a dimensionless ratio. Although some atomic masses differ from modern values, most of the proposed values ​​are in the range of compatibility with current values. This shows that his ideas and experiments are correct.

ü  Mr. Dan Ar

Understanding Relative Atomic Mass (Ar)

The relative atomic mass is the average mass of an atom divided by the mass of the reference atom C-12 atoms.

ü  Relative Molecular Mass (Mr)

The realmic molecular mass is the average mass of the molecule divided by the mass of the reference atom C-12 atoms.

Empirical Formulas And Molecular Formulas

Molecular formula is a formula that states the number and type of atoms that make up a compound, for example water, water has the formula H2O molecule which means the water compound is composed by two types of atoms H and O atoms and the number of each atom consists of 2 H atoms And 1 O atom, so this is what is meant by the molecular formula.

While the empirical formula is not much different from the molecular formula, since both show the number and type of atoms that make up a compound, the difference in this empirical formula comparison of the number of atoms simplified as small as possible, for example in benzene compounds, benzene has the formula C6H6 molecule, to make The empirical formula then each number of atoms must be divided by 6 for the ratio of the number of atoms to be more simple, so the empirical formula of the compound C6H6 is CH.

ü  Hydrate Compounds

Hydrate compounds are one of the chemical compounds that molecules bind to water molecules. Analoginya like this, the chemical compound is like a cotton that can absorb water / moisture around the cotton, the more water is absorbed, the more water molecules contained in the cotton

The concept of Mol

Mol is the number of substances of an element that contains some form of elements such as atoms, molecules, ions or electrons. Here is the relationship of mole with several categories including:

• The Relation of Moles With Particles

• The connection of Mol With Mass

• Relation of Mol With Volume

• The Relation of Moles With Chemical Reactions

Problem No. 1

Determine the number of moles contained in:

A) 96 grams of oxygen (O2)

B) 88 grams of carbon dioxide (CO2)

(Mr. O2 = 32; Mr. CO2 = 44)

Discussion

Determine the known mol mass of the substance.

A) 96 grams of oxygen (O2)

M = 96 grams

Mr = 32

N = ....

N = m / mr

N = 96/32 = 3 mol

B) 88 grams of carbon dioxide (CO2)

M = 84 grams

Mr = 44

N = ....

N = m / mr

N = 88/44 = 2 mol

Problem No. 2

Determine the amount of mass contained in:

A) 0.2 mol of oxygen (O2)

B) 0.04 moles of carbon dioxide (CO2)

(Mr. O2 = 32; Mr. CO2 = 44)

Discussion

Determining the mass

A) 0.2 mol of oxygen (O2)

N = 0.2 mol

Mr. O2 = 32

M = ....

M = n × Mr

M = 0.2 × 32

M = 6.4 grams

B) 0.04 moles of carbon dioxide (CO2)

N = 0.04 mol

Mr. CO2 = 44

M = ....

M = n × Mr

M = 0.04 × 44

M = 1.76 grams

Problem No. 3

Determine the number of particles contained in

A) 0.1 mol of oxygen (O2)

B) 0.02 mol of carbon dioxide (CO2)

Discussion

Determine the number of particles

A) 0.1 mol of oxygen (O2)

N = 0.1 mol

X = ......

X = n × 6.02 × 1023

X = 0.1 × 6.02 × 1023

X = 6.02 × 1022 particles

B) 0.02 mol of carbon dioxide (CO2)

N = 0.02 mol

X = ......

X = n × 6.02 × 1023

X = 0.02 × 6.02 × 1023

X = 1.204 × 1022 particles