Stoichiometry is the basic chemical calculation that states quantitative relation of chemical formulas and chemical equations. Here are the materials you need to know to understand, from the concept of moles and molar masses, empirical formulas and molecular formulas, basic stoichiometry of solutions and ideal gases, and the writing and equalization of reactions, with examples of problems and discussions.
Molar and Molar Mass Concepts (})
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.
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.
The molar mass,},
is defined as the mass of 1 mol entity (atom, ion, molecule, unit of formula)
of matter. The unit of the molar mass (}) is a gram / mol.
m
= n
N
= nNA
Stoichiometric
Material: Mass relation, number of moles, and number of atoms of the element
(Source: Chang, Raymond, 2010. Chemistry (10th edition) New York: McGraw Hill)
(Source: Chang, Raymond, 2010. Chemistry (10th edition) New York: McGraw Hill)
Empirical Formulas and Molecular Formulas
The empirical formula is the simplest integer ratio of the number of moles of each element in a compound. The molecular formula represents the true number of moles of each element in 1 mole of the compound. The molecular formula may be identical to the empirical formula or an integer multiple of the empirical formula. For example, phosphoric acid (H3PO4) has a molecular formula and an identical empirical formula. Glucose has a molecular formula C6H12O6 which is a folding of 6 times its empirical formula, CH2O.
The empirical formula is the simplest integer ratio of the number of moles of each element in a compound. The molecular formula represents the true number of moles of each element in 1 mole of the compound. The molecular formula may be identical to the empirical formula or an integer multiple of the empirical formula. For example, phosphoric acid (H3PO4) has a molecular formula and an identical empirical formula. Glucose has a molecular formula C6H12O6 which is a folding of 6 times its empirical formula, CH2O.
Molecular formula ≡
(empirical formula) n
{Molecular formula = n ×} empirical formula, n = 1, 2, 3, ...
{Molecular formula = n ×} empirical formula, n = 1, 2, 3, ...
Basic Stoichiometric Solution
The term "concentration" of the solution expresses the amount of solute dissolved in a certain amount of solvent or a certain amount of solution. The concentration of the solution can be expressed in molarity. Molarity (M) is defined as the number of moles of solute per liter of solution.
The term "concentration" of the solution expresses the amount of solute dissolved in a certain amount of solvent or a certain amount of solution. The concentration of the solution can be expressed in molarity. Molarity (M) is defined as the number of moles of solute per liter of solution.
Basic Ideal
Gas Stoichiometry
The molar volume, Vm, is defined as the volume of 1 mol entity (atom, ion, molecule, unit of formula) of matter. The unit of the molar volume (Vm) is L / mol.
The molar volume, Vm, is defined as the volume of 1 mol entity (atom, ion, molecule, unit of formula) of matter. The unit of the molar volume (Vm) is L / mol.
Avogadro's law states that at certain pressures and
temperatures and fixed, the volume of the gas is directly proportional to the
amount of gas.
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In the case of RTP / ATP (P = 1 atm, T = 298 K), ideal gas Vm = 24 L / mol
In certain circumstances, an ideal gas law applies:
Where R is the gas constant, R =
0.08206 L ∙ atm / mol ∙ K = 8.314 J / mol ∙ K
Writing and Equating Chemical Reaction Equations
The chemical reaction equation is a statement written with a molecular formula that provides information on the identity and quantity of substances involved in a chemical or physical change. All reacting substances, called reactants, are placed to the left of the arrow, whose direction of the arrow to the right refers to the product, ie all the substances produced from the reaction.
In the equation of the reaction, the reaction coefficient is known, which is the number on the left of the molecular formula for multiplying all the atoms in the molecular formula. The comparison of the reaction coefficients can be interpreted as the mole ratio of the substances in the reaction. In each of the reactants and products, it is written in the form of the substance (s (solid), l (liquid), g (gas), or aq (solution with water solvent)) in brackets to the right of each molecular formula.
