Category: Chemistry
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What is Limiting Reagent? How to Calculate...
Continue ReadingWhat is Limiting Reagent?
Limiting reagent in a reaction is the first to be fully used up and thus stops any further reaction from occurring.
In a given chemical reaction, the limiting reagent, or it is also known as limiting reactant, is the substance that has been fully consumed when the chemical reaction is completed.
With the help of stoichiometry, the exact amount of reactant which is needed to react with another element can be calculated.
Though, if the reagents are not mixed or are present in these precise stoichiometric proportions, then the limiting reagent will be fully consumed and the given reaction will not go to stoichiometric completion.
Thus, in simple terms, the reactant that is fully used up in a reaction is called a limiting reagent.
This reactant called limiting reagent generally governs when the reaction will stop.
The exact amount of reactant which will be required to react with another element can be calculated from the reaction stoichiometry.
The limiting reagent depends only on the mole ratio, not on the masses of the reactants present in the given chemical equation.
1. Limiting Reagent Examples
The formation of ammonia or NH3
3H2 + N2 → 2NH3
In the reaction given above, 3 moles of Hydrogen gas are reacted with 1 mole of nitrogen gas to form 2 moles of ammonia but from the above-mentioned balanced chemical equation, one mole of N2 requires only three moles of H2.
Thus, the limiting reagent in the above-mentioned reaction is H2.
2. Limiting Reagent Examples
Given below is another example of limiting reagent; Assume 1 mol of oxygen and 1 mol of hydrogen are present in a chemical reaction.
The following equation for the same is given below;
2H2 + O2 → 2H2O
Since the above reaction uses up hydrogen twice as fast as oxygen, the limiting reactant, in this case, would be hydrogen.
How to Find Limiting Reagent?
In most limiting reactant stoichiometry problems, the actual purpose is to determine how much product could be formed from a specific reactant mixture.
Thus, the limiting reactant or reagent can be determined by two methods mentioned below;
1. By using the mole ration
2. Using the product approach
First, to calculate the mass of the product, write the balanced equation and then find out which reagent is present in excess quantity. Then, by using the limiting reagent calculate the mass of the product formed in a chemical reaction.
1. How to Identify Limiting Reagent?
The following points given below should be considered while trying to identify the limiting reagent:
• When only two reactants are given in a chemical reaction, first write the balanced chemical equation and then check the amount of reactant B which is required to react with the given reactant A. When the quantity of reactant B is greater, then reactant A is the limiting reagent in that particular reaction.
• The reactant which is present in a lesser amount in a chemical reaction than required by stoichiometry is thus the limiting reactant.
• In another method of finding the limiting agent is by calculating the amount of product formed by each reactant.
• The limiting reactant is the reactant from which the least amount of product is formed in a chemical reaction Thus, the required limiting reagent for the reaction can be recognized using the points mentioned above. These reagents are a vital part of a chemical reaction while calculating the percentage yield of a given chemical reaction.
2. How to Identify Limiting Reagent?
We can also find the limiting reagent by observing the number of moles of each reactant in a given chemical equation.
Steps are mentioned below;
1. Firstly, determine the balanced chemical equation for the given chemical reaction.
2. Then, convert all given data into moles (most probably, with the use of molar mass as a conversion factor).
3. Now, calculate the mole ratio from the given data. Then, compare the calculated ratio to the actual ratio.
4. Use the amount of limiting reactant thus, to calculate the amount of product formed in the chemical reaction.
5. If needed, calculate how much is left in the surplus of the non-limiting reagent.
Limiting Reactant and Theoretical Yield
A reactant in a chemical reaction can also limit the quantities of products formed by the given reaction. When this occurs, we refer to the reactant as the limiting reactant (or limiting reagent) and the amount of a product formed when the limiting reactant is entirely consumed in a reaction is known as the theoretical yield.
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What is Stoichiometry? Definition, Calculation, & Worksheet
Continue ReadingWhat is Stoichiometry?
In simple words, Stoichiometry is defined as the calculation of products and reactants in a given chemical reaction. It is principally concerned with numbers.
The word “ stoichiometry” originated from the Greek word “stoikhein” which means element and the word “metron” means to measure.
This term Stoichiometry was first proposed by a well-known German chemist Jeremias Richter.
Stoichiometry is a significant concept in chemistry that helps us balanced calculate the quantities of reactants and products in a given chemical equation.
For instance, oxygen and hydrogen react to form a substance called water in which one mole of oxygen reacts with two moles of hydrogen and hence forms two moles of water.
In addition to this, stoichiometry can also be used to find quantities such as the amount of products that can be formed with a given quantity of reactants and percent yield.
Importance of Stoichiometry
Stoichiometry helps us govern how much substance is needed or is present in the given compound.
