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Sajjad Hussain Ali Mirani
Introduction to fundamental concepts of chemistry
Chemistry
It is a branch of science which deals with the properties, composition and the structure of matter.
Atom
It is the smallest particle of an element which can exist with all the properties of its own element but it cannot exist in atmosphere alone.
Molecule
When two or more than two atoms are combined with each other a molecule is formed. It can exist freely in nature.
Formula Weight
It is the sum of the weights of the atoms present in the formula of a substance.
Molecular Weight
It is the sum of the atomic masses of all the atoms present in a molecule.
Stat Chapter 1 Fundamental Concept Of Chemistry
Chemistry Sajjad Hussain Ali Mirani (SHAM)
Chemistry is the branch of science which deals with the properties, composition and structure of matter. Study of chemistry also includes the laws and principles related to the structure and inter-relations of elements and compound.
Chemistry has the task of investigating the materials of which our universe is made. Chemistry investigates chemical changes, conditions under which chemical changes occur. Chemistry also deals with the way in which similar changes can be brought about in laboratory and on a large scale in industries.
Chemistry is a very vast field. Chemistry is divided into a number of branches such as Organic chemistry, Inorganic chemistry, Physical chemistry, biochemistry, Applied chemistry, Nuclear chemistry etc.
Significant Figures
Significant figures are the reliable digits in a number or measurement
which are known with certainty.
Significant figures show the accuracy in measurements. We can understand the precision of a measurement if we know exactly the significant figures in the measurement.
A measurement that contains more number of significant figures is more accurate than a measurement that contains less number of Significant figures.
For example:Radius of a bob is 3.3679 cm and that of the other is 3.36 cm. In this situation the first measurement is the most accurate as it has more number of significant figures.
Rules Of Significant Figures
In order to determine significant figures in a number we must follow the following rules:
(1) All the non-zero digits are significant figures.
For Example:
3.456 has four significant figures.
12.3456 has six significant figures.
0.34 has two significant figures.
(2) Zeros between non-zero digits are significant.
For Example:
2306 has four significant figures.
20,0894 has six significant figures.
(3) Zeros locating the position of decimal in numbers of magnitude less than one are not significant.
For Example:
0.2224 has only one significant figures.
0.0000034 has two significant figures.
(4) Final zeros to the right of the decimal point are significant.
For Example:
3.0000 has five significant figures.
1002.00 has six significant figures.
(5) Zeros that locate decimal point in numbers greater than one are not significant.
For Example:
30000 has only one significant figure.
120000 has two significant figures.
Rules Of Rounding Off Data
Rule # 1:
If the digit to be dropped is greater than 5, then add "1" to the last digit to be retained and drop all digits farther to the right.
For example:
3.677 is rounded off to 3.68 if we need three significant figures in measurement.
3.677 is rounded off to 3.7 if we need two significant figures in measurement.
Rule # 2:
If the digit to be dropped is less than 5, then simply drop it without adding any number to the last digit.
For example:
6.632 is rounded off to 6.63 if we need three significant figures in measurement.
6.632 is rounded off to 6.6 if we need two significant figures in measurement.
Rule # 3:
If the digit to be dropped is exactly 5 then:
(A) If the digit to be retained is even, then just drop the "5".
For example:
6.65 is rounded off to 6.6 if we need two significant figures in measurement.
3.4665 is rounded off to 6.466 if we need four significant figures in measurement.
(B) If the digit to be retained is odd, then add "1" to it.
For example:
6.35 is rounded off to 6.4 if we need two significant figures in measurement.
3.4675 is rounded off to 6.468 if we need four significant figures in measurement.
Remember: Zero is an even number
3.05 is rounded off to 3.0 if we need two significant figures in measurement.
Use of significant figures in
addition and subtraction
In addition and subtraction we consider the significant figures on the right side of decimal point. This
means that only as many digits are to be retained to the right side of decimal point as the number with fewest digits to the right of the decimal point.
For example:
4.345 + 23.5 =27.845 (actual answer by using calculator)
Answer after rounding off: 27.8
Use of significant figures in
multiplication and division
In multiplication and division , the number obtained after calculation of two or more numbers must have
no more significant figure than that number used in multiplication or division.
For example:
4.3458 x 2.7 =11.73366(actual answer by using calculator)
Answer after rounding off: 12(because 2.7 has only two significant figures)
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Sahm Notes www.icusj.blogspot.com
Error
An error is defined as:
"The difference between the measured value and the actual value."
If two persons use the same instrument for measurement for finding the same measurement, it is not essential that they may get the same results. There may arises a difference between their measurements. This difference is referred to as an "ERROR".
Types Of Error
Errors can be divided into three categories:
(1) Personal Error
(2) Systematic Error
(3) Random Error
Personal Error
An error comes into play because of faulty procedure adopted by by the observer is called "PERSONAL ERROR".
Personal error comes into existence due to making an error in reading a scale. It is due to faulty procedure adopted by the person making measurement.
Systematic Error
The type of error arises due to defect in the measuring device is known as "SYSTEMATIC ERROR"
Generally it is called "ZERO ERROR". it may be positive or negative error. Systematic error can be removed by correcting measurement device.
Random Error
The error produced due to sudden change in experimental conditions is called "RANDOM ERROR".
For example:
During sudden change in temperature, change in humidity, fluctuation in potential difference(voltage).
It is an accidental error and is beyond the control of the person making measurement.
Atomic Mass
Atomic mass is defined as :
"The mass of one atom of the element compared with the mass of one atom of C12"
Atomic mass is a ratio therefore it has no unit. Generally atoms mass is expressed in ATOMIC MASS UNIT(a.m.u).