Example of equivalent chemical reaction equation:
1.
2.
3.
4.
Stoichiometry Reaction
In chemical reactions, the amount of reacting reactants is sometimes incompatible with the stoichiometric amount of the reaction (not in accordance with the coefficient ratio of the equivalent equation). Therefore, there will be reactants that have reacted first compared to other reactants. The reactants that remain after reacting are called excess reagents. The out-of-date reactant is called a limiting reagent. After the barrier reagents are exhausted, no more reaction products are formed. Thus, the number of limiting reagents determines the amount of product produced.
The principle that underlies stoichiometry?
BalasHapusLaws and Principles of Chemistry
HapusIn general, there are three most important and influential laws, with the most basic first law, namely:
· The Law of Conservation of Mass
· Comparable Law of Comparison
Multiple Multiple Laws
The Law of Conservation of Mass
The law of mass conservation was first proposed by Lomonov and Antoine Lavoisier rediscovered (both unrelated). This law explains that all reactions, in closed systems, of mass of substances before and after the reaction are the same. Nothing was created or destroyed. Everything is just a new combination. Lavoisier conducted the following experiments. He heats a piece of lead in a sealed tube. Gray ash appears on the tin surface. He heated it for a day and a half until the ashes did not appear. Then, he weighs the mass of the tin-covered tube before being burned with a burned tube (which contains lead oxide, tin left, and oxidized air). The result, he found that the mass before and after the reaction is the same.
His discovery was a big leap for the chemistry world because he was the first to prove air involvement in the reaction. However, in the 20th century, this law was revised to accommodate nuclear and relativity conditions. This law remains the basis for many important laws such as equilibrium, thermodynamics, and the basis of stoichiometry.
Fixed Comparison Law
Comparative law is still found by Joseph Proust, a French chemist. Comparative law still states, as the name implies, the ratio of the mass of elements in a compound is certain and fixed. Thus, any compound anywhere must consist of a definite mass ratio. For example, the ratio of sodium and chlorine mass to NaCl by 2 grams is 0.768 grams and 1,124 grams. Then the mass ratio is 1: 1.54 or simplified 2: 3. If the same compound is taken from another source as much as 2.5 grams with sodium 0.983 gram, then found 0.983: 1.517 or 1: 1.54 or 2: 3.
This law breaks the opinion Archimedes used chemists from Arabs to Europe for hundreds of years, that the compound is merely the origin of the mix with the original comparison. Although thereafter found a very small error, this law opens the path of developing the reaction of compounds in modern chemistry.
Multiple Multiple Laws
At the time of this law, the chemical formula of the compound is unknown. This law was put forward by John Dalton, a British chemist and inventor of modern atomic theory. This law states that if the mass of one element in two compounds is equal, then the ratio of the mass of the other elements is a simple integer. For example, the ratio of carbon (C) and oxygen (O) elements to sequential carbon monoxide and carbon dioxide is 3: 4 and 3: 8. If the mass of C is equal, then the ratio of mass O to carbon monoxide and carbon dioxide is 4: 8 or 1: 2.
It should be noted, however, that this law is the development of Proust's law, although it was found before Proust's own law. This law also states that atoms can not be shaped like half shards, must be integers. This law is strong because it is supported by atomic theory.
In addition to the three previously mentioned, there are still other theories such as Avogadro Hypothesis, Dalton atomic theory, equilibrium theory, Hammond-Leffler Hypothesis, and others. Yet these three laws underlie them.
Give examples of exaggerated reagents and examples limiting reagent
BalasHapusReagents or also known as Reactan is a term that is often used in the chemical world. Reagents have many uses and most involve lifesaving applications. Substances or two substances make, measure or build up the existence of chemical reactions with the help of reagents. Organic chemistry may also define reagents as mixtures or different substances that will make changes to the substrate under certain conditions.
HapusExamples of natural reagents:
* Fenton reagents - reagent-style reagent reagents are utilized to eradicate certain natural and organic chemicals such as tetrachloroethylene (PCE) and trichlorethylene (TCE).