Things that can be measured using stoichiometry are;
1. Reactants’ mass and Products mass
2. The molecular weight of a given compound
3. Chemical equations
4. Formulas of different compounds or elements
Stoichiometric Coefficient or Stoichiometric Number
Stoichiometric coefficient or stoichiometric number is defined as the number of molecules that participate in the given chemical reaction. A balanced equation has an equal number of elements on both sides. The stoichiometric coefficient is mostly the number that is present before atoms, molecules, or ions.
Stoichiometric coefficients can be in the form of fractions as well as whole numbers. In addition to that, the stoichiometric coefficients help to establish the mole ratio between reactants and the products of a given chemical equation.
Stoichiometry in Chemical Analysis
Stoichiometric calculations follow a numerical analysis methodology that is frequently used by pharmacists to determine the concentrations of materials present in a given sample.
There are chiefly two main types of analysis given below;
1. Gravimetric Analysis
Gravimetric analysis describes the numerical determination of the analyte based on the mass of the given solid compound.
The gravimetric analysis gives the most precise results as compared to other analytical analyses Gravimetric analysis can be categorized into the following types mentioned below;
1. Precipitation gravimetry involves isolation of ions in a given solution by a precipitation method then filtering, washing the precipitate, and lastly weighing the precipitate and determining its mass by difference.
2. Volatilization gravimetry involves separating components of a mixture either by heating or chemically decomposing the given sample.
3. Electrogravimetry involves the electrochemical reduction of metal ions at the cathode and the immediate deposition of ions on the cathode.
This cathode is weighed before and after the electrolysis is done and the weight, difference thus parallels to the mass of analyte primarily present in the given sample.
2. Volumetric Analysis
The volumetric analysis involves the numerical measurement of a substance in terms of its volume.
In volumetric analysis, a known volume (V1) of the given substance whose concentration (N1) is known is made to react with the unknown volume (V2) of the given solution of the substance whose concentration(N2) is to be calculated. The volume, V1 is defined as the endpoint of the reaction.
Thus, the concentration N2 is calculated with the help of following equation mentioned below;
N1x V1 = N2 x V2
The endpoint of such a reaction is directed by a change in a color with the help of indicators or precipitation etc.
Stoichiometric Values in a Chemical Reaction
In some cases, it might be necessary to calculate the number of moles of a reagent or product under certain given conditions. To do this properly, the reaction needs to be balanced.
The law of conservation of matter thus states that the quantity of each element does not change in a given chemical reaction. Consequently, a chemical equation is balanced to make the quantity of each element in the equation the same.
Chemical reactions are balanced by adding coefficients in front of the given reactants and products.
Example: Balance the given equation;
N2 + O2 → NO
The given equation is not considered balanced since there are more N and O atoms on the left side of the equation as compared to the right side or product of the reaction. Allocate a stoichiometric coefficient of 1 to the given compound, NO.
N2 + O2 → 1NO
Now, balance the remaining element. To do this, we used fractional coefficients.
1/2N2 + 1/2O2 → 1NO
To get rid of the fractional coefficients multiply the reactants by 2 to this equation balanced.
N2 + O2 → 2NO
Thus, the above equation is balanced.
Limiting Reagent
In a chemical reaction, if any one of the reactants is present in a surplus amount then some of these excess reactants will, consequently, be left over when the given reaction is completed.
The given reaction will stop instantly as soon as one of the reactants is fully consumed. The substance that is fully consumed in a given reaction is known as a limiting reagent.
Example;
N2 + 3H2 ➝ 2NH3
Assume we have one mole of N2 that reacts with one mole of H2. But from the above-mentioned balanced chemical equation, one mole of N2 requires only three moles of H2. Thus, the limiting reagent in the above reaction is H2.
Stoichiometry Citations
- “Why not stoichiometry” versus “Stoichiometry–why not?” Part II: GATES in context with redox systems. Crit Rev Anal Chem . 2015;45(3):241-69.
- Stoichiometry and physical chemistry of promiscuous aggregate-based inhibitors. J Am Chem Soc . 2008 Jul 23;130(29):9606-12.
- Determining the stoichiometry and interactions of macromolecular assemblies from mass spectrometry. Nat Protoc . 2007;2(3):715-26.
- Defining the stoichiometry and cargo load of viral and bacterial nanoparticles by Orbitrap mass spectrometry. J Am Chem Soc . 2014 May 21;136(20):7295-9.
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What is Atomic Mass? How to Find...
Continue ReadingWhat is Atomic Mass?
A property that is thoroughly related to an atom’s mass number is its atomic mass. T
he atomic mass of a single atom is basically defined as its total mass and is naturally expressed in its atomic mass units or is also called as amu.
In other words, an atom of carbon has six neutrons, carbon-12, with an atomic mass of 12 amu.
In general, however, an atom’s atomic mass is very close to its given mass number but will have some deviation in its decimal places.
Since an element’s isotopes have different atomic masses, researchers may also govern the relative atomic mass which is sometimes called as the atomic weight of an element.
