LEC:2 LAWS OF CHEMICAL COMBINATIONS
The combination of elements to form compounds is governed by the following five basic laws.
1) Law of Conservation of Mass
by Antoine Lavoisier in 1789 "It states that matter can neither be created nor destroyed."
it means mass before the reaction and after the reaction are same.
2) Law of Definite Proportions/Composition.
by French chemist, Joseph Proust. He stated that "a given compound always contains exactly the same proportion of elements by weight"
it means if we take some sample of compound from different source then weight of all elements constituent in the compound is same in all sample, in simple way we can say that quantity of elements in a compound is same in all cases or sources.
3) Law of Multiple Proportions
This law was proposed by Dalton in 1803. According to this law
"if two elements can combine 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"
For example, hydrogen combines with oxygen to form two compounds, namely, water and hydrogen peroxide.
Hydrogen + Oxygen → Water
Hydrogen + Oxygen → Hydrogen Peroxide
Here, the masses of oxygen (i.e. 16 g and 32 g)
which combine with a fixed mass of hydrogen
(2g) bear a simple ratio, i.e. 16:32 or 1: 2.
4) Gay Lussac’s Law of Gaseous Volumes
This law was given by Gay Lussac in 1808. He observed 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 same temperature and pressure"
Examples
in this reaction
N2(g) + 3H2(g) -----> 2NH3(g)
1 litre of nitrogen gas reacts with 3 litres of hydrogen gas to produce 2 litres of ammonia gas or
1 volume of Nitrogen gas reacts with 3 volume of hydrogen gas to produce 2 volume of ammonia gas.
now consider next example
HCl(g) + NH3(g) ==> NH4Cl(s)
1 volume of hydrogen chloride will react with 1 volume of ammonia to form solid ammonium chloride or
1 mole hydrogen chloride gas combines with 1 mole of ammonia gas to give 1 mole of ammonium chloride solid.
5) Avogadro Law
In 1811, Avogadro proposed that "equal volumes of gases at the same temperature and pressure should contain equal number of molecules""6) Dalton’s Atomic Theory
In 1808, Dalton proposed the following :1. Matter consists of indivisible atoms.
2. All the 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 combine in a fixed ratio.
4. Chemical reactions involve reorganisation of atoms. These are neither created nor destroyed in a chemical reaction.
Dalton’s theory could explain the laws of chemical combination.
Atomic mass, Average mass and molecular mass
Atomic mass: the present system of atomic masses is based on carbon - 12 (Carbon - 12 is one
of the isotopes of carbon) as the standard and has been agreed upon in 1961. In this system, 12C is assigned a mass of exactly 12 atomic mass unit (amu) and masses of all other atoms are given relative to this standard.
"One atomic mass unit is defined as a mass exactly equal to one twelfth the mass of one carbon - 12 atom"
1 amu = (mass of 12C) / 12 = 1.66056×10–24 g
Mass of an atom of hydrogen = 1.6736×10–24 g
Thus, in terms of amu, the mass of hydrogen atom = 1.0078 amu = 1.0080 amu
Today, ‘amu’ has been replaced by ‘u’ which is known as unified mass.
When we use atomic masses of elements in calculations, we actually use average atomic masses of elements which are explained below.
Average atomic mass: Many naturally occurring elements exist as more than one isotope. When we take into account the existence of these isotopes and their relative abundance (per cent occurrence) the average atomic mass of that element can be computed. For example, carbon has the following three isotopes with relative abundances and masses as shown
From the above data, the average atomic mass of carbon will come out to be :
= 12.011 u
Molecular mass: Molecular mass is the sum of atomic masses of the elements present in a molecule from molecular formula.
for example
Molecular mass of water (H2O) = 2 × atomic mass of hydrogen + 1 × atomic mass of oxygen
Formula mass: formula mass is the sum of atomic masses of the element present in a molecules from its empirical formula.
now question arise in mostly student that
what is difference between Molecular formula and Empirical formula?
The molecular formula indicates the type and number of atoms in a molecule. The molecular formula of glucose is C6H12O6, which indicates that one molecule of glucose contains 6 atoms of carbon, 12 atoms of hydrogen, and 6 atoms of oxygen.
The empirical formula is also known as the simplest formula. It is used to indicate the mole ratio of elements present in a compound. The empirical formula of glucose would be CH2O.
The formula mass and molecular mass of water (H2O) are same.
Both molecular mass and formula mass signify the same thing that is the total mass of a substance. While molecular mass is used for covalent compounds, the term formula mass is used for ionic compounds. This is because in covalent compounds the constituent particles are molecules, while in ionic compounds like NaCl the constituent particles are ions. In other words, the constituent particles of both type of compounds are different, hence different terms are used.
in NaCl ions are located in 3D like as shown in fig.
Mole Concept
Atoms and molecules are extremely small in size and their numbers in even a small amount of any substance is really very large. To handle such large numbers, a unit of similar magnitude is required. In SI system, mole (symbol, mol) was introduced as seventh base quantity for the amount of a substance.
"One mole is the amount of a substance that contains as many particles or entities as there are atoms in exactly 12 g (or 0.012 kg) of the 12C isotope."
In 12 g of 12C there are 6.022 x 1023 atoms, This number of entities in 1 mol is so important that it is given a separate name and symbol, known as ‘Avogadro’ constant, denoted by (NA) in honour of Avogadro.
It means in one mole of any elements there are 6.022 x 1023 atoms/molecules/ions.
And "The mass of one mole of a substance in grams is called its molar mass."
To solve mole concept problems we have to use some basic formula in steps like
1) No of moles(n) of any substance by dividing its mass(m) to molar mass(M).
for example if we wants to calculate moles of 64 g of oxygen then
molar mass (M) of oxygen is O2 = 2 x 16 = 32 g
so No of moles (n) of O2 is = 64/32= 2 moles
2) No of atoms/molecules/ions from moles (n) is calculated by multiply No of moles (n) to Avogadro No.
No of Particles = No of moles (n) x Avogadro No (6.022 x 1023 )
for example if we wants to calculate No of molecules in 2 moles of Oxygen then
No of molecules = 2 x 6.022 x 1023 = 12.044 x 1023
3) No of particles from mass of a components then first of all convert mass to moles and then multiply it by 6.022 x 1023 .
for example if we wants to calculate no of molecules in 64 g of Oxygen then
No of moles in 64 g of Oxygen = 2 moles (as we have done in previous example)
No of molecules of Oxygen = 2 x 6.022 x 1023 = 12.044 x 1023 molecules.
4) At Standard temperature and pressure (STP) means 273.15 K (0°C) temperature and 1 bar (105 pascal) pressure, the molar volume (V) of an ideal gas is 24.789 L mole-1 .
it means Volume of 1 mole of ideal gas is 24.789 L.
for example if we consider oxygen as a ideal gas then in 64 g of oxygen there is 2 moles and so
Volume of 2 moles of Oxygen is = 2 x 24.789 = 49.578 L.
Or
in reverse if we have 49.578 L of Oxygen (ideal gas) it means we have 2 moles Oxygen.
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