Download NOTES: ATOMIC THEORY

NOTES: ATOMIC THEORY
400 B.C. --Greek philosopher, Democritus of Abdera, proposed that all elements of matter (earth, wind,
fire, water) must be made up of the same basic particle (tiny and indivisible).
ATOM = from the Greek ‘ATOMOS’, meaning INDIVISIBLE. It is the smallest particle of an element
that retains the properties of that element. Also called the building blocks of matter.
1800’S -- JOHN DALTON (Englishman) proposed the first ATOMIC THEORY.
1. All elements are composed of tiny INDIVISIBLE particles called ATOMS.
2. Atoms of the same element are IDENTICAL. Atoms of any one element are different from the atoms
of any other element.
3. Atoms of different elements can combine in simple whole number ratios to form compounds.
4. Chemical reactions occur when atoms are SEPARATED, JOINED, OR REARRANGED. Atoms of
one element are not changed into atoms of another element (BY A CHEMICAL REACTION).
Dalton’s Law of DEFINITE Proportions: When two elements combine to form ONE compound, the
elements combine in the same proportion by mass.
Ex: C + O2 (limited) --> CO
1n:1n ratio
All molecules of CO are in a 12g:16g ratio by mass of carbon atoms to oxygen atoms.
Dalton’s Law of MULTIPLE Proportions: Two elements can combine in different whole number ratios
(proportion) to form MORE THAN ONE compound; the different masses of one element that combine
with the same mass of the other are in small whole number ratios.
Ex: C + O2 (limited)  CO
1n:1n ratio
12g 32g
28g
12g:16g ratio
C + O2 (excess)  CO2
12g 32g
44g
1n:2n ratio
12g:32g ratio
Revision to Dalton’s Atomic Theory:
1. Atoms are DIVISIBLE into over 200 different sub-atomic particles. FOUR of concern to
chemistry are the ELECTRON, the PROTON, the NEUTRON and QUARKS.
2. Due to the existence of ISOTOPES, atoms of the same element are SIMILAR not identical.
3. During a NUCLEAR REACTION, atoms of one element are changed into atoms of a different
element.
PARTICLE
Electron
SYMBOL CHARGE
0
-1
e- or −1 e
Proton
p+ or
1
+1
Neutron
no or
1
0
Positron €
Up Quark
€Quark
Down
Antineutrino
e+ or +1 e
u
d
Neutrino
ν=
€
€
€
_
+1
1.007 amu
nucleus
n
0
1.009 amu
nucleus
+1
0.00055 amu
e--cloud
.336 amu
.336 amu
0 amu
p+’s/no’s
p+’s/no’s
e--cloud
0 amu
e--cloud
ν
+2/3
-1/3
0
ν
0
0
0
0
0
LOCATION
e--cloud
p
0
ν=
MASS
0.00055 amu
_
€
Subatomic Particles (in order of discovery): ELECTRONS (1897), nucleus (1911), PROTONS (1918),
NEUTRONS (1932),€
POSITRONS (1932), QUARKS (1964), anti-neutrino (1934,1956).
€
©T. M. Chipi
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NOTES: ATOMIC THEORY
(1897) J. J. THOMSON = Discovered the ELECTRON, a negatively charged sub-atomic particle, by
experimenting with cathode rays [a glowing beam which travels from the cathode (-) to the anode (+)];
his experiment is know as the cathode –ray tube experiment. He found that the e- was almost 2000 times
lighter than a hydrogen atom, which was thought to be the lightest component of matter.
(1911, 1918) ERNEST RUTHERFORD = Reasoning, just like Thomson, that since the atom is an
electrically NEUTRAL particle, then there must be a POSITIVE particle inside the atom to cancel the
NEGATIVE charge of the e-. The theory was that e- and p+ were evenly distributed throughout the atom.
He discovered the NUCLEUS of the atom (small dense center with positive charge) in 1911 during his
gold-foil experiment and the PROTON (positively charged particles in the atom) in 1918; he noticed that
when alpha particles were shot into nitrogen gas, his scintillation detectors showed the signatures of hydrogen
nuclei. He determined that the only place this hydrogen could have come from was the nitrogen, and therefore
nitrogen must contain hydrogen nuclei (a proton). He also found that the atom was mostly empty space.