One atomic mass unit is equal to 1/12 of the mass of a C12 atom.
Empirical Formula
"Empirical Formula is that formula which expresses the relative number
of each kind of atoms present in the molecule of a compound"
OR
"The formula of a compound which expresses the ratio in which atoms of
different elements are combined in a molecule"
Empirical formula only indicates atomic ratios but it does not indicate actual number of atoms of different kinds present in the molecule of a compound.
Two or more compound may have same empirical formula.
Empirical formula is determined by experiment.
Molecular Formula
The molecular formula of a compound is defined as:
"The formula of a compound which not only expresses the relative number of atoms of each
kind but also expresses the actual number of atoms of each element present in one molecule".
Molecular formula and empirical formula of a compound are related as:
MOLECULAR FORMULA = (EMPIRICAL FORMULA)n
Where "n" is an integer and is given by:
n = molecular mass of compound / Empirical formula mass of compound
Molecular formula of propane = C3H8.
Molecular formula of sugar = C12H22O11.
Limiting Reactant
The limiting reactant is defined as:
" The substance which produces least amount of products when it is
completely consumed in a chemical reaction".
In stoichiometry when more than one reactant is involved in a chemical reaction, it is not so simple to get actual result of the stoichiometric problem by making relationship between any one of the reactant and product, which are involved in the chemical reaction. As we know that when any one of the reactant is completely used or consumed the reaction is stopped no matter the other reactants are present in very large quantity. This reactant which is totally consumed during the chemical reaction due to which the reaction is stopped is called limiting reactant.
Limiting reactant help us in calculating the actual amount of product formed during the chemical reaction. To understand the concept the limiting reactant consider the following calculation.
Problem
We are provided 50 gm of H2 and 50 gm of N2. Calculate how many gm of NH3 will be formed when the reaction is irreversible.
The equation for the reaction is as follows.
N2 + 3 H2 ----> 2 NH3
Solution
In this problem moles of N2 and H2 are as follows
Moles of N2 = Mass of N2 / Mol. Mass of N2
= 50 / 28
= 1.79
Moles of H2 = Mass of H2 / Mol. Mass of H2
= 50 / 2
= 25
So, the provided moles for the reaction are
nitrogen = 1.79 moles and hydrogen = 25 moles
But in the equation of the process 1 mole of nitrogen require 3 mole of hydrogen. Therefore the provided moles of nitrogen i.e. 1.79 require 1.79 x 3 moles of hydrogen i.e. 5.37 moles although 25 moles of H2 are provided but when nitrogen is consumed the reaction will be stopped and the remaining hydrogen is useless for the reaction so in this problem N2 is a limiting reactant by which we can calculate the actual amount of product formed during the reaction.
N2 + 3 H2 ----> 2 NH3
1 mole of N2 gives 2 moles of NH3
1.79 mole of N2 gives 2 x 1.79 moles of NH3
= 3.58 moles of NH3
Mass of NH3 = Moles of NH3 x Mol. Mass
= 3.58 x 17
= 60.86 gm of NH
Sahm Notes www.icusj.blogspot.com
Define the following terms:
Mole
"The atomic mass or molecular mass of a substance expressed in grams is called mole".
OR
A mole may also be defined as
" the gram atomic mass or gram molecular mass or gram formula mass
of a substance that contains 6.02 x 1023 particles."
OR
"The mass in grams of the atoms or molecules or ions which contains
Avogadro's number of particles i.e. 6.02 x 1023 particles."
For example:
(1) Atomic mass of Carbon = 12 a.m.u.
Therefore 12 gram of carbon = 1 mole of carbon.
(2) molecular mass of nitrogen = 28 a.m.u.
Therefore 28 gram of nitrogen = 1 mole of nitrogen.
(3) Formula mass of NaCl = 58.5 a.m.u.
Therefore 58.5 gram of NaCl = 1 mole of NaCl.
Mole is denoted by "n".
Formula
Number of moles of substance = mass of substance (in grams) / molecular mass or atomic mass or formula mass
Avogadro's Number
One mole of any substance contains equal number of particles (atoms or molecules or ions).Value of this number is 6.02 x 1023. This constant value or number is referred to as "AVOGADRO'S NUMBER"
Example
1 gm mole of Na contain 6.02 x 10(23) atoms of Na.
1 gm mole of Sulphur = 6.02 x 10(23) atoms of Sulphur.
1 gm mole of H2SO4 = 6.02 x 10(23) molecules H2SO4
1 gm mole of H2O = 6.02 x 10(23) molecules of H2O
On the basis of Avogadro's Number "mole" is also defined as
Mass of 6.02 x 10(23) molecules, atoms or ions in gram is called mole.
Determination Of The Number Of Atoms Or Molecules In The Given Mass Of A Substance
Example 1
Calculate the number of atoms in 9.2 gm of Na.
Solution
Atomic mass of Na = 23 a.m.u
If we take 23 gm of Na, it is equal to 1 mole.
23 gm of Na contain 6.02 x 10(23) atoms
1 gm of Na contain 6.02 x 10(23) / 23 atoms
9.2 gm of Na contain 9.2 x 6.02 x 10(23) /23
= 2.408 x 10(23) atoms of Na
Determination Of The Mass Of Given Number Of Atoms Or Molecules Of A Substance
Example 2
Calculate the mass in grams of 3.01 x 10(23) molecules of glucose.
Solution
Molecular mass of glucose = 180 a.m.u
So when we take 180 gm of glucose it is equal to one mole So,
6.02 x 10(23) molecules of glucose = 180 gm
1 molecule of glucose = 180 / 6.02 x 10(23) gm
3.01 x 10(23) molecules of glucose = 3.01 x 10(23) x 180 / 6.02 x 10(23)
= 90 gm
It is denoted by "NA". Sahm Notes www.icusj.blogspot.com
Stoichiometry
Stoichiometry refers to chemical calculations.