The relative atomic mass is defined as an average of the atomic masses of all the different isotopes present in a sample.
Celestial objects such as asteroids or meteors may have a very different isotope abundance.
How to Find Atomic Mass?
The figure given below explains how to find atomic mass and atomic number from a given element;
Atomic Mass Definition
Atomic mass is defined as the total mass of an atom. For example; the mass of 12C (carbon twelve) is inscribed as 12 amu.
All the masses of the elements are determined with 12C.
Periodic Table With Atomic Mass
The figure (periodic table with atomic mass) below shows the atomic mass of some of the frequently used elements;
What is Atomic Number?
The atomic number of an element describes the element’s identity and indicates the number of protons present in the nucleus of that particular atom.
For instance, the element hydrogen always has one proton present in its nucleus. The element helium always has two protons present in its nucleus.
Similarly, from the above figure element molybdenum have 42 protons present in its nucleus.
Atomic Mass and Isotopes
Atoms of the same element can, have different numbers of neutrons present in their nucleus.
For any assumed isotope, the summation of the numbers of protons and neutrons in the nucleus is known as the mass number.
This is because each proton and each neutron present in the nucleus weigh one atomic mass unit (amu).
By totaling together the number of protons and neutrons and multiplying it by 1 amu, we can calculate the mass of the given atom.
The word ‘isotope’ originates from the Greek word ‘isos’ which means ‘same’ and ‘topes’ means ‘place’.
Thus the name isotopes to given to elements having the same number of protons because these elements can occupy a similar place on the periodic table while being different in their subatomic construction.
For instance, Carbon exists as two main isotopes, 12C, and 14C carbon atom which has the same number of protons and electrons, 12C has 6 neutrons, and 14C has 8 neutrons and is a radioactive isotope.
Isotopes Types
• Few isotopes do have unstable atomic nuclei that undergo radioactive decay. Particularly these isotopes are radioactive and are, consequently, known as radioisotopes or radionuclides.
Examples for the same include carbon-14, chlorine-36, uranium-235, and uranium-238, etc.
• Some isotopes are identified to have tremendously long half-lives (hundreds of millions of years of life). Such isotopes are generally stated as stable nuclides or stable isotopes.
Some examples for the same include carbon-12, carbon-13, oxygen-17, and oxygen-18, etc.
• Primordial nuclides are defined as the nuclides that have existed since the formation of this solar system.
Out of 339 naturally occurring isotopes present on the Earth, a whole of 286 isotopes are recognized to be primordial isotopes.
Why Atomic Mass is Useful?
The atomic mass is beneficial in chemistry when it is connected with the mole concept.
The atomic mass of an element that is measured in amu, is similar as the mass present in grams of one mole of an element.
For instance, the atomic mass of iron is 55.847 amu and one mole of iron atoms also weigh 55.847 grams.
Atomic Mass Citations
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What is Law of Conservation of Mass?...
Continue ReadingHistory of Law of Conservation of Mass
The ancient Greeks initially proposed the idea that the total sum of matter in the universe is always constant. Though, Antoine Lavoisier stated the law of conservation of mass as a vital principle of physics in 1789.
What is Law of Conservation of Mass?
The Law of Conservation of Mass states that matter can neither be created nor it can be destroyed in a given chemical reaction. Thus, the total mass of reactants is always equal to the total mass of products.
This law of conservation of mass also stated that the mass of the system cannot change over a period of time in a closed system.
Law of Conservation of Mass example; the burning of a candle, a candle is made up of wax, thus When the lit candle, the oxygen gas present in the room will react with the candle wax to produce water vapor and carbon dioxide gas.
Thus the mass of reactants which are oxygen and wax will be equal to the mass of the products which are water and carbon dioxide.
The matter changes from one form to other but the mass of the matter always remains the same before and after the change.
For instance, if 44 grams of reactants go into a given chemical reaction, then 44 grams of products are being produced.
Law of Conservation of Mass Importance
The Discovery of this law of conservation of mass assisted to turn chemistry into the reputable science it is today.
With the initiation of this law of conservation of mass, chemists brought certainty and consistency to the science of chemistry.
If scientists know the quantities and characteristics of reactants for a specific reaction, then they can foresee the amounts of products that will be formed.
Law of Conservation of Mass Examples
I. Combustion Process
The burning of wood is an example of conservation of mass as the burning of wood comprises of production of gases such as Oxygen, Carbon dioxide, and water vapor, ashes.
II. Chemical Reactions
Water, for instance, comprises of two hydrogen atoms and one oxygen atom. Water is the only known substance on Earth that occurs naturally in three states: solid, liquid, and gas.
For water to change amongst these states, it must undergo physical changes. Thus, when water freezes, it changes into this hard and less dense substance known as ice, but the number of water molecules present before and after the change remains chemically the same.
Water gives a very vibrant example of how the cycle of matter works through our world, often changing form but never vanishing.
