Category: Chemistry Notes PPT

In the physical definition of any matter, most of the physical properties are subdivided into intensive and extensive properties. The identity and function of any substance or system are defined by these properties. Let us have a look at these two sub-categories in detail.

What is an Intensive Property?

Intensive properties of any matter are those physical properties that are independent of the mass of the substance or the system. Intensive properties are also known as bulk properties. Most intensive properties are used to define the identity of that substance or system. 

Intensive Property Examples

Pressure (P), temperature (T), color are all intensive properties. Other examples include density, melting point, boiling point, etc. All these parameters do not change with the mass of the body. For example, the melting point of 1 kg ice and 1 gm of ice is the same= 0ᴼC.

Chemical potential, refractive index, specific heat capacity, thermal conductivity, viscosity are all examples of intensive properties.

What is an Extensive Property?

Extensive properties of any matter are those physical properties that depend on the mass of that matter. These properties are proportional to the size or mass of the system.

Extensive Property Examples

The weight of the system increases with the mass. Similarly, the volume also increases with the mass of the substances. The heat capacity is directly proportional to the mass of a system. The energy stored in a system is dependent on the mass of the system. For example, two boxes of the same material but the different weights will also differ in their properties.  

Some other examples of extensive properties are enthalpy, entropy, Gibb’s energy, internal energy, etc.

Differences between Extensive and Intensive Properties

How to Differentiate between Intensive and Extensive Properties?

It is easy to distinguish between intensive and extensive properties. One needs to double the mass of the system. The physical properties that change with an increase in mass are extensive properties. However, those physical properties that do not change with an increase in mass are intensive properties.

Other Examples of Properties

Both the intensive and extensive properties are useful in understanding the thermodynamics of a system. Thermodynamics is the study of the flow and transformation of heat forms of any matter. It depends on the matter and the factors determining its state. Parameters that define the thermodynamic properties are:

• Path function- the parameter defined by the path taken by the matter or the system to reach the current state. Work done due to frictional force is an example of path function.

• State functions, also known as state variables, are defined by the current state, and not the path that is taken to reach that state. Temperature is an example of state function. 

The state function of the system depends on the initial and final position of the system. However, the path functions of the system depend on the path taken by the system to reach from the initial to the final state. Both the state and path functions are important parameters to study the thermodynamic properties of a system.

Ionisation energy is the amount of energy required to remove an electron from a specific gaseous atom or ion. It applies to all the elements on the periodic table and not just the atoms that are gases at room temperature.

Trends and Periods

Looking at the periodic trend, as the students go from lithium over to neon, across the periodic table, the students can notice that there’s an increase in the ionisation energy. Lithium is positive 520 kilojoules per mole, and Beryllium’s goes up to 900 kilojoules per mole, and then again, in general, there’s an increase in ionisation energies going over to neon. That is because there’s also a relative increase in the effective nuclear charge. An ion is just an atom or a molecule with a charge, and it’ll have a charge if the protons are not equal to electrons. Neutrons are also composed of atoms but are neutral. The charge is given from protons or electrons, which is a net charge for an atom or molecule. A molecule’s just a cluster of atoms bonded together. The negative ions are more significant in the number of electrons than protons. So, for example, Hydrogen in its neutral state has one proton and one electron. Still, even if one of the electrons is taken away, then Hydrogen would have a positive charge, and essentially, it would just be, in its most common isotope, it would just be a proton by itself. And so, when it’s a positive ion where the number of protons is more than electrons, it is called Cations. Cation is just another word for positive ions. Likewise, we can have negative ions. For example, Fluorine. When fluorine gains an electron, it will have a negative charge. A negative ion is named an Anion. 

With the help of Ionisation, one can ionise different elements in the periodic table and turn them into cations. However, turning the element into gas is necessary before moving onto the electron.

Metals have low ionisation energy, whereas nonmetals have high ionisation energy. Ionisation energy will increase from left to right, and it will rise from the bottom to the top on the periodic table. Therefore, the lowest ionisation energy will be Francium, and the highest ionisation energy will be Helium. 