Stoichiometry is defined as:
"The quantitative relationship among the reactants and products in a balanced chemical equation".
Assumption of Stoichiometric Calculations
There are two important assumptions for stoichiometric calculations:
(1) Reactants are completely converted into products.
(2) No supplementary or side-reactions occur.
Suppose we want to calculate the mass of CO2 formed when a given mass of carbon burns in air.
The reaction is:
C + O2 CO2
In actual practice, it is possible that we get less amount of CO2 than the calculated mass of CO2.This is because that the given mass of carbon can also form CO in addition to CO2.
2C + O2 2CO
This means that we have to avoid the formation of carbon monoxide.
Types of stoichiometric calculations
Stoichiometric calculations can be divided into three categories.
(1) Mass - Mass Relationship
(2) Mass - Volume Relationship
(3) Volume - Volume Relationship
Mass - Mass Relationship
In this relationship we can determine the unknown mass of a reactant or product from a given mass of teh substance involved in the chemical reaction by using a balanced chemical equation.
Example
Calculate the mass of CO2 that can be obtained by heating 50 gm of limestone.
Solution
Step I - Write a Balanced Equation
CaCO3 ----> CaO + CO2
Step II - Write Down The Molecular Masses And Moles Of Reactant & Product
CaCO3 ----> CaO + CO2
Method I - MOLE METHOD
Number of moles of 50 gm of CaCO3 = 50 / 100 = 0.5 mole
According to equation
1 mole of CaCO3 gives 1 mole of CO2
0.5 mole of CaCO3 will give 0.5 mole of CO2
Mass of CO2 = Moles x Molecular Mass
= 0.5 x 44
= 22 gm
Method II - FACTOR METHOD Sahm Notes www.icusj.blogspot.com
From equation we may write as
100 gm of CaCO3 gives 44 gm of CO2
1 gm of CaCO3 will give 44/100 gm of CO2
50 gm of CaCO3 will give 50 x 44 / 100 gm of CO2
= 22 gm of CO2
Mass - Volume Relationship
The major quantities of gases can be expressed in terms of volume as well as masses. According to Avogardro One gm mole of any gas always occupies 22.4 dm3 volume at S.T.P. So this law is applied in mass-volume relationship.
This relationship is useful in determining the unknown mass or volume of reactant or product by using a given mass or volume of some substance in a chemical reaction.
Example
Calculate the volume of CO2 gas produced at S.T.P by combustion of 20 gm of CH4.
Solution
Step I - Write a Balanced Equation
CH4 + 2 O2 ----> CO2 + 2 H2O
Step II - Write Down The Molecular Masses And Moles Of Reactant & Product
CH4 + 2 O2 ----> CO2 + 2 H2O
Method I - MOLE METHOD
Convert the given mass of CH4 in moles
Number of moles of CH4 = Given Mass of CH4 / Molar Mass of CH4
From Equation
1 mole of CH4 gives 1 moles of CO2
1.25 mole of CH4 will give 1.25 mole of CO2
No. of moles of CO2 obtained = 1.25
But 1 mole of CO2 at S.T.P occupies 22.4 dm3
1.25 mole of CO2 at S.T.P occupies 22.4 x 1.25
= 28 dm3
Method II - FACTOR METHOD Sahm Notes www.icusj.blogspot.com
Molecular mass of CH4 = 16
Molecular mass of CO2 = 44
According to the equation
16 gm of CH4 gives 44 gm of CO2
1 gm of CH4 will give 44/16 gm of CO2
20 gm of CH4 will give 20 x 44/16 gm of CO2
= 55 gm of CO2
44 gm of CO2 at S.T.P occupy a volume 22.4 dm3
1 gm of CO2 at S.T.P occupy a volume 22.4/44 dm3
55 gm of CO2 at S.T.P occupy a volume 55 x 22.4/44
= 28 dm3
Volume - Volume Relationship
This relationship determine the unknown volumes of reactants or products from a known volume of other gas.
This relationship is based on Gay-Lussac's law of combining volume which states that gases react in the ratio of small whole number by volume under similar conditions of temperature & pressure.
Consider this equation
CH4 + 2 O2 ----> CO2 + 2 H2O Sahm Notes www.icusj.blogspot.com
In this reaction one volume of CH4 gas reacts with two volumes of oxygen gas to give one volume of CO2 and two volumes of H2O
Examples
What volume of O2 at S.T.P is required to burn 500 litres (dm3) of C2H4 (ethylene)?
Solution
Step I - Write a Balanced Equation
C2H4 + 3 O2 ----> 2 CO2 + 2 H2O
Step II - Write Down The Moles And Volume Of Reactant & Product
C2H4 + 3 O2 ----> 2 CO2 + 2 H2O
According to Equation
1 dm3 of C2H4 requires 3 dm3 of O2
500 dm3 of C2H4 requires 3 x 500 dm3 of O2
= 1500 dm3 of O2
Gay-Lussac's Law Of Combining Volume
According to Gay-Lussac's Law Of Combining Volume:
"Gases react in the ratio of small whole numbers by volume under
similar conditions of temperature and pressure".
For example:
Nitrogen and hydrogen react to form ammonia gas.
N2 + 3H2 2NH3
In this example one volume of nitrogen gas reacts with three volumes of hydrogen gas to produce two volumes of ammonia gas.
Fill In The Blanks Give Suitable Ans And Fill With Pencil
A Product Of Sham Notes
1. 35 gm of oxygen at S.T.P occupies a volume of __________ dm3.