Both hydrogen and oxygen are diatomic atoms means they occur naturally as bonded pairs (H2 and O2, correspondingly). In the precise conditions and with an adequate amount of energy, these diatomic bonds will break and the atoms present individually will combine to form H2O (water).
The chemical equation can be written as:
2H2 + O2 = 2H2O
Note: There is the same number of hydrogen atoms and oxygen atoms on both sides of the equation that is reactants are always equal to the products.
III. Photosynthesis
Another Example of the same include Photosynthesis. In this process of photosynthesis, the plants convert light energy from the sun into chemical energy and store it in the form of sugars.
Although, light energy simply provides some energy for a chemical change to occur. The atoms required for photosynthesis are derived from carbon dioxide present in the air and water in the soil.
Then rearranges the atoms and the molecules into glucose C6H12O6 and oxygen O2.
6CO2 + 6H2O + sunlight = C6H12O6 (sugar or carbohydrates) + 6O2
This equation states that the six-carbon dioxide molecules combine with six water molecules thus forming one sugar or carbohydrate molecule and six molecules of oxygen.
By totaling or adding all the carbon, hydrogen, and oxygen atoms on both sides of the equation, the number of atoms or molecules would be equal.
Thus, the matter is preserved in this chemical change as well.
Law of Conservation of Mass Citations
- Scaling and scale invariance of conservation laws in Reynolds transport theorem framework. Ann Sci . 2015 Apr;72(2):153-69.
- The conservation of matter. Ann Emerg Med . 2002 Jan;39(1):91-3.
- The application of mass and energy conservation laws in physiologically structured population models of heterotrophic organisms. J Theor Biol . 1999 Apr 7;197(3):371-92.
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Dalton Atomic Theory: Definition, Examples, and Types
Continue ReadingWhat is Dalton Atomic Theory?
Dalton atomic theory is based on two of the well-known laws:
• The law of conservation of mass
• Law of constant composition.
The law of conservation of mass was formulated by Antoine Lavoisier states that matter can neither be created nor it can be destroyed in a closed system.
This means in a chemical reaction, the amount of each element involved in the given reaction must be the same as in the starting materials (reactants) and the products. This law of conservation of mass is used every time to balance chemical equations.
While researching about the same, Dalton also discovered that some gases having the same common element in given compounds could only be combined in certain proportions. Thus, he formulated the law of constant composition.
Postulates of Dalton Atomic Theory
The postulates of Dalton Atomic Theory are mentioned below;
• The matter is composed of atoms, and these atoms are the undividable building blocks of matter and thus cannot be destroyed.
Dalton theorized that the law of conservation of mass and the law of definite proportions could be described using the knowledge of atoms.
He proposed that all matter is made of tiny undividable particles which are known as atoms, and he assumed them as solid, hard, impassable, movable particles.
• The properties of all the atoms of an assumed element have the same mass. This can also be defined as the atoms of a given element have identical mass whereas the atoms of different elements have different masses.
Dalton projected that every single atom of an element, for instance, gold, is as same as every other atom of that particular element. He also stated that the atoms of one particular element vary from the atoms of all other elements.
For example; a sodium atom is very different from a given carbon atom. Elements may share few similar properties such as boiling points, melting points, etc. but no two elements can have the same set of exact properties.
• The Compounds are formed through different whole-number combinations of given atoms. In the next part of Dalton atomic theory, he theorized that compounds are the groupings of two or more different types of atoms.
For instance; the compound is table salt or NaCl. Table salt is a mixture of two separate elements the first is sodium which is a highly reactive metal and the second is chlorine which is a toxic gas.
When they react, the given atoms combine in a ratio of 1:1 to form white crystals thus forming table salt.
• A chemical reaction always results in the rearrangement of atoms of a given compound in the reactant and product formed.
In this part of Dalton’s atomic theory, he suggested that chemical reactions don’t terminate or create atoms. They just reorganize the atoms of compounds.
• Atoms of the same element usually combine to form two or more compounds in more than one ratio.
• The atom is thus the tiniest unit of matter and takes part in a different chemical reaction.
Advantages of Dalton Atomic Theory
• The three given laws were not exploited by Dalton; The law of multiple proportions, the law of conservation of mass, and the law of constant proportions
• The Dalton Atomic Theory also provides a basis to distinguish between elements and compounds.
Limitations of Dalton Atomic Theory
• Dalton atomic theory does not account for subatomic particles. Dalton atomic theory proposed that the atoms were undividable. Though, the discovery of subatomic particles (for instance, protons, electrons, and neutrons) refuted this postulate. Of Dalton.
• This theory fails to explain the existence of isotopes: Dalton atomic theory stated that all atoms present in the given element have identical masses and densities. Though, this postulate was disapproved as different isotopes of elements have different atomic masses
For instance: isotopes of hydrogen are; hydrogen, deuterium, and tritium.
