Options not available — refer to textbook.
Options not available — refer to textbook.
Options not available — refer to textbook.
Options not available — refer to textbook.
Options not available — refer to textbook.
Options not available — refer to textbook.
Options not available — refer to textbook.
Options not available — refer to textbook.
Options not available — refer to textbook.
| Column A | Column B |
|---|---|
| 8 g of O2 | 0.25 moles |
| 4 g of H2 | 2 moles |
| 52 g of He | 13 moles |
| 112 g of N2 | 4 moles |
| 35.5 g of Cl2 | 0.5 moles |
| # | Statement | Answer | Correction (if False) |
|---|---|---|---|
| 1 | Two elements sometimes can form more than one compound. | True | - |
| 2 | Noble gases are Diatomic | False | Noble gases are monoatomic. |
| 3 | The gram atomic mass of an element has no unit | False | The gram atomic mass of an element is expressed in grams. |
| 4 | 1 mole of Gold and Silver contain same number of atoms | True | - |
| 5 | Molar mass of CO2 is 42g. | False | Molar mass of CO2 is 44 g mol^-1. |
Answer: A and R are correct, and R explains A.
Answer: A and R are correct, and R explains A. Chlorine has isotopes Cl-35 and Cl-37 in unequal natural abundance, so its average relative atomic mass is 35.5.
The relative atomic mass of an element is a dimensionless quantity that represents the average mass of atoms of that element compared to a standard reference. Specifically, it is defined as the ratio of the average mass of one atom of the element to one-twelfth (1/12) of the mass of an atom of the carbon-12 isotope. Carbon-12 is used as the standard because it is abundant and stable. This definition allows for a consistent and comparable scale for the masses of different elements, facilitating stoichiometric calculations and understanding chemical composition. It is important to note that relative atomic mass is a ratio and therefore has no units.
Oxygen, a vital element for life, exists in nature as a mixture of three stable isotopes. These isotopes have the same number of protons but differ in their number of neutrons, leading to slightly different atomic masses. The most abundant isotope is Oxygen-16 (¹⁶O), which has a mass of approximately 15.9949 atomic mass units (amu) and constitutes about 99.757% of naturally occurring oxygen. The second isotope is Oxygen-17 (¹⁷O), with a mass of approximately 16.9991 amu, and it is present in a very small abundance of about 0.038%. The third stable isotope is Oxygen-18 (¹⁸O), having a mass of approximately 17.9992 amu, and it accounts for about 0.205% of natural oxygen. These isotopic abundances are crucial for various scientific applications, including dating and tracer studies.
Atomicity refers to the number of atoms that are chemically bonded together to form a single molecule of a particular element or compound. For instance, in the case of elements, if a molecule consists of only one atom, it is monatomic (like Helium, He). If it consists of two atoms, it is diatomic (like Oxygen, O₂). Molecules made of three atoms are triatomic (like Ozone, O₃), and so on. For compounds, atomicity is the total count of all atoms of different elements present in one molecule. For example, a molecule of water (H₂O) has an atomicity of three (two hydrogen atoms and one oxygen atom).
Heterodiatomic molecules are molecules composed of two atoms of different elements chemically bonded together. These molecules are diatomic because they contain exactly two atoms. Examples of such molecules include hydrogen chloride (HCl), where one hydrogen atom is bonded to one chlorine atom, and carbon monoxide (CO), where one carbon atom is bonded to one oxygen atom. In both cases, the molecule consists of two distinct types of atoms, making them heterodiatomic. Other examples include nitric oxide (NO) and hydrogen bromide (HBr).
The molar volume of a gas is a fundamental concept in chemistry that defines the space occupied by a specific amount of gas under standard conditions. It is defined as the volume that one mole of any gas occupies at Standard Temperature and Pressure (STP). At STP, which is typically defined as a temperature of 0°C (273.15 K) and a pressure of 1 atmosphere (atm), one mole of any ideal gas occupies a volume of 22.4 liters (L) or 22,400 milliliters (mL). This principle, known as Avogadro's Law, is extremely useful for calculations involving gases, allowing chemists to relate the amount of a gas (in moles) to its volume without needing to know its specific identity.
Percentage of nitrogen = (14/17) x 100 = 82.35%.
Number of moles = 0.18/18 = 0.01 mol.
Number of molecules = 0.01 x 6.023 x 10^23 = 6.023 x 10^21 molecules.
1 mole of nitrogen gas = 28 g; 3 moles of hydrogen gas = 3 x 2 = 6 g; 2 moles of ammonia = 2 x 17 = 34 g.
So, 28 g of N2 reacts with 6 g of H2 to form 34 g of NH3.
(ii) For 1.51 x 10^23 molecules of NH4Cl, number of moles = number of molecules/Avogadro number = (1.51 x 10^23)/(6.023 x 10^23) = 0.25 mole approximately.
(i) Atom is divisible into subatomic particles such as electrons, protons and neutrons.
(ii) Atoms of the same element may have different atomic masses; these are isotopes, such as Cl-35 and Cl-37.
(iii) Atoms of different elements may have the same mass number; these are isobars, such as Ar-40 and Ca-40.
(iv) Atoms of one element can be transmuted into atoms of another element by nuclear reactions.
(v) Atom is the smallest particle that takes part in a chemical reaction.
(vi) Mass can be converted into energy in nuclear reactions, E = mc^2.
By Avogadro's law, equal volumes contain equal number of molecules. Since hydrogen is diatomic, the mass of one molecule of hydrogen is twice the mass of one hydrogen atom.
Therefore, vapour density = molecular mass/2.
Hence, relative molecular mass = 2 x vapour density.
(i) Number of moles of CaCO3 involved = 1 mole.
(ii) Gram molecular mass of CaCO3 = 40 + 12 + (16 x 3) = 100 g mol^-1.
(iii) Number of moles of CO2 formed = 1 mole.
(ii) 3 moles of Cl2: molar mass = 70.9 g mol^-1, mass = 3 x 70.9 = 212.7 g.
(iii) 5 moles of S8: molar mass = 32 x 8 = 256 g mol^-1, mass = 5 x 256 = 1280 g.
(iv) 4 moles of P4: molar mass = 31 x 4 = 124 g mol^-1, mass = 4 x 124 = 496 g.
Percentage of Ca = (40/100) x 100 = 40%.
Percentage of C = (12/100) x 100 = 12%.
Percentage of O = (48/100) x 100 = 48%.
Mass of oxygen = 16 x 12 = 192 g.
Percentage of oxygen = (192/342) x 100 = 56.14%.
Average atomic mass = [10.01294x + 11.009305(100 - x)]/100 = 10.804.
Solving, x = 20.636%. Therefore B-10 = 20.636% and B-11 = 79.364%.
Study Smarter, Score Higher.
Revise this Samacheer Class 10 Science topic, then continue with the Revision Challenge.