2. 0.00051 contains __________ significant figures.
3. 0.00054471 contains __________ significant figures.
4. The volume of 1 gm of hydrogen gas at S.T.P is __________.
5. The oxidation number Mn in KMnO4 is __________.
6. The product of ionic concentration in a saturated solution is called __________.
7. 16 gm of oxygen at S.T.P occupies a volume of __________ dm3.
8. The shape of the orbital for which l = 0 is __________.
9. The radius of Cl-1 is __________ than the radius of Cl0.
10. Sp2 hybridization is also known as __________.
11. The value of 1 Debye is __________.
12. The reactions catalyzed by sunlight are called __________.
13. The blue colour of CuSO4 is due to the presence of __________.
14. The force of attraction between the liquid molecules and the surface of container is called __________.
15. The heat of neutralization of a strong acid and a strong base is __________.
16. C º C triple bond is __________. C = C double bond length.
17. The ions having the same electronic configuration are called iso electronic.
18. On heating, if a solid changes directly into vapours without changing into the liquid state, the phenomenon is called __________.
19. Each orbital in an atom can be completely described by __________.
20. In a molecule of alkene, __________ restricts the rotation of the group of atoms at either end of the molecule.
21. Density, refractive index and vapour pressure are __________ properties.
22. The addition of HCl to H2 solution __________ the ionization of H2S.
23. The reaction of cation or anion (or both) with water so as to change its __________ is known as Hydrolysis.
24. A reaction with higher activation energy will start at __________ temperature.
25. 6.02 x 1023 has __________ significant figures.
26. The internal resistance in the flow of liquid is called __________.
27. A catalyst increases the velocity of a reaction but decreases the __________.
Chapter 1
Introduction to Fundamental Concepts
1. 1 mole of a gas at S.T.P occupies a volume of __________.
2. A gas occupying a volume of 22.4 dm3 at S.T.P contains __________ molecules.
3. A formula, which gives the relative number of atoms in the molecule of a compound, is called __________.
4. A formula which gives the actual number of all kinds of atoms present in the molecule of compound is termed as __________.
5. The chemical formula that not only gives the actual number of atoms but also shows the arrangement of different atoms present in the molecule is called __________.
6. Atomic weight or molecular weight expressed in grams is known as __________.
7. 2 moles of H2O contain __________ grams and __________ number of molecules.
8. Any thing that occupies space and has __________ is called matter.
9. Volume of one __________ mole of a gas at S.T.P is 22.4 cubic feet.
10. A ton mole of iron is equal to __________ tons.
11. The force with which the earth attracts a body is called the __________ of the body.
12. A pure substance contains __________ kind of molecules.
13. The smallest indivisible particle of matter is called __________.
14. The atomic number is equal to the number of __________ in nucleus.
15. The atomic mass is the total number of protons and __________ in an atom of the element.
16. The average weight of atoms of an element as compared to the weight of one atom of __________ is called the atomic mass.
17. 1.0007 contains __________ significant figures.
18. The figure 24.75 will be rounded off to __________.
19. __________ means that the readings and measurements obtained in different experiments are very close to each other.
20. __________ means that the results obtained in different experiments are very close to the accepted values.
21. The degree of a measured quantity __________ with increasing number of significant figures in it.
22. The atomic mass of sodium is __________. Sahm Notes www.icusj.blogspot.com
23. The symbolic representation of a molecule of a compound is called __________.
24. Molecular formula of CHCl3 and its Empirical formula is __________.
25. Molecular formula of benzene is C6H6 and its empirical formula is __________.
26. 58.5 is the __________ of NaCl.
27. 4.5 gms of nitrogen will have __________ molecules.
28. 28 gms of nitrogen will have __________ molecules.
29. 2 moles of SO2 is equal to __________ gms.
30. 1000 gms of H2O is equal to __________ moles.
31. The reactions, which proceed in both directions, are called __________.
32. The reactions, which proceed in forward directions only, are called __________ reactions.
33. The __________ reactions are completed after some time.
34. 0.0006 has __________ significant figures
35. 7.40 x 108 has __________ significant figures.
36. 7 x 108 has __________ significant figures.
37. Usually Molecular formula is simple multiple of the __________.
38. 0.1 mole of H2O contains __________ molecules of H2O.
39. Mass of 3.01 x 1022 molecules of CO2 is __________.
40. __________ is the branch of science which deals with the properties, composition and structure of matter.
41. None zero digits are all __________.
42. The integer part of logarithm is called __________.
43. The decimal fraction of logarithm is called __________.
44. __________ is the amount of substance, which contains as many number of particles as there are in 12 gms of Carbon.
45. 6.02 x 1023 is called the __________.
46. The accuracy of measurement depends on the number of __________.
47. __________ is the branch of chemistry that deals with quantitative relationships among the substances undergoing chemical changes.
48. The sum of atomic weights of all the elements present in molecular formula is called the __________.
49. __________ is the sum of atomic weights of the elements represented by the Empirical formula of the compound.
50. Very small and very large quantities are expressed in terms of __________.
51. In rounding off __________ figure is dropped.
52. Mole is the quantity, which has __________ particle of the substance.
53. For three significant figures, 25.55 is rounded off to __________.
54. The S.I unit of a mass is __________.
55. Mass of 6.02 x 1023 molecules of NaCl is __________ gm.
56. 1 mole of NaOH is __________ gm of NaOH.
57. Formula weight is used for __________ substances.
58. The word S.I stands for __________. Sahm Notes www.icusj.blogspot.com
59. 4.5 gms of water will have __________ molecules.