• Dalton atomic theory fails to explain isobars also. This theory of Dalton states that the masses of the given atoms of two different elements must vary. Thus, two different elements can have the same mass number. Such atoms are known as isobars.
For Instance: 40Ar and 40Ca.
• This theory of Dalton does not account for allotropes as well.
For instance; the differences in the properties of diamond and graphite, containing a single atom of carbon, cannot be described by Dalton’s atomic theory.
In spite of of these drawbacks, Dalton atomic theory is generally true, and it forms the basis of modern chemistry.
Dalton Atomic Theory Citations
- John Dalton and the origin of the atomic theory: reassessing the influence of Bryan Higgins. Br J Hist Sci . 2017 Dec;50(4):657-676.
- A letter signed: the very beginnings of Dalton’s atomic theory. Ambix . 2010 Nov;57(3):301-10.
- Atomic Theory and Multiple Combining Proportions: The Search for Whole Number Ratios. Ann Sci . 2015 Apr;72(2):153-69.
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Law of Definite Proportions: Definition, Examples, and...
Continue ReadingWhat is Law Of Definite Proportions?
This law of definite proportions is the basis for the study of stoichiometry in Chemistry.
The law of constant proportions states that chemical compounds are made up of elements that are present in a fixed ratio in terms of their mass.
This specifies that any pure sample of a compound, (no matter what is the source of that compound), will always encompass the same elements that are present in a similar ratio by their mass.
The Law of Definite Proportions also shows that whatever the quantity of water is, whether it be 4 moles or 72 grams, the ratio of the amount of hydrogen to oxygen by its weight will always remain the same, just like that of egg:sugar: butter ratio in the chocolate cake is always constant to maintain the delicious taste of the cake.
Law of Definite Proportions Examples
• For instance, in a nitrogen dioxide (NO2) molecule, the ratio of the number of nitrogen atoms present in given compound and oxygen atoms is 1:2 always. This ratio between the nitrogen and oxygen molecules would always remain the same.
• Salt which is written as the chemical compound NaCl comprises of atoms of Sodium (Na) and Chlorine (Cl). Thus, to create salt the exact same proportions of sodium and chlorine must always be combined.
• Sulfuric acid comprises of the specific elements of hydrogen, sulfur, and oxygen. The chemical formula of sulfuric acid is written as H₂SO₄. Thus, to create sulfuric acid the same proportions of hydrogen, sulfur, and oxygen must be combined.
• Ammonia which is a common household item is made up of the elements of hydrogen and nitrogen. The chemical formula for Ammonia is written as NH3, meaning that there is one atom of nitrogen which is combined with 3 atoms of hydrogen. Anhydrous ammonia consists of 82% nitrogen and 18% hydrogen. Any other grouping of hydrogen and nitrogen would result in a completely different type of chemical compound.
Law of Definite Proportions Applications
The law of definite composition has applications of both the compounds;
• molecular compounds having a fixed composition
• whereas ionic compounds require certain ratios to achieve electrical neutrality
• There are certain exceptions to the law of definite composition. These compounds are thus known as non-stoichiometric compounds, and examples of the same include ferrous oxide
• Further, this law of definite composition does not justify for isotopic mixtures.
Law of Definite Proportions and Non-Stoichiometric Compounds / Isotopes
There exist certain non-stoichiometric compounds whose elemental composition can vary from one to another. Such kinds of compounds obey the Law of multiple proportions.
For example, the iron oxide, which may hold iron atoms between 0.83 and 0.95 for each of the oxygen atoms, and consequently comprise anywhere between a percentage of 23 and 25 oxygen by mass.
The ideal chemical formula is given as FeO, but because of the crystallographic vacancies, it is up to FeO.95O.
Basically, Proust’s measurements were not precise enough to spot such small differences. The element’s isotopic composition can also differ based on its source; therefore, its involvement in even a pure stoichiometric compound mass can vary.
This discrepancy can be used in radiometric dating since atmospheric, astronomical, crustal, oceanic, and deep Earth processes can concentrate limited environmental isotopes favorably.
With the hydrogen and its isotope exception, generally, the effect is slight, but it is assessable with modern-day instrumentation.
Law of Definite Proportions Citations
- The Law of Definite Proportions: Nature volume 84, page364 (1910)
- The Berthollet-Proust Controversy and the Law of Definite Proportions. Nature volume 50, pages149–150 (1894)
- How Important are the Laws of Definite and Multiple Proportions in Chemistry and Teaching Chemistry? – A History and Philosophy of Science Perspective. Science & Education volume 10, pages243–266 (2001)
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Law of Chemical Combination: Definition, Types, and...
Continue ReadingWhat is Law of Chemical Combination?
Chemistry is the study of the change of a substance from one form to the other.
These transformations frequently occur as a result of the combination of two diverse types of matter.
The combination of various elements to form compounds is ruled by certain basic rules. These rules are stated as laws of chemical combination.
