Basic Concepts Of Chemistry

Question 1

Define Matter

Answer 1

Anything which has mass and occupies space is called matter.

Question 2

What are the 3 major states of matter? Define their properties.

Answer 2

Solids - Particles and held very closely and tightly. Have definite volume and shape.
Liquids - Particles are close to each other but they can move around. Have definite volume but takes shape of container.
Gases - Particles are very far apart and movement is easy and fast. Neither definite shape nor volume.

Question 3

What are the 3 major states of matter? Define their properties.

Answer 3

Solids - Particles and held very closely and tightly. Have definite volume and shape.
Liquids - Particles are close to each other but they can move around. Have definite volume but takes shape of container.
Gases - Particles are very far apart and movement is easy and fast. Neither definite shape nor volume.

Question 4

Matter can be divided into 2 categories:

Answer 4

Pure Substance: All constituent particles of a substance are same in chemical nature.
Mixture: Contains particles of two or more pure Substances which may be present it in any ratio.

Question 5

A mixture may be:/Define types of mixtures.

Answer 5

Homogeneous: components mix completely. Uniform distribution of particles.

Heterogeneous: Not uniform composition, seperate particles may be visible.

Question 6

Pure Substances can be classified as:

Answer 6

Elements: consists of only one type of atoms. May exist as atom or molecule. Properties of atoms of an element are constant throughout.
Compounds: Two or more atoms of different elements combine together in a definite ratio, molecule of a compound is obtained. Properties of a compound are different from those of it’s constituent elements.

Pure substances can’t be separated by normal physical methods.

Question 7

Define SI unit and it’s constituents:

Answer 7

International System of Units or SI have seven base or fundamental quantities which are:
Length [L] (metre)
Mass [M] (kg)
Time [T] (secs)
Temprature [θ] or [0] (kelvin)
Current [A] or [I] (ampere)
Amount of Substance [N] (mole)
Luminous Intensity [J] or [I] (candela)

Question 8

Prefixes for SI Units?

Answer 8

femto 10-¹⁵
pico 10-¹²
nano 10-⁹
micro 10-⁶

milli 10-³
centi 10-²
deci 10-¹
deca 10
hecto 10²
kilo 10³

mega 10⁶
giga 10⁹
tera 10¹²
peta 10¹⁵

Question 9

Differentiate between Mass and Weight? How is mass of a substance determined?

Answer 9

Mass of a substance is the amount of matter present in it, while weight is the force exerted by gravity on an object.

The mass of a substance is constant, whereas, it’s weight may vary from one place to another due to change in gravity.

Mass of a substance can be determined accurately in the laboratory by using an Analytical Balance.

Question 10

Define Volume. How is it measured in labs?

Answer 10

Volume is the amount of space occupied by a substance, it has the units of l^3.

In the labs, volume of liquids or solutions can be measured by graduated cylinders, burette, pipette, etc. Volumetric Flask is used to prepare a known volume of a solution.

Question 11

Temperature Info?

Answer 11

Temprature has 3 common scales to be measured by: Celsius, Kelvin, Fahrenheit.

Celsius scale is calibrated from 0° to 100° (freezing point to boiling point).
Fahrenheit scale is calibrated between 32° to 212°
The Kelvin scale is a temperature scale that starts at absolute zero, the coldest possible temperature. It’s the base unit of temperature in the International System of Units (SI)

Relations:
C/5 = (F-32)/9
K = C + 273
K = (F − 32) × 5⁄9 + 273.15

In kelvin scale, negative values are not possible.

Question 12

What is exponential notation?

Answer 12

N × n^10

Question 13

The uncertainty in the experimental or the calculated values is indicated by mentioning the number of ________________?

Answer 13

Significant Figures

Question 14

Differentiate between Precision and Accuracy.

Answer 14

Precision refers to the closeness od various measurements for the same quantity.

Accuracy: is the agreement of a particular value to the true value if the result.

Question 15

Rules for addition and subtraction of Significant Figures?

Answer 15

Result cannot have more digits to the right of the decimal point then either of the original numbers.

Question 16

Rules for multiplication and division of Significant Figures?

Answer 16

In these operations, the result must be reported with no more significant figures as in the measurement with the few significant figures.

Question 17

Dimensional Analysis?

Answer 17

Often while calculating there is a need to convert units from one system to another. Method used to accomplish this is called factor label method or unit factor method or dimensional analysis.

Question 18

Laws of Chemical Combinations?

Answer 18

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)Avagadro’s Law

Question 19

Define Law of Conservation of Mass

Answer 19

Put forth by Antoine Lavoiser, this law state that there is no net change in mass during the process meaning matter can neither be created nor destroyed.

This means that the mass of all reactants in a reaction should be equal to mass of all products in a reaction.

2(H2) + (O2) → 2(H2O)

Question 20

Explain volume and shape in Solid, Liquids and Gases.

Answer 20

Solids have definite shape and volume.
Liquids have definite volume but not shape as they take shape of container in which they’re contained.
Gases nether have definite volume nor definite shape.

Question 21

When are the values of Fahrenheit and Celcius scale same?

Answer 21

-40°C = -40F

Question 22

When are the values of F and K scale equal?

Answer 22

574.25

Question 23

What is Law of Definite Proportions?

Answer 23

Given by Joseph Proust, it states that a given compound always contains exactly the same proportion of elements by weight.

Also referred to as the Law of Definite Composition.

Ex: Water (H₂O)
Water is always composed of hydrogen and oxygen in a fixed 1:8 mass ratio.
Hydrogen’s atomic mass = 1 g/mol
Oxygen’s atomic mass = 16 g/mol
In H₂O, the mass of hydrogen is 2 × 1 = 2 g, and oxygen is 16 g, giving a ratio of 2:16 or 1:8.
This ratio remains constant regardless of the water’s source, proving the law.

