The Atomic Theory:
In 1911, Rutherford proposed a revolutionary view of the atom. He suggested that the atom consisted of
a small, dense core of
positively charged particles in the center (or nucleus) of the atom, surrounded
by a swirling ring of electrons. The nucleus was
so dense that the alpha particles would bounce off of it, but the
electrons were so tiny, and spread out at such great distances, that the alpha
particles would pass right through this area of the atom. Rutherford 's
atom resembled a tiny solar system with
the positively charged nucleus always at the center and the electrons revolving
around the nucleus.
Interpreting Rutherford 's
Gold Foil Experiment: The
positively charged particles in the nucleus of
the atom were
called protons. Protons carry an equal, but
opposite, charge to electrons, but protons are much larger and
heavier than electrons.
In 1932,
James Chadwick discovered a third type of subatomic particle, which he named the
neutron. Neutrons help stabilize the protons in
the atom's nucleus. Because the nucleus is so tightly
packed together, the positively charged protons would tend to repel each other
normally. Neutrons help to reduce the repulsion between protons and stabilize
the atom's nucleus. Neutrons always reside in the nucleus of atoms and they are
about the same size as protons. However, neutrons do not have any electrical charge; they are
electrically neutral.
Atoms are
electrically neutral because
the number of protons (+
charges) is equal to the number of electrons (-
charges) and thus the two cancel out. As the atom gets larger, the number
of protons increases, and so does the number of electrons (in the neutral state
of the atom).
Each element has
its own distinct line spectra.
To Bohr,
the line spectra phenomenon
showed that atoms could
not emit energy continuously, but
only in very precise quantities (he described the energy emitted as quantized). Because the emitted light was
due to the movement of electrons, Bohr suggested that electrons could
not move continuously in the atom (as Rutherford
had suggested) but only in precise steps. Bohr hypothesized that electrons
occupy specific energy levels.
When an atom is excited, such as during heating, electrons can jump to higher
levels. When the electrons fall back to lower energy levels, precise quanta of
energy are released as specific wavelengths (lines) of light.
Under
Bohr's theory, an electron's energy levels
(also called electron shells) can be imagined as concentric circles around the nucleus. Normally, electrons exist in the ground state, meaning they occupy the lowest
energy level possible (the electron shell closest to the nucleus). When an
electron is excited by adding energy to an atom (for
example, when it is heated), the electron will absorb energy, "jump"
to a higher energy level, and spin in the higher energy level. After a short
time, this electron will spontaneously "fall" back to a lower energy
level, giving off a quantum of light energy.
Key to Bohr's theory was the fact that the electron could only "jump"
and "fall" to precise energy levels, thus emitting a limited spectrum
of light.
Properties of elements repeat periodically. Similar elements form a group.
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I
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II
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III
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IV
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V
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VI
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VII
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VIII
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H
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He
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1
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2
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Li
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Be
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B
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C
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N
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O
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F
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Ne
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3
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4
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5
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6
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7
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8
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9
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10
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Na
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Mg
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Al
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Si
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P
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S
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Cl
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Ar
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11
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12
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13
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14
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15
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16
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17
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18
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K
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Ca
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Ga
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Ge
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As
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Se
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Br
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Kr
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19
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20
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Transition
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31
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32
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33
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34
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35
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36
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Rb
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Sr
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metals
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In
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Sn
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Sb
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Te
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I
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Xe
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37
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38
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49
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50
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51
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52
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53
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54
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Cs
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Ba
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Tl
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Pb
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Bi
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At
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Rn
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55
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56
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81
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82
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83
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84
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85
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86
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Formation
of compounds
All things are made up of tiny particles called atoms.
Helium, neon, argon, krypton are inert gases. They have 2, 8, 8, and 8 electrons
respectively in their outermost shell. If any atom gets eight electrons in the outer shell, it becomes stable. This is the octet rule.
Oxygen has six electrons in the outermost shell. They are
called valence electrons. If however, there are eight electrons in its valence
shell, it becomes stable. The electron affinity of oxygen is more. By losing
electrons into the neighborhood or gaining electrons from the neighborhood, atoms
become stable in form [compounds].