The law of chemical combination describes the basic principles that are followed when atoms and molecules interact.
This remarkable variety of interactions allows for a diverse variety of chemical reactions and compounds.
Though chemical reactions are complex yet they are all primarily ruled by some guiding laws of chemical combination, which lay the basis for the investigation or analysis of chemical reactions.
Law of Chemical Combination Rule
The combination of elements to form compounds is ruled by the five basic laws which are mentioned below:-
1. Law of Conservation of Mass.
2. Law of Definite Proportions.
3. Law of Multiple Proportions.
4. Gay Lussac’s Law of Gaseous Volumes.
5. Avogadro Law.
1. Law of Conservation of Mass
In 1789, this law of conservation of mass was framed by Antoine Lavoisier. In simple terms, this law of conservation of mass states that matter can neither be created nor it can be destroyed.
In other words, the total mass, that is, the sum of the mass of reacting mixture that is the reactants and the products formed always remains constant.
For instance; the Burning of wood Thus, when wood burns the mass of the dust, ashes, and fumes is always equal to the original mass of the charcoal and the oxygen present when it initially reacted.
So the mass of the product which is ashes and gases equals the mass of the reactant which is charcoal and oxygen.
2. Law of Definite Proportions
A French well-known chemist, Joseph Proust stated that the proportion of elements by mass in a given compound will remain the same at all times.
In simple terms, This law of constant proportions states that chemical compounds are made up of elements that are present in a fixed ratio in terms of their masses.
For example; A nitrogen dioxide (NO2) molecule is taken, thus, the ratio of the number of nitrogen atoms and oxygen atoms present in a given compound is 1:2 always. This ratio of 1:2 between the molecules of nitrogen and oxygen would always remain the same.
3. Law of Multiple Proportions
In 1803, This law of multiple proportions which was given by Dalton stated that if two elements combine to form more than one compound, the masses of these elements present in the reaction are defined in the ratio of small whole numbers.
For instance, the masses of molecules of oxygen is combined with a fixed mass of carbon in CO2 (carbon dioxide ) and CO (carbon monoxide) are 32 and 16, respectively. Consequently, these masses of oxygen display a simple ratio of 32: 16 or it can be written as 2: 1.
4. Gay Lussac Law of Gaseous Volumes
In 1808, Gay Lussac gave this law built on his observations. In simple words, Gay Lussac law or Amonton’s law states that the pressure exerted by a gas is directly proportional to the temperature of the gas when the mass is fixed and the volume is kept constant.
The mathematical expression of the above-given law can be inscribed as follows:
P ∝ T;
P = kT
Thus,
P/T = k
• P is as defined the pressure applied by the gas on the walls of its vessel
• T is defined as the absolute temperature of the gas
For example, when aerosol cans are kept under warm conditions they eventually burst, because by heating the aerosol can pressure of the contents increases ultimately causing the can to burst.
When a pressurized aerosol can for example a deodorant can or a spray-paint can is heated, the consequent increase in the pressure is applied by the gases on the walls of its vessel can result in a blast.
This is the reason why numerous pressurized bottles or vessels have alerting labels stating that the given vessel must be kept away from fire and should be stored in a cool atmosphere.
5. Avogadro’s Law
Avogadro anticipated this law in the year 1811. Avogadro’s law is a relationship between the volume of gas and the number of moles.
This law states that at a constant temperature and pressure the total number of atoms or molecules present in a gas is directly proportional to the volume employed by that gas.
This indicates that 4 litres of hydrogen will have the same number of molecules as 4 litres of oxygen assumed that both the gases are at the same temperature and pressure.
This equation is mentioned below;
V ∝ n
V= kn
or
k = V/n
V is defined as the volume of gas
n is defined as the number of moles present in a given gas
k is defined as the proportionality constant
For instance, A compressed pool tube becomes portable when the number of air particles inside the tube is decreased which in turn decreases its volume and makes it compact.
During inflation, when the tube is filled with air thus it increases the number of air molecules in it which in turn increases the volume and size of the pool tube.
Hereafter, Avogadro’s law can be useful to inflate or deflate the given pool tube as per our need.
Law of Chemical Combination Citations
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Gay Lussac Law: Formula, Equation, Calculation, and...
Continue ReadingWhat is Gay Lussac Law?
The law is named after a French chemist and physicist named Joseph Gay-Lussac. He articulated this law in 1808.
In simple words, Gay Lussac law or Amonton’s law states that the pressure exerted by a gas is directly proportional to the temperature of the gas when the mass is fixed and the volume is kept constant.
Gay Lussac Law Equation
The mathematical expression of the given law can be written as follows:
P ∝ T; P=kT
Thus, P/T = k
• P is defined as the pressure exerted by the gas on the walls of a container
• T is defined as the absolute temperature of the gas
• k is defined as a proportionality constant.
Gay Lussac Law Examples
The heating gas in a sealed vessel causes its pressure to rise, whereas cooling a gas lowers or reduces its pressure.