60. 0.0087 has __________ significant figure.
icusj.boards.net
Sajjad Hussain Ali Mirani
Introduction to fundamental concepts of chemistry
Chemistry
It is a branch of science which deals with the properties, composition and the structure of matter.
Atom
It is the smallest particle of an element which can exist with all the properties of its own element but it cannot exist in atmosphere alone.
Molecule
When two or more than two atoms are combined with each other a molecule is formed. It can exist freely in nature.
Formula Weight
It is the sum of the weights of the atoms present in the formula of a substance.
Molecular Weight
It is the sum of the atomic masses of all the atoms present in a molecule.
Stat Chapter 1 Fundamental Concept Of Chemistry
Chemistry Sajjad Hussain Ali Mirani (SHAM)
Chemistry is the branch of science which deals with the properties, composition and structure of matter. Study of chemistry also includes the laws and principles related to the structure and inter-relations of elements and compound.
Chemistry has the task of investigating the materials of which our universe is made. Chemistry investigates chemical changes, conditions under which chemical changes occur. Chemistry also deals with the way in which similar changes can be brought about in laboratory and on a large scale in industries.
Chemistry is a very vast field. Chemistry is divided into a number of branches such as Organic chemistry, Inorganic chemistry, Physical chemistry, biochemistry, Applied chemistry, Nuclear chemistry etc.
Significant Figures
Significant figures are the reliable digits in a number or measurement
which are known with certainty.
Significant figures show the accuracy in measurements. We can understand the precision of a measurement if we know exactly the significant figures in the measurement.
A measurement that contains more number of significant figures is more accurate than a measurement that contains less number of Significant figures.
For example:Radius of a bob is 3.3679 cm and that of the other is 3.36 cm. In this situation the first measurement is the most accurate as it has more number of significant figures.
Rules Of Significant Figures
In order to determine significant figures in a number we must follow the following rules:
(1) All the non-zero digits are significant figures.
For Example:
3.456 has four significant figures.
12.3456 has six significant figures.
0.34 has two significant figures.
(2) Zeros between non-zero digits are significant.
For Example:
2306 has four significant figures.
20,0894 has six significant figures.
(3) Zeros locating the position of decimal in numbers of magnitude less than one are not significant.
For Example:
0.2224 has only one significant figures.
0.0000034 has two significant figures.
(4) Final zeros to the right of the decimal point are significant.
For Example:
3.0000 has five significant figures.
1002.00 has six significant figures.
(5) Zeros that locate decimal point in numbers greater than one are not significant.
For Example:
30000 has only one significant figure.
120000 has two significant figures.
Rules Of Rounding Off Data
Rule # 1:
If the digit to be dropped is greater than 5, then add "1" to the last digit to be retained and drop all digits farther to the right.
For example:
3.677 is rounded off to 3.68 if we need three significant figures in measurement.
3.677 is rounded off to 3.7 if we need two significant figures in measurement.
Rule # 2:
If the digit to be dropped is less than 5, then simply drop it without adding any number to the last digit.
For example:
6.632 is rounded off to 6.63 if we need three significant figures in measurement.
6.632 is rounded off to 6.6 if we need two significant figures in measurement.
Rule # 3:
If the digit to be dropped is exactly 5 then:
(A) If the digit to be retained is even, then just drop the "5".
For example:
6.65 is rounded off to 6.6 if we need two significant figures in measurement.
3.4665 is rounded off to 6.466 if we need four significant figures in measurement.
(B) If the digit to be retained is odd, then add "1" to it.
For example:
6.35 is rounded off to 6.4 if we need two significant figures in measurement.
3.4675 is rounded off to 6.468 if we need four significant figures in measurement.
Remember: Zero is an even number
3.05 is rounded off to 3.0 if we need two significant figures in measurement.
Use of significant figures in
addition and subtraction
In addition and subtraction we consider the significant figures on the right side of decimal point. This
means that only as many digits are to be retained to the right side of decimal point as the number with fewest digits to the right of the decimal point.
For example:
4.345 + 23.5 =27.845 (actual answer by using calculator)
Answer after rounding off: 27.8
Use of significant figures in
multiplication and division
In multiplication and division , the number obtained after calculation of two or more numbers must have
no more significant figure than that number used in multiplication or division.
For example:
4.3458 x 2.7 =11.73366(actual answer by using calculator)
Answer after rounding off: 12(because 2.7 has only two significant figures)
|
Sahm Notes www.icusj.blogspot.com
Error
An error is defined as:
"The difference between the measured value and the actual value."
If two persons use the same instrument for measurement for finding the same measurement, it is not essential that they may get the same results. There may arises a difference between their measurements. This difference is referred to as an "ERROR".
Types Of Error
Errors can be divided into three categories:
(1) Personal Error
(2) Systematic Error
(3) Random Error
Personal Error
An error comes into play because of faulty procedure adopted by by the observer is called "PERSONAL ERROR".
Personal error comes into existence due to making an error in reading a scale. It is due to faulty procedure adopted by the person making measurement.
Systematic Error
The type of error arises due to defect in the measuring device is known as "SYSTEMATIC ERROR"
Generally it is called "ZERO ERROR". it may be positive or negative error. Systematic error can be removed by correcting measurement device.
Random Error
The error produced due to sudden change in experimental conditions is called "RANDOM ERROR".
For example:
During sudden change in temperature, change in humidity, fluctuation in potential difference(voltage).
It is an accidental error and is beyond the control of the person making measurement.
Atomic Mass
Atomic mass is defined as :
"The mass of one atom of the element compared with the mass of one atom of C12"
Atomic mass is a ratio therefore it has no unit. Generally atoms mass is expressed in ATOMIC MASS UNIT(a.m.u).