Isotones definition is given as any of either two or more species of nuclei or atoms that contain a similar neutron count. Since the nucleus of this species of chlorine has 17 protons and 20 neutrons, while the nucleus of this species of potassium has 19 protons and 20 neutrons, potassium-39 and chlorine-37 are isotones.

About Isotone

Two nuclides are said to be isotones if they contain similar neutron number N, but with different proton numbers Z. For example, carbon-13, and boron-12 nuclei both have 7 neutrons, and so they are called isotones. In the same way, 37Cl, 36S, 39K, 38Ar, and the 40Ca nuclei are all isotones of 20 since they all hold 20 neutrons. About its similarity to the Greek word for “same stretching,” the word “isotope” was created by German physicist K. Guggenheimer by changing the letter “p” in “isotope” from “proton” to “neutron.”

Observationally, the largest numbers of the stable nuclides there exist for isotones 50 (which are five: 88Sr, 86Kr, 90Zr, 89Y, and 92Mo) and 82 (which are six: 139La, 138Ba, 141Pr, 140Ce, 144Sm, and 142Nd). The neutron numbers where there are no stable isotones are given as 19, 21, 35, 39, 45, 61, 89, and 115 or even more. In contrast, the proton numbers, where there are no stable isotopes, are given as 43, 61, and 83 or even more.

This is related to the nuclear magic numbers, which are the number of nucleons that form complete shells within the nucleus. For example, 2, 8, 28, 50, 82. Not more than one stable nuclide contains a similar odd neutron number, except for 1 (it means 2H and 3He), 5 (which means 9Be and 10B), 55 (the 97Mo and 99Ru), and also 107 (the 179Hf and 180mTa). 27 (50V), 65 (113Cd), 81 (138La), 85 (147Sm), and 105 are the odd neutron numbers of a primordial radionuclide and a stable nuclide (176Lu). Neutron numbers, where there exist two primordial radionuclides, are given as 88 (151Eu and 152Gd) and 112 (187Re and 190Pt).

About the discovery of Neutron, Isotopes, Isobars, and Isotones

Introduction

It’s a remarkable fact that the neutron existence was not discovered until 1932. Protons and electrons are the most general atomic imagination of the time. Through the alpha scattering experiments of Rutherford, it was found that the Atomic mass number, which is ‘A’ of an element, is a bit greater than twice the atomic number, which is ‘Z’ for the most atoms and that importantly all the mass of an atom was concentrated in a tiny space at the center of the atom. The alpha particles, which took a turn of 180-degree stand as proof of this.

Until 1930, some of the electrons were thought to coexist with protons in the dense nucleus, whereas the huge amount of energy needed to sustain such a type of system was way beyond atomic energies. If we take the Hydrogen atom size as 0.2 nanometers, then the electron confinement energy is given as 38eV, which is the exact magnitude for the atomic electrons. Whereas, if the electron were to coexist with protons present in the nucleus, the electron confinement energy is approximately given as 250Mev! Several magnitudes are higher than the 38eV.

What is Neutron?

James Chadwick has provided an answer to this puzzle, who boldly stated this was a new fundamental particle type, that is neutral, and he called them Neutrons. From the conservation of momentum and energy, he was able to derive with considerable accuracy this new particle’s mass. He also found that a neutron’s mass was closer to that of a proton.

MN = 1.00866 

U = 1.6749 x 10^-27 Kg

So, the nucleus had another resident now, and the proton-neutron pair was known as a Nucleon. The Neutron discovery led to a better understanding of atomic number and atomic mass also with the isotopes; that is what the radioactivity is based on.

N – Neutron Number = Number of Neutron

Z – Atomic Number = means, the number of protons or electrons

A – Atomic Mass Number = Z + N = it means, the total number of protons and neutrons

Hence, the elements of the periodic table now had a new form of representation;

For example, the Uranium atom nucleus can be represented by a 23592 U, which means one atom of Uranium 235 comprises 235 nucleons, where 92 are the protons, and the rest of 143 is said to be neutrons.

What are Isotopes?