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Sl no
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Name
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Symbol
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configuration
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1
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Hydrogen
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H
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1
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2
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Carbon
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C
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2, 4
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3
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oxygen
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O
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2, 6
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4
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Sodium
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Na
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2, 8,
1
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5
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Chlorine
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Cl
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2, 8,
7
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Atoms can share electrons to form molecules.
Example:- H2, O2, N2, CO2 [covalent bond formation]
H─H, O═O, N≡N, O═C═O
Atoms can donate electrons to form molecules.[transfer of
electrons]
Example:- NaCl , NaOH, CaCO3, NH4OH [
ionic bonds]
Na+
Cl- , Na+ OH- , Ca2+CO32- NH4+ OH-
Sodium donated its outer shell electron to chlorine and both
have now eight electrons in their outer shell and therefore they are stable.
The earth's atmosphere has N2, O2, CO2, and H2O molecules as
gases.
Sea water is H2O with many dissolved compounds like NaCl,
MgCl2, CaCl2, MgSO4 CaSO4 etc. Water is a polar molecule and a good solvent.
When oxides of non-metals dissolve into water, the water
becomes acidic due to an excess of H+ ions in water. [Acid is formed]. The acids
are more reactive.
When hydroxides dissolve in water, the water becomes basic
due to an excess of OH-
ions.
[Base is formed]. The bases are more reactive.
Therefore compounds can form by;
- Only covalent bonds.
- Only ionic bonds.
- By the combination of both, covalent and ionic bonds.
Chemistry is all about understanding the nature of atoms and the formation of bonds.
Enthalpy change:
While forming new bonds, energy is liberated. Therefore bond
formation is exothermic.
(for Carbon to Carbon bond formation, energy is needed from an external source.)
Breaking bonds need energy. Therefore bond breaking is
endothermic.
In a chemical change either energy is given out or energy is
taken in. Energy is conserved.
Hydrogen and oxygen combine to form water. This is an exothermic reaction.
2H2(g) + O2(g) → 2H2O(l). two molecules of hydrogen gas
combine with one molecule of oxygen to form two molecules of water.
Oxidation and
reduction:
Oxidation: - loss of
electrons by any species is oxidation.
4Al (s) + 3O2 (g) → 2Al2O3
(s)
Aluminum is oxidized to Al2O3
Reduction: - a gain of
electrons by any species is reduction.
ZnO (s) + C
(s) →Zn (s) + CO2 (g)
Zinc oxide
is reduced to Zn
The Earth's crust:
According to calculations by F. W. Clarke, a
little more than 47 percent of Earth's crust consists of oxygen. It occurs mainly in the form of
oxides, particularly silica, alumina, iron oxides, lime, magnesia, potash, and soda. Silica functions principally
as an acid, forming silicates, and the most common minerals of igneous rocks
are silicates. From a computation based on 1,672 analyses of all kinds of
rocks, Clarke arrived at the following values for the average percentage
composition: SiO2=59.71; Al2O3=15.41; Fe2O3=2.63;
FeO=3.52; MgO=4.36; CaO=4.90; Na2O=3.55; K2O=2.80; H2O=1.52;
TiO2=0.60; and P2O5=0.22. (The total of these
is 99.22 percent). All other constituents occur in very small quantities,
usually much less than one percent.
The oxides combine in various ways. Some
examples are given below.
§
Potash and soda combine to produce mostly feldspars, but may also produce nepheline,
leucite, and muscovite.
§
Phosphoric acid with lime forms apatite.
§
Titanium dioxide with ferrous oxide gives rise to ilmenite.
§
Magnesia and iron oxides with silica crystallize as olivine or
enstatite, or with alumina and lime form the complex ferromagnesian silicates
(such as the pyroxenes, amphiboles, and biotites).
§
Any silica in excess of that required to neutralize the bases
separates out as quartz; excess alumina crystallizes as corundum.
These combinations must be regarded only as
general tendencies, for there are numerous exceptions to the rules. The
prevalent physical conditions also play a role in the formation of rocks.
Clarke also calculated the relative abundances
of the principal rock-forming minerals and obtained the following results:
apatite=0.6 percent, titanium minerals=1.5 percent, quartz=12.0 percent,
feldspars=59.5 percent, biotite=3.8 percent, hornblende, and pyroxene=16.8
percent, for a total of 94.2 percent. These figures, however, can only be
considered rough approximations