The happens because the increasing temperature imparts thermal kinetic energy to the given gas molecules.
As a result when the temperature increases, molecules present in the gas collide more often with the walls of the vessel.
These increased collisions are perceived as increased pressure.
Gay Lussac Law Graph
The association between the pressure and absolute temperature of a given mass of gas present at a constant volume can be demonstrated graphically as shown below;
From the above graph, it can be seen that the pressure of a gas that is taken at constant volume decreases with temperature as this pressure is directly proportional to the temperature.
Gay Lussac Law Formula and Derivation
Gay Lussac law indicates that the ratio of the initial pressure and temperature is always equivalent to the ratio of the final pressure and temperature for molecules of gas having fixed mass and is kept at a constant volume.
This formula can be written as given below:
(P1/T1) = (P2/T2)
• P1 is defined as the initial pressure
• T1 is defined as the initial temperature
• P2 is defined as the final pressure
• T2 is defined as the final temperature
This expression is formed from the pressure-temperature proportionality law of gas molecules.
Since,
P ∝ T for gases,
Thus,
P1/T1 = k
(initial pressure/ initial temperature = proportionality constant)
P2/T2 = k
(final pressure divided by final temperature = proportionality constant)
Consequently,
P1/T1 = P2/T2 = k
Or this can be written as,
P1T2 = P2T1
Gay Lussac Law Examples
Here are few examples of Gay Lussac law observed in our everyday life:
I. The Pressure
Automobile tire pressure drops on a cold day and climbs on a hot day. After driving, the air pressure present in a car’s tires goes up. This happens because the frictional force present between the tires and the road causes the air which is present inside the tires to heat up.
This air present in tires cannot expand further because the tires have a fixed-volume container, Thus, the pressure increases. The above example defines Gay Lussac law clearly.
II. Pressure Cooker
When heat is applied to a pressure cooker it increases the pressure inside the cooker. Thus, the Increasing pressure leads to an increase in the boiling point of water, as a result, it shortens the cooking time.
As the temperature of the liquid water rise, thus leading to the formation of water vapour (that is water present in its gas state). This vapour cannot escape the pressure cooker as it is a closed container, which means the volume is constant.
At the higher temperature which is the normal boiling point of water (100°C) food can be cooked sooner
III. Aerosol Can
It is generally advised not to store the aerosol cans under warm conditions or because by heating the aerosol can pressure of the contents increases eventually causing the can to burst.
When a pressurized aerosol can for instance a deodorant can or a spray-paint can is heated, the subsequent increase in the pressure applied by the gases on the walls of the container can result in an explosion.
This is the reason why several pressurized bottles or containers have cautioning labels stating that the purchased container must be kept away from fire and should be stored in a cool environment.
IV. Water Heater
An electric water heater is a lot similar to a pressure cooker. A pressure-relief valve present in the water heater prevents the steam formed from accumulating.
If the valve present in the heater fails to function due to one reason or the other then the heat drives up the steam pressure inside the heater, ultimately bursting it.
Gay Lussac Law Limitations
The limitations of Gay Lussac law are as follows:
The Gay Lussac law is only valid to ideal gases.
Gay Lussac law holds moral for real gases either at high temperatures or low pressure.
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Law of Multiple Proportions: Example, Definition, and...
Continue ReadingWhat is Law of Multiple Proportions?
o The law of multiple proportion is also known as Dalton’s law which was proposed by English chemist and metrologist named John Dalton in 1804.
o This law was based on the observation of Dalton which states that when the elements combine to form two or more compounds the proportion of the elements in those chemical compounds can be articulated in this small whole number ratios.
o This is law of multiple proportions is also a part of the basis of modern atomic theory.
o This law of multiple proportion only applies to compounds which are composed of same elements.
o When SO2 and CO2 are taken together then one compound has sulfur (S) and one has carbon (C), thus, here the law of multiple proportions is not valid as compounds composed of only same elements is applicable to law of multiple proportions.
Postulates of Dalton's Atomic Theory
o All matter consists of tiny particles named atoms.
o Atoms cannot be created nor it can be destroyed.
o The atom is the smallest unit present in matter and take part in chemical reaction.
o Atoms of the same element are identical in their shape and mass.
o Atoms of different elements when combined in a constant ratio form compound.
o Atoms of the same element can also combine in more than one ratio to form different compounds
Law of Multiple Proportions Examples
o Assume two molecules of CO that is carbon monoxide and CO2 that is carbon dioxide
o CO contains 12 grams of carbon and 16 grams of oxygen
o CO2 contains 12 gram of carbon plus 32 grams of oxygen
o The ratio of mass of oxygen in each of the two compounds is expressed in the ratio 16 :32 which is equal to 1: 2
o Therefore, the law of multiple proportions is proved here.