Isotopes are the variants of a specific element with a different number of neutrons. For example, the two isotopes of Uranium are given as 23992 U and 23592 U. We can also notice that the number of protons is similar in both the isotopes, but they have 147 and 143 neutrons, respectively. An extra neutron presence significantly changes the behaviour of that specific atom. There exist two different types of isotopes, radioactive and stable. Radioactive isotopes are defined as the ones that are more unstable to sustain themselves, and they break down into two lighter daughter elements spontaneously with the particle emission, such as the rays of alpha, beta, gamma. Stable isotopes are the ones that can exist in their free state without spontaneously breaking down.

Kohlrausch law states that at infinite dilution, when dissociation is complete, each ion makes a definite contribution towards equivalent conductance of the electrolyte irrespective of the nature of the ion with which it is associated and the value of equivalent conductance at infinite dilution for any electrolyte is the sum of the contribution of its constituent ions (cations and anions). Thus, we can say it states that ‘conductivity of ions of an electrolyte at infinite dilution is constant and it does not depend on nature of co-ions.’

[lambda_{eq}^{infty} = lambda_{c}^{infty} + lambda_{a}^{infty}]

[lambda_{eq}^{infty}] = Molar conductivity at infinite dilution 

[lambda_{C}^{infty}]= Conductivity of cation at infinite dilution 

[ lambda_{a}^{infty}]= Conductivity of anion at infinite dilution 

When the concentration of the electrolyte is almost zero, at that point, molar conductivity is called limiting molar conductivity. 

The molar conductivity of the solution can be defined as the volume of the solution that is conducting that also contains one mole of electrolyte when kept between two electrodes with a unit area of cross-section and one unit length of distance. With the decrease in the concentration, the molar conductivity increases. The increase in the molar conductivity is due to the increase in the volume that comprises one mole of electrolytes. The molar conductivity is known as limiting molar conductivity, Ëm°, when the concentration of the electrolyte approaches zero. 

Explanation of Kohlrausch Law

Any random electrolyte is the general case of this law which can be denoted as [ A_{x}B_{y}]. 

Thus mathematically, the limiting molar conductivity of [ A_{x}B_{y}] can be represented as:

[ lambda _{AxBy}^{infty} = 2 lambda_{A + B}^{infty} + lambda_{B – x}^{infty}]

Where [lambda^{infty}] is the limiting molar conductivity of the electrolyte chosen. 

When a cation is the same in both the electrolytes, then the difference in the molar conductivity of the two electrolytes does not depend upon the cation and is only dependent on the change that happens in their anions. The statement mentioned is also true if the anions are the same and the cations are different.  

For instance, if there are two pairs of electrolytes with the same cation A and D in each pair, then the difference between their limiting molar conductivities is not affected by A or D. This can be mathematically represented as,

[ lambda _{AB}^{infty} – lambda _{AC}^{infty} = lambda _{DB}^{infty} – lambda _{DC}^{infty}]

Where

[ lambda _{AB}^{infty}] is limiting molar conductivity of AB,

[lambda _{AC}^{infty}] is limiting molar conductivity of AC,

[ lambda _{DB}^{infty}] is limiting molar conductivity of DB, and

 [ lambda _{DC}^{infty}] is limiting molar conductivity of DC.

Uses of Kohlrausch’s Law 

• Kohlrausch’s law is used to calculate molar conductivity at infinite dilution for the weak electrolytes. It’s very difficult or impossible to calculate the molar conductivity of weak electrolytes at infinite dilution. as the conductance of these types of solutions is very low and dissociation of these electrolytes is not completed at high dilutions as well. For example, acetic acid is a weak electrolyte and its molar conductivity at infinite dilution can be calculated by Kohlrausch’s Law. It can be represented as follows: 

[mu^{infty}] = Molar conductance at infinite dilution

[mu^{infty} = mu^{infty} = m lambda _{+}^{infty} + m lambda _{-}^{infty}]

• m and n are a number of ions formed. 

• For example – molar conductance of aluminium sulphate at infinite dilution can be written as follows – 

• The formula of aluminium sulphate is [Al_{2}(SO_{4})_{3}.]

