PERTEMUAN13_ STOICHIOMETRY

STOICHIOMETRY

Stoichiometry comes from the Greek word stoicheion which means element and metron which means measure. Stoichiometry discusses the relation of mass between elements in a compound (stoichiometric compound) and interactivity in a reaction (reaction stoichiometry). Mass measurements in chemical reactions were initiated by Antoine Laurent Lavoisier (1743 - 1794) who found that in chemical reactions there was no mass change (mass conservation law). Furthermore Joseph Louis Proust (1754 - 1826) found that the elements form compounds in certain comparisons (fixed comparison law). Furthermore, in order to construct his atomic theory, John Dalton discovered the third basic chemical law, called the law of multiples of comparison. These three laws are the basis of the first chemical theory, the atomic theory proposed by John Dalton around 1803. According to Dalton, every material consists of atoms, elements composed of similar atoms, whereas compounds composed of different atoms in certain comparisons . However, Dalton has not been able to determine the ratio of the atoms in the compound (the chemical formula of the substance). Determination of chemical formula of substances can be done thanks to the discovery of Gay Lussac and Avogadro. Once the chemical formula of the compound can be determined, then the ratio of antaratome (Ar) and intermolecular (Mr) masses can be determined. Knowledge of relative atomic mass and chemical formula of compounds is the basis of chemical calculations.

Basic Laws of Chemistry

1. The Law of Conservation of Mass (Lavoisier Law)

"The mass of the substance before the reaction equals the mass of the substance after the reaction" 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. Comparable Law (Proust Law)

"The mass ratio of the constituent elements is always fixed, even if it is made in a different way" 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 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, it takes 50 grams of O2 to form 100 grams of 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 of O2 or 1: 8. This means, every one gram of H2 precisely reacts with 8 gram of O2 forming 9 gram H2O. If provided 24 grams of O2, it takes 3 grams of H2 to form 27 grams of H2O.

3.The Law of Multiple Comparisons (Dalton's Law)

 The Law of Multiple Comparisons: If an element reacts with other elements, then the ratio of the weight of the element is a simple integer

 So from persmaaan:

2Na (s) + 2HCl (aq) → 2NaCl (aq) + H2 (g)

We can know that 2 moles of HCl react with 2 moles of Na to form 2 moles of NaCl and 1 mole of H2. By equalizing this reaction, it can be known the quantity of each substance involved in the reaction.

Hence the equalization of this reaction is very important in solving stoichiometric problems.

Example:

Lead (IV) Hydroxide reacts with Sulfuric Acid, by reaction as follows:

Pb (OH) 4 + H2SO4 → Pb (SO4) 2 + H2O

If we look good either:

Reactant Element

(Number of moles) Product

(Number of moles)

Pb 1 1

O 8 9

H 6 2

S 1 2

Then this equation is not equivalent. Therefore we need to equate this equation. In the reactant there are 16 atoms, but in its product there are only 14 atoms. This equation needs to add coefficients so that the number of atoms of the elements is the same.

In front of H2SO4 it is necessary to add coefficient 2 so that the number of sulfur atoms corresponds, then in front of H2O it is necessary to add coefficient 4 so that the number of oxygen atoms is appropriate. Then the equivalent reaction is:

Pb (OH) 4 + 2H2SO4 → Pb (SO4) 2 + 4H2O

Reactant Element

(Number of moles) Product

(Number of moles)

Pb 1 1

O 12 12

H 8 8S 2 2

The condition in which the equation of the reaction is equal is when it satisfies the following two criteria:

1. The number of atoms of each element on the left and right sides of the equation has been the same.

2. The number of ions on the left and right has been the same (using redox reaction equation)

3. Comparative Law of Volume (Gay Lussac Law)

Applies only to chemical reactions that involve the gas phase

"At the same temperature and pressure, the ratio of 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.

4. Avogadro's Law

Applies only to chemical reactions that involve the gas phase

"At the same temperature and pressure, the same volumes of gas contain the same number of moles" .Avogadro's law is closely related to Gay Lussac's Law : Example:

N2 (g) + 3 H2 (g) → 2 NH3 (g)

The mole ratio is equal to the ratio of the reaction coefficient. This means that every 1 mole of precise N2 gas reacts with 3 moles of H 2 gas to form 2 moles of NH 3 gas. The gas volume ratio is equal to the ratio of the reaction coefficient. This means that every 1 L of N2 gas precisely reacts with 3 L of H 2 gas to form 2 L of NH3 gas. Thus, if at a certain temperature and pressure, 1 mole of gas is equivalent to 1 L of gas, then 2 moles of gas is equivalent to 2 L of gas. In other words, the mole gas ratio is equal to the ratio of gas volume.

Here are some examples of problems and chemical calculations that use the basic laws of chemistry:

1. 20 gram calcium powder (Ar Ca = 40) is reacted with 20 grams of sulfur (Ar S = 32) according to the Ca + S → CaS reaction equation. What substance is left after the reaction is completed? How much mass of the substance is left after the reaction is complete?

Resolution:

The ratio of Ca mole to S is 1: 1. This means that every 40 grams of Ca exactly reacts with 32 grams of S to form 72 grams of CaS. The ratio of mass of Ca to S is 40: 32 = 5: 4.

If 20 grams of S just exhausts, it takes (5/4) x 20 = 25 grams Ca, to form 45 grams of CaS. Unfortunately, the amount of Ca provided is insufficient.

Therefore, 20 grams of Ca will be properly reacted. The required S mass of (4/5) x 20 grams = 16 grams. Thus, the remaining substance is sulfur (S). The remaining sulfur mass is 20-16 = 4 grams.

Equalization of Chemical Reaction

Chemical reactions are often written in bentu equations using element symbols. The reactants are the substances that are on the left, and the product is the substance that is on the right, then both are separated by arrows (can be one or two alternating arrows). Example:

2Na (s) + HCl (aq) → 2NaCl (aq) + H2 (g)

The equation of a chemical reaction is like a prescription in the reaction, thus indicating everything associated with the reaction, whether it is an ion, an element, a compound, a reactant or a product. All.

Then as in the recipe, there is a proportion of the equation shown in the figures in front of the molecular formula.

When considered again, the number of H atoms on the reactant (left) is not equal to the number of H atoms on the product (right). Then this reaction needs to be synchronized. The equalization of chemical reactions must satisfy some chemical laws of matter.

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