(1911) CARL D. ANDERSON = In 1932, positrons were observed by Carl D. Anderson, although the
existence of positrons was first postulated in 1928 by Paul Dirac as an inevitable consequence of the
Dirac equation. Anderson gave the positron its name; he also unsuccessfully suggested renaming
electrons "negatrons." The positron was the first evidence of antimatter and was discovered by passing
cosmic rays through a gas chamber and a lead plate surrounded by a magnet to distinguish the particles by
bending differently charged particles in different directions. Today, positrons are produced in enormous
numbers in accelerator physics laboratories and used in electron-positron colliders.
(1932) JAMES CHADWICK = Confirmed the existence of the NEUTRON (subatomic particle with no
electrical charge). The neutron has mass nearly that of the proton [about 1 atomic mass unit (amu)].
(1964) MURRAY GELL-MANN (American) = suggested the existence of particles with 1/3 and 2/3
electrical charges; these are the QUARKS.
STRUCTURE OF THE ATOM:
1. NUCLEUS: small, dense center = 99% of mass of the atom.
2. ELECTRON CLOUD: mostly empty space where e- revolve around the nucleus = 99% of
the volume of the atom.
OTHER SUBATOMIC PARTICLES- over 200 found to date. Separated into two families of particles:
1. LEPTONS = Include electrons, muons, tau-mesons, neutrinos. (Affected by weak force)
2. HADRONS = include protons, neutrons, quarks. (Affected by strong force)
QUARKS (1964) = subatomic particle that makes up the protons and the neutrons.
SIX TYPES, OR “FLAVOURS”, OF QUARKS:
1. Up (u)
3. Top (t)
5. Charm (c)
2. Down (d)
4. Bottom (b)
6. Strange (s)
TWO CHARGES, THREE “COLOURS”, OF QUARKS:
1. +2/3  up, top, and charm
2. -1/3  down, bottom, and strange
Proton = up (+2/3) + up (+2/3) + down (-1/3) = +1
Neutron = up (+2/3) + down (-1/3) + down (-1/3) = 0
The colours of quarks are red, green and blue; combinations of colors always end up in white.
There are also subatomic particles that make-up ANTIMATTER, for example the anti-particle of the
electron is the POSITRON (electron with a +1 charge).
**When a particle collides with an antiparticle, both are annihilated, ENERGY is released in the form of
two gamma rays and new particles are formed.
©T. M. Chipi
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NOTES: ATOMIC THEORY
Number of elementary particles:
6 types of quarks (make-up p+ and no)
6 leptons (includes electrons)
24 antiparticles (anti-quarks, positrons, etc.)
8 gluons (mass-less, carry the strong nuclear force)
3 bosons (carry the strong and *weak nuclear forces {*w+, *w-, *z}, mass is 100 times the mass
of protons)
1 photon (no mass/no charge, carries electromagnetic force)
Four Universal Forces (in order of strength, from the strongest to the weakest):
1. Strong nuclear force (holds p+ & no together in nucleus since same charges repel)
2. Weak nuclear force (responsible for nuclear (radioactive) decay)
3. Electromagnetic force (opposite charge attraction between e- and p+)
4. Gravitational force/gravity (attraction between two large bodies)
Atomic number = the number that identifies each element as a distinct and unique element. It represents
the number of protons in the nucleus of an atom.
Example:
carbon (atomic # 6) = 6 p+ in the nucleus of every carbon atom.
gold (atomic # 79) = 79 p+ in the nucleus of every gold atom.
*In a neutral atom, the atomic # also indicates the number of e- around the nucleus.
**All atoms are electrically neutral, therefore, the # of p+ must equal the # of e-.
Ion = a charged particle arising from the loss of one or more e- (characteristic of metals) or from the gain
of one or more e- (characteristic of non-metals).
1. An atom that loses electrons becomes positively charged and is called a cation.
2. An atom that gains electrons becomes negatively charged and is called an anion.
3. An ion may be mono-atomic (contains only one element), poly-atomic (contains two or more
elements creating a group that behaves as ONE unit), or complex (contains central metal ion
surrounded by group of molecules or ions).
Element
Iron
Sulfur
Chemical Symbol
Fe
S
p+
26
16
e¯
26
16
e¯ lost/gained
3 e¯
2 e¯
Sum of Particles
26p+ + 26e¯ – 3e¯ =
16p+ + 16e¯ + 2e¯ =
Electrical Charge
3+
2¯
Symbol
Fe3+
S 2¯
Mass number = the sum of the total number of p+ and no in the nucleus of an atom; it identifies isotopes.