Question 24

What is Law of Multiple Proportions?

Answer 24

Proposed by Dalton, it state s that if two elements can combined to form more than one compound the masses of one element that combine with a fixed mass of the other element are in the ratio of small whole numbers.

Ex: Water and Hydrogen Peroxide

2(H2) + O2 → 2(H20)
2g 16g 18g

H2 + O2 → H2O2
2g 32g 34g

Here, masses of oxygen (16g and 32g) combine with fixed mass of hydrogen (2g) bearing a simple ratio (16:32 = 1:2).

Question 25

Define Gay Lussac's Law of Combined Volumes.

Answer 25

Proposed by Gay Lusaac, it states that **when gases combine or are produced in a chemical reaction, they do so in a simple ratio by volume. Provided all gases are at the same temperature and pressure.** Ex: Water Hydrogen + Oxygen→Water 100m 50ml 100ml Volumes lf hydrogen and oxygen combine and bear simple ratio- 2:1 This is only the Law of definite proportions by volume. As law of definite proportions was only stated for masses earlier.

Question 26

State postulates of Dalton's Atomic Theory.

Answer 26

1.Matter consists of indivisible atoms. 2.All atoms of a given element have identical properties, including identical mass. Atoms of different elements differ in mass. 3. Compounds are formed when atoms of different elements combined in a fixed ratio. 4. Chemical reactions involve reorganisation of atoms. They are neither created nor destroyed in a chemical reaction.

Question 27

State Avogadro's Law.

Answer 27

Avogadro proposed that: **Equal volumes of all gases at the same temperature and pressure should contain equal number of molecules.** He made a clear distinction between atoms and molecules which is quite understandable in present times.

Question 28

State the drawbacks of Dalton's Atomic Theory.

Answer 28

1. Could not explain the law of gaseous volumes. 2. Could not provide a reason for combining of atoms. 3. Did not explain the existence of isomers and isobars. 4. Contradicts the existence of nuclear reactions which leads to destruction of atoms. 5. Does not explain existence of complex organic compounds (do not form by combining in whole no. ratios). 6. Does not explain existence of sub-atomic particles.

Question 29

Define Atomic Mass.

Answer 29

The present system of atomic masses is based on Carbon-12 as the standard and has been agreed upon. Here, Carbon-12 is one of the isotopes of carbon and can be represented as ¹²C. In this system, **¹²C is assigned the mass of exactly 12 atomic mass unit (amu) and mass of all other atoms are given relative to this standard.** **One atomic mass unit is defined as a mass exactly equal to one-twelfth (1/12) of the mass of one ¹²C atom.** At present, "amu" is replaced by 'u' (unified mass). We use avg atomic masses of elements for calculation of atomic mass.

Question 30

Average Atomic Mass?

Answer 30

Elements exist as more than one isotype. When we take into account the existence of these isotypes and their relative abundance, the average atomic mass of a element can be computed. In the periodic table, the atomic masses represented for each element are actually there average atomic masses.

Question 31

Define Formula Mass.

Answer 31

Formula mass refers to the sum of the atomic masses of all atoms present in a chemical formula, essentially representing the mass of a single "formula unit" of a compound.

Question 32

Explain mole concept.

Answer 32

In the SI system, mole (mol) was introduced as the 7th base quantity for the amount of a substance. One mole contains exactly **6.022 × 10²³** elementary entities. This number is the fixed numerical value of the **Avogadro constant (NA), when expressed in the unit mol-¹ and is called the avogadro number.**

Question 33

Explain molar mass.

Answer 33

**The mass of one mole of a substance in grams is called its molar mass**. The molar mass in grams is numerically equal to atomic/molecular/formula mass in u. **molar mass = w/n Where, w = mass in grams, n = no. of moles**

Question 34

What is Percentage Composition/ Mass % of an element in a compound?

Answer 34

Mass % of an element: **(mass of element in that compound × 100)/(molar mass of compound)** Or, =

Question 35

Define Empirical Formula.

Answer 35

**Empirical formula represents the simplest whole number ratio of various atoms present in a compound.**

Question 36

Define Molecular Formula

Answer 36

**Molecular formula shows the exact number of different types of atoms present in a molecule of a compound.**

Question 37

Difference between Molecular and Empirical Formula?

Answer 37

An empirical formula represents the simplest whole number ratio of various atoms present in a compound, whereas, the molecular formula shows the exact number of different types of atoms present in a molecule of a compound.

Question 38

Define Stoichiometry.

Answer 38

Stoichiometry deals with the calculation of masses (sometimes volume also) of the reactants and the products involved in the chemical reaction. Coefficients such as 2 for O2 in formation of water are known as **Stoichiometry Coefficients**. Stoichiometry coefficient represent the number of molecules and mole taking part in the reaction or formed in the reaction by ratio.

Question 39

What is limiting reagent.

Answer 39

The reactant which is present in the least amount gets consumed after some time. Then after that the reaction does not take place whatever be the amount of the other reactant. Hence, **the reactant which gets consumed first in a reaction and limits the amount of product formed is called the limiting reagent.**

Question 40

What is Mass per cent?

Answer 40

Mass % = (Mass of Solute/Mass of Solution) × 100

Question 41

Define Mole Fraction?

Answer 41

Mole fraction can be defined as the number of moles of a A upon the number of moles of solution for mole fraction of A. Mole Fraction of A = **(no. of moles of A/ no. of moles in a solution )** = **(Na/Na+Nb)**

Question 42

Define Molarity.

Answer 42

Molarity (M) = **(No. of moles of solute/ Volume of solution in litres)** =

Question 43

What is Molality?

Answer 43

Molality(m) = **(No. of moles of solute / Mass of solvent in kg)**