One atomic mass unit is equal to 1/12 of the mass of a C12 atom.
Empirical Formula
"Empirical Formula is that formula which expresses the relative number
of each kind of atoms present in the molecule of a compound"
OR
"The formula of a compound which expresses the ratio in which atoms of
different elements are combined in a molecule"
Empirical formula only indicates atomic ratios but it does not indicate actual number of atoms of different kinds present in the molecule of a compound.
Two or more compound may have same empirical formula.
Empirical formula is determined by experiment.
Molecular Formula
The molecular formula of a compound is defined as:
"The formula of a compound which not only expresses the relative number of atoms of each
kind but also expresses the actual number of atoms of each element present in one molecule".
Molecular formula and empirical formula of a compound are related as:
MOLECULAR FORMULA = (EMPIRICAL FORMULA)n
Where "n" is an integer and is given by:
n = molecular mass of compound / Empirical formula mass of compound
Molecular formula of propane = C3H8.
Molecular formula of sugar = C12H22O11.
Limiting Reactant
The limiting reactant is defined as:
" The substance which produces least amount of products when it is
completely consumed in a chemical reaction".
In stoichiometry when more than one reactant is involved in a chemical reaction, it is not so simple to get actual result of the stoichiometric problem by making relationship between any one of the reactant and product, which are involved in the chemical reaction. As we know that when any one of the reactant is completely used or consumed the reaction is stopped no matter the other reactants are present in very large quantity. This reactant which is totally consumed during the chemical reaction due to which the reaction is stopped is called limiting reactant.
Limiting reactant help us in calculating the actual amount of product formed during the chemical reaction. To understand the concept the limiting reactant consider the following calculation.
Problem
We are provided 50 gm of H2 and 50 gm of N2. Calculate how many gm of NH3 will be formed when the reaction is irreversible.
The equation for the reaction is as follows.
N2 + 3 H2 ----> 2 NH3
Solution
In this problem moles of N2 and H2 are as follows
Moles of N2 = Mass of N2 / Mol. Mass of N2
= 50 / 28
= 1.79
Moles of H2 = Mass of H2 / Mol. Mass of H2
= 50 / 2
= 25
So, the provided moles for the reaction are
nitrogen = 1.79 moles and hydrogen = 25 moles
But in the equation of the process 1 mole of nitrogen require 3 mole of hydrogen. Therefore the provided moles of nitrogen i.e. 1.79 require 1.79 x 3 moles of hydrogen i.e. 5.37 moles although 25 moles of H2 are provided but when nitrogen is consumed the reaction will be stopped and the remaining hydrogen is useless for the reaction so in this problem N2 is a limiting reactant by which we can calculate the actual amount of product formed during the reaction.
N2 + 3 H2 ----> 2 NH3
1 mole of N2 gives 2 moles of NH3
1.79 mole of N2 gives 2 x 1.79 moles of NH3
= 3.58 moles of NH3
Mass of NH3 = Moles of NH3 x Mol. Mass
= 3.58 x 17
= 60.86 gm of NH
Sahm Notes www.icusj.blogspot.com
Define the following terms:
Mole
"The atomic mass or molecular mass of a substance expressed in grams is called mole".
OR
A mole may also be defined as
" the gram atomic mass or gram molecular mass or gram formula mass
of a substance that contains 6.02 x 1023 particles."
OR
"The mass in grams of the atoms or molecules or ions which contains
Avogadro's number of particles i.e. 6.02 x 1023 particles."
For example:
(1) Atomic mass of Carbon = 12 a.m.u.
Therefore 12 gram of carbon = 1 mole of carbon.
(2) molecular mass of nitrogen = 28 a.m.u.
Therefore 28 gram of nitrogen = 1 mole of nitrogen.
(3) Formula mass of NaCl = 58.5 a.m.u.
Therefore 58.5 gram of NaCl = 1 mole of NaCl.
Mole is denoted by "n".
Formula
Number of moles of substance = mass of substance (in grams) / molecular mass or atomic mass or formula mass
Avogadro's Number
One mole of any substance contains equal number of particles (atoms or molecules or ions).Value of this number is 6.02 x 1023. This constant value or number is referred to as "AVOGADRO'S NUMBER"
Example
1 gm mole of Na contain 6.02 x 10(23) atoms of Na.
1 gm mole of Sulphur = 6.02 x 10(23) atoms of Sulphur.
1 gm mole of H2SO4 = 6.02 x 10(23) molecules H2SO4
1 gm mole of H2O = 6.02 x 10(23) molecules of H2O
On the basis of Avogadro's Number "mole" is also defined as
Mass of 6.02 x 10(23) molecules, atoms or ions in gram is called mole.
Determination Of The Number Of Atoms Or Molecules In The Given Mass Of A Substance
Example 1
Calculate the number of atoms in 9.2 gm of Na.
Solution
Atomic mass of Na = 23 a.m.u
If we take 23 gm of Na, it is equal to 1 mole.
23 gm of Na contain 6.02 x 10(23) atoms
1 gm of Na contain 6.02 x 10(23) / 23 atoms
9.2 gm of Na contain 9.2 x 6.02 x 10(23) /23
= 2.408 x 10(23) atoms of Na
Determination Of The Mass Of Given Number Of Atoms Or Molecules Of A Substance
Example 2
Calculate the mass in grams of 3.01 x 10(23) molecules of glucose.
Solution
Molecular mass of glucose = 180 a.m.u
So when we take 180 gm of glucose it is equal to one mole So,
6.02 x 10(23) molecules of glucose = 180 gm
1 molecule of glucose = 180 / 6.02 x 10(23) gm
3.01 x 10(23) molecules of glucose = 3.01 x 10(23) x 180 / 6.02 x 10(23)
= 90 gm
It is denoted by "NA". Sahm Notes www.icusj.blogspot.com
Stoichiometry
Stoichiometry refers to chemical calculations.