Example:
hydrogen (mass # 1) = (1 p+)+(0 no) = 1
lithium (mass # 7) = (3 p+)+(4 no) = 7
*To determine the number of neutrons in an atom subtract the atomic # from the mass #
mass # - atomic # = # of neutrons in atom
Example:
59 - 27 = 32 neutrons in an atom of Co-59
Atomic Mass Unit (amu) = unit of comparison, based on carbon-12, to determine the relative masses of
atoms. One amu is defined as one-twelfth the mass of a carbon atom containing 6 protons and 6 neutrons.
Isotopes = atoms of the same element that have the same number of protons, but different number of
neutrons. (Remember that to be from the same element, atoms must contain the same number of protons).
Example:
carbon-12 = 6 p+ and 6 no
carbon-13 = 6 p+ and 7 no
carbon-14 = 6 p+ and 8 no
©T. M. Chipi
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NOTES: ATOMIC THEORY
Isotope
C-12
C-13
C-14
p+ no
6 6
6 7
6 8
e6
6
6
Atomic #
6
6
6
Mass #
12
13
14
Atomic Mass
12 amu
13 amu
14 amu
Chemical
Same
Same
Same
12
Isotopic notation = 6 C : the superscript is the mass #, the subscript is the atomic #, the letter is the
chemical symbol (also C-12).
Hydrogen Isotopes:
1
1. €Hydrogen-1, (H-1), ( 1 H ) or Protium has 1 p+, 0 neutrons
2
2. Hydrogen-2, (H-2), ( 1 H ) or Deuterium has 1 p+, 1 neutron (used in nuclear reactors)
3
3. Hydrogen-3, (H-3), ( 1 H ) or Tritium has 1 p+, 2 neutrons (nuclei are unstable)
€
Many isotopes with unstable nuclei are used in the diagnosis and treatment of disease.
€ treatment of goiter, and Co-60 in the treatment of cancer.
e.g.: I-131 is used in the
€
Atomic Mass = the relative weighted average of the masses of the isotopes of an element.
The Facts:
1. The mass of an atom is concentrated in the nucleus.
2. Each p+ and each no weighs approximately one atomic mass unit.
3. Atomic masses are not whole numbers; they are almost always in decimal form (ie. 12.011)
4. Most elements occur as two or more isotopes in nature, therefore, the percent abundance for each
isotope must be taken into account.
5. Finally, the “average mass” of all the isotopes is assigned as the atomic mass of the element.
Calculating the Atomic Mass of an Element: Example 1
63
65
1. The element copper has the two naturally occurring isotopes: 29Cu (copper-63) and 29Cu (copper-65).
2. The relative abundances are 69.1% and 30.9%, respectively.
3. The mass of copper-63 is approximately 63 amu. The mass of copper 65 is approximately 65 amu.
4. Calculate the atomic mass of copper.
€
€
A. Turn percent to decimal dividing by 100.
i.
69.1/100 = .691
ii.
30.9/100 = .309
*(This changes % abundance to fractional abundance)
B. Multiply each fractional abundance times the isotope’s mass.
i.
63 amu x .691 = 43.533 amu
ii.
65 amu x .309 = 20.085 amu
C. Add (fractional) weighted masses to obtain the atomic mass of the element.
i.
43.533 amu + 20.085 amu = 63.618 amu
Calculating the Percent Abundance of each Isotope: Example 2
63
65
1. The element copper has the two naturally occurring isotopes: 29Cu (copper-63) and 29Cu (copper-65).
+
2. Find the number of p and n° in each isotope: 63 (p+ + n°) - 29 p+ = 34 n°; 65 (p+ + n°) - 29 p+ = 36 n°
3.
4.
5.
6.
7.
8.
Multiply the total of each particle times its mass to find the mass of each individual isotope:
a. 29p+ x 1.0 amu = 29 amu
29p+ x€
1.0 amu = 29 amu
€
b. 34n° x 1.0 amu = 34 amu
36n° x 1.0 amu = 36 amu
c. Cu-63 weighs (sum)= 63 u
Cu-65 weighs (sum) = 65 u
If we let the % abundance of Cu-63 be x, then since % is out of 100, the % abundance of Cu-65 is (100-x).
We need the fractional abundance of each isotope. Hence since a whole = 1, if Cu-63 is x, then Cu-65 is (1-x).
[Isotope1 Mass x Fractional Abundance] + [Isotope2 Mass x Fractional Abundance] = Atomic Mass of Element
For Cu: [(63 u)(x)] + [(65 u)(1-x)]= 63.546 u; x = 0.73, 1-x = 0.27
The % abundance of Cu-63 is 0.73 x 100 = 73%, that of Cu-65 is 0.27 x 100 = 27% {Why is this different?}
©T. M. Chipi
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