• So, molar conductance at infinite dilution = [ mu _{Al_{2}}^{infty } (SO_{4})_{8} = 2 lambda _{Al^{s}}^{infty } + 3lambda _{SO{_{4}}^{2-}}^{infty }] 

• The knowledge of molar conductivities at infinite dilution of the strong electrolyte like HCl, [ (CH_{3}COONa)], and NaCl and the molar conductivity of acetic acid at infinite dilution can be obtained as follows:

[ lambda_{m(HCl)}^{infty } = lambda _{H^{+}}^{infty }  + lambda _{Cl^{-}}^{infty } ]

[lambda_{m(NaCl)}^{infty } = lambda _{Na^{+}}^{infty }  + lambda _{Cl^{-}}^{infty }]

[lambda_{m(CH_{3}COONa)}^{infty } = lambda _{CH_{3}Coo^{-}}^{infty }  + lambda _{Na^{+}}^{infty }]

[lambda_{m(CH_{3}COOH)}^{infty } = lambda _{CH_{3}Coo^{-}}^{infty }  + lambda _{H^{+}}^{infty }]

[lambda_{m(CH_{3}COOH)}^{infty } = [lambda _{CH_{3}Coo^{-}}^{infty } + lambda _{Na^{+}}^{infty }]  – [ lambda _{Na^{+}}^{infty } + lambda _{Cl^{-}}^{infty }] [ lambda _{H^{+}}^{infty } + lambda _{Cl^{-}}^{infty }]

[ alpha = frac{lambda _{m}}{ lambda _{m}^{o}} ]

[ K_{a} = frac{Lambda _{m}^{2} .C}{Lambda _{m}^{2}left [ 1 – frac{Lambda _{m}}{Lambda _{m}^{o}} right ] } ]

[ Rightarrow K_{a} = frac{Lambda _{m}^{2}.C}{Lambda _{m}^{o2}left ( frac{Lambda _{m}^{o} – Lambda _{m}}{Lambda _{m}} right )}]

[ Rightarrow K_{a} = frac{Lambda _{m}^{2}C}{Lambda _{m}^{o}(Lambda _{m}^{o} – Lambda _{m})} ]

[ K_{c} = frac {Ca ^{2}} {1 – alpha} ]

Where

K = dissociation constant,

C = concentration of the solution, and 

α = degree of dissociation. 

• Kohlrausch’s law is used for the calculation of solubility of moderately soluble salt. Some salts that dissolve in very small quantities in water are called moderately or sparingly soluble salts. For example – silver chloride, barium sulphate, lead sulphate, etc. 

• Acid dissociation constant [K_{a}] can also be calculated by Kohlrausch’s law. 

• When the concentration of the electrolyte is almost zero, at that point, molar conductivity is called limiting molar conductivity. By Kohlrausch’s law, we can determine limiting molar conductivity for an electrolyte. 

Chemistry is the study of matter, the composition of matter and its different forms. Matter changes into different forms by chemical combinations. These chemical combinations of different elements and compounds follow a set of laws. This is the reason why we always balance chemical equations. 

Here in this article, we are providing you with five basic laws of chemical combination which rule the chemical combinations of elements.

Basic Laws of Chemical Combinations

The five basic laws of chemical combination for elements and compounds are given below.

1. Law of Conservation of Mass 

The Law of conservation of mass states that “Mass in an isolated system can neither be created nor be destroyed but can be transformed from one form to another”. It was given by French Chemist Antoine Lavoisier in 1789. Mass of reactants and mass of products will always be equal in a chemical reaction. According to this law, mass can neither be created nor destroyed. This is the reason why we always balance a chemical equation. 

Thus, for any chemical reaction or chemical change, the mass of the reactants is equal to the mass of the product formed. This law can be explained by Dalton’s atomic theory. According to Dalton’s atomic theory –Atoms are indivisible particles, which cannot be created or destroyed in a chemical reaction. 