Stoichiometry is defined as:
"The quantitative relationship among the reactants and products in a balanced chemical equation".
Assumption of Stoichiometric Calculations
There are two important assumptions for stoichiometric calculations:
(1) Reactants are completely converted into products.
(2) No supplementary or side-reactions occur.
Suppose we want to calculate the mass of CO2 formed when a given mass of carbon burns in air.
The reaction is:
C + O2 CO2
In actual practice, it is possible that we get less amount of CO2 than the calculated mass of CO2.This is because that the given mass of carbon can also form CO in addition to CO2.
2C + O2 2CO
This means that we have to avoid the formation of carbon monoxide.
Types of stoichiometric calculations
Stoichiometric calculations can be divided into three categories.
(1) Mass - Mass Relationship
(2) Mass - Volume Relationship
(3) Volume - Volume Relationship
Mass - Mass Relationship
In this relationship we can determine the unknown mass of a reactant or product from a given mass of teh substance involved in the chemical reaction by using a balanced chemical equation.
Example
Calculate the mass of CO2 that can be obtained by heating 50 gm of limestone.
Solution
Step I - Write a Balanced Equation
CaCO3 ----> CaO + CO2
Step II - Write Down The Molecular Masses And Moles Of Reactant & Product
CaCO3 ----> CaO + CO2
Method I - MOLE METHOD
Number of moles of 50 gm of CaCO3 = 50 / 100 = 0.5 mole
According to equation
1 mole of CaCO3 gives 1 mole of CO2
0.5 mole of CaCO3 will give 0.5 mole of CO2
Mass of CO2 = Moles x Molecular Mass
= 0.5 x 44
= 22 gm
Method II - FACTOR METHOD Sahm Notes www.icusj.blogspot.com
From equation we may write as
100 gm of CaCO3 gives 44 gm of CO2
1 gm of CaCO3 will give 44/100 gm of CO2
50 gm of CaCO3 will give 50 x 44 / 100 gm of CO2
= 22 gm of CO2
Mass - Volume Relationship
The major quantities of gases can be expressed in terms of volume as well as masses. According to Avogardro One gm mole of any gas always occupies 22.4 dm3 volume at S.T.P. So this law is applied in mass-volume relationship.
This relationship is useful in determining the unknown mass or volume of reactant or product by using a given mass or volume of some substance in a chemical reaction.
Example
Calculate the volume of CO2 gas produced at S.T.P by combustion of 20 gm of CH4.
Solution
Step I - Write a Balanced Equation
CH4 + 2 O2 ----> CO2 + 2 H2O
Step II - Write Down The Molecular Masses And Moles Of Reactant & Product
CH4 + 2 O2 ----> CO2 + 2 H2O
Method I - MOLE METHOD
Convert the given mass of CH4 in moles
Number of moles of CH4 = Given Mass of CH4 / Molar Mass of CH4
From Equation
1 mole of CH4 gives 1 moles of CO2
1.25 mole of CH4 will give 1.25 mole of CO2
No. of moles of CO2 obtained = 1.25
But 1 mole of CO2 at S.T.P occupies 22.4 dm3
1.25 mole of CO2 at S.T.P occupies 22.4 x 1.25
= 28 dm3
Method II - FACTOR METHOD Sahm Notes www.icusj.blogspot.com
Molecular mass of CH4 = 16
Molecular mass of CO2 = 44
According to the equation
16 gm of CH4 gives 44 gm of CO2
1 gm of CH4 will give 44/16 gm of CO2
20 gm of CH4 will give 20 x 44/16 gm of CO2
= 55 gm of CO2
44 gm of CO2 at S.T.P occupy a volume 22.4 dm3
1 gm of CO2 at S.T.P occupy a volume 22.4/44 dm3
55 gm of CO2 at S.T.P occupy a volume 55 x 22.4/44
= 28 dm3
Volume - Volume Relationship
This relationship determine the unknown volumes of reactants or products from a known volume of other gas.
This relationship is based on Gay-Lussac's law of combining volume which states that gases react in the ratio of small whole number by volume under similar conditions of temperature & pressure.
Consider this equation
CH4 + 2 O2 ----> CO2 + 2 H2O Sahm Notes www.icusj.blogspot.com
In this reaction one volume of CH4 gas reacts with two volumes of oxygen gas to give one volume of CO2 and two volumes of H2O
Examples
What volume of O2 at S.T.P is required to burn 500 litres (dm3) of C2H4 (ethylene)?
Solution
Step I - Write a Balanced Equation
C2H4 + 3 O2 ----> 2 CO2 + 2 H2O
Step II - Write Down The Moles And Volume Of Reactant & Product
C2H4 + 3 O2 ----> 2 CO2 + 2 H2O
According to Equation
1 dm3 of C2H4 requires 3 dm3 of O2
500 dm3 of C2H4 requires 3 x 500 dm3 of O2
= 1500 dm3 of O2
Gay-Lussac's Law Of Combining Volume
According to Gay-Lussac's Law Of Combining Volume:
"Gases react in the ratio of small whole numbers by volume under
similar conditions of temperature and pressure".
For example:
Nitrogen and hydrogen react to form ammonia gas.
N2 + 3H2 2NH3
In this example one volume of nitrogen gas reacts with three volumes of hydrogen gas to produce two volumes of ammonia gas.
Fill In The Blanks Give Suitable Ans And Fill With Pencil
A Product Of Sham Notes
1. 35 gm of oxygen at S.T.P occupies a volume of __________ dm3.