o Another example of law of multiple proportions includes; Hydrogen and oxygen combine to form two compounds of H2O and one compound of hydrogen peroxide.
o Thus, the ratio of different weights of oxygen is 16 and 32 combining with a fixed weight of two hydrogen thus,16 I; 32 that is 1; 2 which is a simple whole number ratio.
o Therefore, Law of multiple proportions is proved here.
o Nitrogen reacts with oxygen to form numerous nitrogen oxides under various reaction conditions.
o Some of them include; nitrogen monoxide (NO), nitrogen dioxide (NO2), dinitrogen trioxide (N2O3), and dinitrogen pentoxide (N2O5).
o The amount of nitrogen taken is 14 g.
o When take the ratio of the amounts of oxygen that is required each reaction, thus the ratios obtained are 8:16:24:32:40 or 1:2:3:4:5.
o Thus, the ratio obtained is in small whole numbers and the given law of multiple proportions is also valid for more than two reactions.
Law of Multiple Proportions Limitations
o The existence of different isotopes of hydrogen like H1 or H2 causes differences. Thus, the same isotope should be used throughout the reaction for preparation of compounds.
o This law of multiple proportions fails often when heavy weight molecules are involved in the given reaction. Example for the same include biochemical reaction.
o Atomic masses of elements are not always in the whole number example the atomic mass of nitrogen. presumed to be 14 amu, but its value is 14.0067 amu. When the ratio of such number is taken, the results obtained are very large whole numbers.
o The law is valid only for normal reactions. It fails repeatedly when a non-stoichiometric compound is involved in the given reaction.
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What is Density? Formula, Calculation, Definition, and...
Continue ReadingWhat is Density?
o The principle of density was discovered by a Greek scientist named Archimedes.
o Density is a physical property that is defined as mass divided by its volume.
o Everything around us has mass, it is a fundamental property and plays a role in two other important properties which are volume and density.
o The SI unit of mass is kilograms or grams and it is a measure of the amount of matter present in an object.
o Volume is defined as a three-dimensional space that an object occupies it is expressed in cubic metre.
o Thus, Density is the ratio of mass to volume and the SI unit of density is expressed as kg per metre cube (kg/m3) or per gram per centimeter cube(g/cm3).
o By definition, density is the measurement of how closely or loosely a matter is packed into a certain volume.
Everyday Examples of Density
o Wood floats on water because it has less density than that of water.
o Helium balloons rise in the air because helium gas has less density as compared to that of surrounding air.
o Ice floating on water is also an example of density which can be attributed to the Archimedes principle.
o Oil spills- when an oil tanker leaks on ocean/water body the oil floats on the water since oil has less density as compared to that of water and this helps to clean up the oil spills from surface of the earth.
Factors Affecting Density
o Density differs according to temperature and pressure.
o Temperature: As the temperature rises, most substances enlarge or increase their volume.
o This results in the decrease in density.
o Similarly, when the temperature goes down, the density of the substance increases.
o Thus, density is inversely proportional to temperature.
o Pressure: As pressure rises, density increases.
o Similarly, when pressure decreases density decreases because the inter molecular force of attraction decreases.
o Thus, density is directly proportional pressure.
o To sum up the above statements, density is inversely proportional to temperature and directly proportional to pressure exerted on a substance.
Change of Phase
o Liquids are less dense as compared to solids because solids have densely packed particles whereas particles of liquid that can slide around.
o Example; Density of water at its freezing point is 0.999 g/cm3, whereas density of ice is 0.92 g/cm3.
o Gases are less dense as compared to liquids because the particles of are free and can move around.
o When a liquid changes into a gas, its volume rises vividly, thus reducing the density of the substance.
o Example; Water has a density of 1.03 g/cm3 at 100°C, whereas the density
Change in Volume
o The volume of a substance changes with change in its temperature and pressure. Thus, in turn, it changes the density of the substance.
Density of Water
o The density of pure water is taken as 1 g/cm3.
o The water density depends on the purity of water.
o As pure water is less dense as compared to the saline water.
o When the purity of water molecules decreases, the density of water increases.
o Contamination in the water interrupts the density of water.
o Due to contamination, the density of water rises and this change in density depends on the degree of contamination and the temperature.
o Water is also identified as the universal solvent and it is the key component of fluids found in all organisms.
o Water is profusely available on the Earth and is found in all forms that is gas, liquid and solid.
o When the temperature reduces, water changes into solid and when the temperature rises, water changes into gas.
o At the room temperature, water is present in the liquid state.
Why Density is Important?
o Density is significant because it defines whether the given substance will rise or sink.
o So, understanding density has main inferences for the motions of substances and gases in the atmosphere and materials floating or sinking in water (H2O).
o By determining of the densities of substances, help us in their separation techniques.
o For instance, separation of oil from the water.
o Another significant application of density is determining whether a substance will float on or will sink in water.
o Example for the same include the floating of ships and diving of submarines which happens due to their density difference.
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