Example – Formation of water

2H2 + O2 🡪 2H2O 

Mass of reactants = 4 + 32 = 36g

Mass of products = 2(2+16) = 2(18) = 36g

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As we can see the mass of reactants and mass of product is equal in the reaction. So, it proves the law of conservation of mass.  

2. Law of Definite Proportions 

The law of definite proportions was proposed by Joseph Proust in 1799. It is also known as the law of constant proportions. According to this law in a chemical substance, the elements are always present in definite proportions by mass. John Dalton’s theory provided an explanation for the law of constant proportions as well. According to John Dalton’s theory, the relative number and kinds of atoms are constant in a given compound. This statement supports the law of constant proportions. 

Example – In a water molecule, the ratio of the mass of hydrogen and mass of oxygen always remains the same which is 1:8. You can take water molecules from any source but hydrogen and oxygen will always remain in the 1:8 ratio by mass in a water molecule. 

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3. Law of Multiple proportions 

Law of multiple proportions was given by John Dalton in 1804. According to this law if elements combine to form two or more than two different kinds of compounds, then the masses of these elements in the compounds are in the ratio of small whole numbers. Dalton’s atomic theory states that atoms combine in the ratio of small whole numbers to form compounds. 

Example – Carbon forms two oxides with oxygen- Carbon monoxide (CO) and carbon dioxide (CO2). In these compounds (one molecule) mass of carbon is 12g (fixed) and the ratio of masses of oxygen in both compounds CO and CO2 are 16:32 or 1:2. 

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4. Gay Lussac’s Law of Gaseous Volumes 

Law of Gaseous Volumes was proposed by Joseph Louis Gay-Lussac in 1808. According to this law when measured at the same temperature and pressure, the ratio of the volumes of reacting gases are small whole numbers. This can be considered as a different form of the law of definite proportions. This law is with respect to volume while the law of definite proportion is with respect to mass. 

Example – 

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5. Avogadro’s Law 

Avogadro’s law was given by Amedeo Avogadro in 1811. According to this law, equal volumes at the same temperature and pressure contain equal numbers of moles of gases. It means that 2 litres of oxygen and 2 litres of nitrogen will contain the same numbers of moles if measured at the same temperature and pressure. 

Conclusion

Above, we have discussed the law of chemical combination in a simple manner, if still you have doubts and want more detailed notes on the topic then register yourself on and get access to free PDFs of NCERT Solutions, study material etc.

In a chemical reaction, the limiting reagent is the reactant that determines the quantity of the products that are produced. The other reactants present in the reactions are sometimes found to be in excess since there is some leftover quantity of them after the limiting reagent is completely used up. The maximum amount of product that is produced is known as the theoretical yield. The limiting reagent should be identified to calculate the percentage yield of a reaction. Given the balanced chemical equation, that describes the reaction, there are many equivalent ways to identify the limiting reagent and calculate the excess quantities of other reagents in the reaction. In this article, we will discuss what is limiting agent is, how to find limiting reagents and some limiting reagent questions.

Limiting Reagent Definition

Limiting reagents are defined as the substances which are entirely consumed in the completion of a chemical reaction. They are also referred to as limiting reactants or limiting agents. According to the stoichiometry of chemical reactions, a fixed amount of reactants is necessary for the reaction to complete.

This reactant usually determines when the reaction would stop. The exact amount of reactant that would be needed to react with another element is calculated from the reaction stoichiometry. The limiting reagent depends on the mole ratio and not on the masses of the reactants present.

Consider the following reaction for the formation of ammonia:

3H₂ + N₂ —> 2NH₃

In the reaction shown above, 3 moles of hydrogen gas is required for the reaction with 1 mole of nitrogen gas for the formation of 2 moles of ammonia. But what if, during the time of the reaction, there are only 2 moles of hydrogen gas available with 1 mole of nitrogen?

In this case, the entire quantity of nitrogen cannot be used since the entirety of nitrogen requires 3 moles of hydrogen gas to react. Therefore, the hydrogen gas is limiting the reaction and is hence called the limiting reagent for this reaction.

Limiting Reagent Examples

Let us now look at some of the limiting reagent examples.