* Grignard reagents - reagents in this kind are specially prepared when using a response produced from a mixture of alkyl and magnesium. Organic compounds all need this particular chemical reaction to create carbon-carbon bonds
Collins reagents - these reagents are used to help some complex substances and alcohols to oxidize.
Fehling reagents - this is a solution of sodium hydroxide, copper sulphate and potassium sodium tartrate which are specially used to test the presence of aldehydes and sugars in certain substances, such as those of urine.
* Millon reagents - investigative reagents in this unique type are made by melting Mercury's metals with nitric acid and then watering down to get the desired density. Reagents are substances used to detect protein soluble.
In a chemical reaction, the mole ratio of the added reagents is not always the same as the ratio of the reaction coefficient. This causes a reagent to be reacted first. This is called pereaksi pembatas. Limiting reagent is a reactant contained in the relatively smallest amount (in the stoichiometric relationship). The limiting reagents will run out, while the other reactions will leave the rest.
Example:
One mole of sodium hydroxide solution (NaOH) is reacted with 1 mole of sulfuric acid (H 2 SO 4) solution according to the reaction
2 NaOH (aq) + H2SO4 (aq) -> Na2SO4 (aq) + 2 H2O (l)
Resolution:
Mol each substance is divided coefficient, then select the small divide as a limiting reagent
- mole NaOH / NaOH coefficient
= 1/2 mol
= 0.5 mol
- mol H2SO4 / H2SO4 coefficient
= 1/1 mol
= 1 mol
Since the yield is NaOH Na2SO4 (aq) + 2 H2O (l)
First: 1 mol 1 mol 0 0
Reacts: (2x0.5) = 1 mol (1x0.5) = 0.5 mol
Residual: 1-1 = 0 moles 1-0,5 = 0.5 mole 0.5 mole 1 mole
The remaining reagent is H2SO4
What is "restricting reagent" and give me an example?
BalasHapusThe Limiting Reagent is a completely discharged reactant that determines when the reaction stops.
Hapusexample:
In a closed container, 20 grams of methane (CH4) burned with 64 grams of oxygen (O2) produces carbon dioxide and water vapor according to the reaction below. Known Mr. methane = 16, Ar oxygen = 16, and Mr. H2O = 18.
CH4 (g) + O2 → CO2 + 2H2O
Determine the limiting reagents
Answer:
Resolve the Reaction (already equivalent → given in the matter)
CH4 (g) + 2O2 → CO2 + 2H2O
Calculate the moles of each reactant
Mol CH4 = 20/16 = 1.25 mol
Mol O2 = 64/32 = 2 mol
Adjust the reaction mole and calculate the result of the reaction
1.25 mol CH4 → 2.5 mol H2O (comparison coefficient 1: 2)
2.5 x 18 = 45 grams of H2O
2 mol O2 → 2 moles H2O (2: 2 mole ratio = 1: 1)
2 x 18 = 36 grams of H2O
So which is a limiting reagent is oxygen (O2)
Provide an example of the Molar concept
BalasHapusThe molar mass is the mass of the substance that is equal to the atomic mass or the mass of the formula of the substance expressed in grams.
HapusIn the mole discussion, we have described the relationship between mass and the number of particles and the relationship between the number of particles and units of moles of substances using the Avogadro constant. These two relationships can be used to express substances in units of grams and mol units, and can be used to connect between grams and moles by applying relative atomic mass or the relative molecular mass of the substance.
An example of a molar mass
How many O2 molecules are contained in 16 grams of O2? Given Mr. O2 = 32.
Resolution:
O2 molar mass = 32 gr per mol
The number of moles of oxygen = (16 g O2) / (32 g / mol O2) = 0.5 mole
The number of O2 molecules in 0.5 mol = 0.5 mol of O2 x 6.02 x 1023 molecule per mol = 3.01 x 1023 molecule O2.