2. 0.00051 contains __________ significant figures.
3. 0.00054471 contains __________ significant figures.
4. The volume of 1 gm of hydrogen gas at S.T.P is __________.
5. The oxidation number Mn in KMnO4 is __________.
6. The product of ionic concentration in a saturated solution is called __________.
7. 16 gm of oxygen at S.T.P occupies a volume of __________ dm3.
8. The shape of the orbital for which l = 0 is __________.
9. The radius of Cl-1 is __________ than the radius of Cl0.
10. Sp2 hybridization is also known as __________.
11. The value of 1 Debye is __________.
12. The reactions catalyzed by sunlight are called __________.
13. The blue colour of CuSO4 is due to the presence of __________.
14. The force of attraction between the liquid molecules and the surface of container is called __________.
15. The heat of neutralization of a strong acid and a strong base is __________.
16. C º C triple bond is __________. C = C double bond length.
17. The ions having the same electronic configuration are called iso electronic.
18. On heating, if a solid changes directly into vapours without changing into the liquid state, the phenomenon is called __________.
19. Each orbital in an atom can be completely described by __________.
20. In a molecule of alkene, __________ restricts the rotation of the group of atoms at either end of the molecule.
21. Density, refractive index and vapour pressure are __________ properties.
22. The addition of HCl to H2 solution __________ the ionization of H2S.
23. The reaction of cation or anion (or both) with water so as to change its __________ is known as Hydrolysis.
24. A reaction with higher activation energy will start at __________ temperature.
25. 6.02 x 1023 has __________ significant figures.
26. The internal resistance in the flow of liquid is called __________.
27. A catalyst increases the velocity of a reaction but decreases the __________.
Chapter 1
Introduction to Fundamental Concepts
1. 1 mole of a gas at S.T.P occupies a volume of __________.
2. A gas occupying a volume of 22.4 dm3 at S.T.P contains __________ molecules.
3. A formula, which gives the relative number of atoms in the molecule of a compound, is called __________.
4. A formula which gives the actual number of all kinds of atoms present in the molecule of compound is termed as __________.
5. The chemical formula that not only gives the actual number of atoms but also shows the arrangement of different atoms present in the molecule is called __________.
6. Atomic weight or molecular weight expressed in grams is known as __________.
7. 2 moles of H2O contain __________ grams and __________ number of molecules.
8. Any thing that occupies space and has __________ is called matter.
9. Volume of one __________ mole of a gas at S.T.P is 22.4 cubic feet.
10. A ton mole of iron is equal to __________ tons.
11. The force with which the earth attracts a body is called the __________ of the body.
12. A pure substance contains __________ kind of molecules.
13. The smallest indivisible particle of matter is called __________.
14. The atomic number is equal to the number of __________ in nucleus.
15. The atomic mass is the total number of protons and __________ in an atom of the element.
16. The average weight of atoms of an element as compared to the weight of one atom of __________ is called the atomic mass.
17. 1.0007 contains __________ significant figures.
18. The figure 24.75 will be rounded off to __________.
19. __________ means that the readings and measurements obtained in different experiments are very close to each other.
20. __________ means that the results obtained in different experiments are very close to the accepted values.
21. The degree of a measured quantity __________ with increasing number of significant figures in it.
22. The atomic mass of sodium is __________. Sahm Notes www.icusj.blogspot.com
23. The symbolic representation of a molecule of a compound is called __________.
24. Molecular formula of CHCl3 and its Empirical formula is __________.
25. Molecular formula of benzene is C6H6 and its empirical formula is __________.
26. 58.5 is the __________ of NaCl.
27. 4.5 gms of nitrogen will have __________ molecules.
28. 28 gms of nitrogen will have __________ molecules.
29. 2 moles of SO2 is equal to __________ gms.
30. 1000 gms of H2O is equal to __________ moles.
31. The reactions, which proceed in both directions, are called __________.
32. The reactions, which proceed in forward directions only, are called __________ reactions.
33. The __________ reactions are completed after some time.
34. 0.0006 has __________ significant figures
35. 7.40 x 108 has __________ significant figures.
36. 7 x 108 has __________ significant figures.
37. Usually Molecular formula is simple multiple of the __________.
38. 0.1 mole of H2O contains __________ molecules of H2O.
39. Mass of 3.01 x 1022 molecules of CO2 is __________.
40. __________ is the branch of science which deals with the properties, composition and structure of matter.
41. None zero digits are all __________.
42. The integer part of logarithm is called __________.
43. The decimal fraction of logarithm is called __________.
44. __________ is the amount of substance, which contains as many number of particles as there are in 12 gms of Carbon.
45. 6.02 x 1023 is called the __________.
46. The accuracy of measurement depends on the number of __________.
47. __________ is the branch of chemistry that deals with quantitative relationships among the substances undergoing chemical changes.
48. The sum of atomic weights of all the elements present in molecular formula is called the __________.
49. __________ is the sum of atomic weights of the elements represented by the Empirical formula of the compound.
50. Very small and very large quantities are expressed in terms of __________.
51. In rounding off __________ figure is dropped.
52. Mole is the quantity, which has __________ particle of the substance.
53. For three significant figures, 25.55 is rounded off to __________.
54. The S.I unit of a mass is __________.
55. Mass of 6.02 x 1023 molecules of NaCl is __________ gm.
56. 1 mole of NaOH is __________ gm of NaOH.
57. Formula weight is used for __________ substances.
58. The word S.I stands for __________. Sahm Notes www.icusj.blogspot.com
59. 4.5 gms of water will have __________ molecules.
60. 0.0087 has __________ significant figure.