Example 

Consider the combustion of benzene which is represented by the following chemical equation:

2C₆H₆(l) + 15 O₂(g) —> 12CO₂(g) + 6HO₂(l)

It means that 15 moles of molecular oxygen O₂ are needed to react with 2 moles of benzene C₆H₆. 

The amount of oxygen that is required for other quantities of benzene is calculated using cross-multiplication. For example, if 1.5 mol C₆H₆ is present, 11.25 mol O₂ is required:

1.5 mol C₆H₆ x [frac{15;mol;O_{2}}{2;mol;C_{6}H_{6}}] = 11.25 mol O₂ 

If in 18 mol O₂ are present, there would be an excess of (18 – 11.25) = 6.75 mol of unreacted oxygen when all of the benzene is consumed. Benzene is, therefore, the limiting reagent.

How to Find Limiting Reagent in a Reaction?

Let us now learn about how to determine limiting reagent in a reaction. 

There are two ways for how to calculate limiting reagent. One method is to find and compare the mole ratio of the reactants that are used in the reaction. Another method is to calculate the grams of products produced from the quantities of reactants in which the reactant which produces the smallest amount of product is the limiting reagent.

Method 1: Finding the limiting reagent by looking at the number of moles of every reactant.

1. First, determine the balanced chemical equation for the given chemical reaction.

2. Then, convert all the given information into moles (by using molar mass as a conversion factor).

3. The next step is to calculate the mole ratio from the given information. Then, compare the calculated ratio to the actual ratio.

4. Use the amount of limiting reactant for calculating the amount of product produced.

5. Lastly, if necessary, calculate how much of the non-limiting agent is left in excess.

Method 2: Finding the limiting reagent by calculating and comparing the amount of product each reactant would produce.

1. The first step is to balance the chemical equation for the given chemical reaction.

2. Then, convert the given information into moles.

3. Use stoichiometry for each individual reactant for finding the mass of product produced.

4. The reactant which produces a lesser amount of product would be the limiting reagent.

5. The reactant which produces a larger amount of product would be the excess reagent.

6. Lastly, for finding the amount of remaining excess reactant, subtract the mass of excess reagent consumed from the total mass given of the excess reagent.

Limiting Reagent Problems

1. Determine the limiting reagent if 76.4 grams of  C₂H₃Br₃ reacts with 49.1 grams of O₂.

4 C₂H₃Br₃ + 11 O₂ —> 8 CO₂ + 6HO₂ + 6Br₂

Solution:

Using method 1,

76.4 g x  [frac{1;mol}{266.72g}] = 0.286 moles of C2H3Br3 

49.1 g x  [frac{1;mol}{32g}] = 1.53 moles of O2 

If you assume that all of the oxygen is used up, 

1.53 x 411, or, 0.556 moles of C2H3Br3 are required. Since there are only 0.286 moles of C2H3Br3 that are available, C2H3Br3 is the limiting reagent here.

Using method 2,

76.4 g C₂H₃Br₃ [frac{1;mol}{266.72g}] x  [frac{8;mol;CO_{2}}{4;mol;C_{2}H_{3}Br_{3}}] x [frac{44.01;g;CO_{2}}{1;mol;CO_{2}}] = 25.2 g CO₂ 

49.1 g O₂ x [frac{1;mol;O_{2}}{32;g;O_{2}}] x [frac{8;mol;CO_{2}}{11;mol;O_{2}}] x [frac{44.01;g;CO_{2}}{1;mol;CO_{2}}] =  49.1 g CO₂ 

Hence, by using any of these methods, C₂H₃Br₃ is the limiting reagent.

In a chemical reaction, the limiting reactant (or limiting reagent) is the reactant that is used first, limiting the amount of product that can be created. There are a variety of methods for determining the limiting reactant, as we saw in the above examples, but they all rely on mole ratios from the balanced chemical equation.

The theoretical yield is the amount of product that can be created depending on the limiting reactant. In fact, the actual yield, or the amount of product obtained, is usually always less than the theoretical yield. The actual yield is commonly represented as a percent yield, indicating how close the actual yield was to the theoretical yield.