Transcript chemistry chapter 5 notes

Chemistry
Chapter 5
Democritus – 4th century B.C. Greek
philosopher who first proposed that
matter is composed of tiny, indivisible
particles called atoms.
The first atomic model was not proposed
until 2000 years later!
John Dalton, an
English
schoolteacher,
proposed the
following atomic
theory, in 1808.
www.infoscience.fr/histoire/biograph/biograph.php3?Ref=96
Dalton’s Atomic Theory:
1. All elements are composed of tiny,
indivisible particles called atoms.
2. Atoms of the same element are
identical and differ from those of
other elements.
3. Atoms cannot be subdivided, created, nor
destroyed.
4. Different elements combine in simple
whole-number ratios to form chemical
compounds.
5. Chemical reactions occur when atoms are
separated, joined, or rearranged.
Law of conservation of mass – mass is
neither created nor destroyed. (See #3
and 5 above)
Law of definite proportions – different
samples of any pure compound contain the
same elements in the same proportions by
mass.
For instance, both 23g of water and 72,000g of
water will be 11.2% hydrogen and 88.9%
oxygen by mass. (See #4 above)
Law of multiple proportions – if two or
more different compounds are composed of
the same two elements, A and B, then the
ratio of the masses of element B that
combine with a certain mass of element A is
always a ratio of small whole numbers. (See
#4 above)
Examples: NO, NO2, N2O4, N2O5
The Modern Atomic Theory
Not all aspects of Dalton’s theory have been
proven correct:
1) atoms are divisible and
2) a given element can have atoms with
different masses.
The important aspects from Dalton’s theory
are that:
1) all matter is composed of atoms and
2) atoms of any one element differ in
properties from atoms of another element.
An atom is defined as the smallest particle of
an element that retains the properties of that
element.
We now know that atoms are not indivisible.
They can be broken down into even smaller,
more fundamental particles.
3 Subatomic Particles: electron,
proton, and neutron.
Discovery of the Electron
Electrons are negatively charged
subatomic particles.
J. J. Thomson, an
English physicist
discovered them in
1897 by doing
experiments
involving passing an
electric current
through a tube
containing gases at
low pressure.
http://www.sciencemuseum.org.uk/collections/treasures/thomp2.asp
Thomson found that a beam, called a cathode
ray tube, traveled from the cathode (-) to the
anode (+), and that this beam was repelled by
negative electrical charge.
Since he knew opposite charges attract and
like charges repel, he proposed that this
cathode ray was a stream of tiny negatively
charged particles moving at high speed. He
named them electrons.
Further studies by Robert Millikan, an
American scientist, led to the determination
of the properties of the electron in 1909 with
his oil-drop experiment .
Oil-drop experiment
Oil-Drop Experiment:

His experiment involved measuring the charge on
tiny oil drops. The charge of each drop was a
multiple of 1.60x10-19 C. This was the charge of an
individual electron. He used this charge, along with
Thomson’s charge/mass ratio, to determine the
mass of an electron.
1.6x10-19C x
1g
1.76x108C
= 9.11x10-28g
An electron has a charge of -1 and a mass of
1/1840, or 9.109x10-31 kg, the mass of a
hydrogen atom.
Discovery of the Atomic Nucleus
Protons: Scientists knew that atoms are neutral
(have no charge). So, if there are electrons
with negative charge, there must also be
particles with a positive charge. This led to
the discovery of the proton.
Protons have a +1 charge and are 1840
times more massive than the electron.
Neutrons: Sir James
Chadwick
discovered them in
1932.
Neutrons have no
charge, but have
essentially the same
mass as a proton.
http://nobelprize.org/physics/laureates/1935/
Once subatomic particles were discovered,
Dalton’s model of the atom had to be
modified.
The second model of the atom (Thomson’s “Plum
Pudding Model”) proposed that electrons were
evenly distributed throughout an atom filled
uniformly with positively charged material.
In 1911, Ernest Rutherford, a New Zealand
native, tested this model with his gold foil
experiment.
http://www.vanderkroft.net/elements/images/portret/ernest_rutherford2.jpg
Gold Foil Experiment:
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

He shot a beam of massive alpha particles
(He+2) at a very thin sheet of gold foil.
He expected the alpha particles to pass easily
through the foil, with little deflection.
He was shocked to see that even though
most passed through without deflection, a
small fraction was deflected at large angles
and some even bounced straight backwards!
To explain this, he modified Thomson’s
atomic model.
Gold-Foil Experiment
Rutherford’s model of the atom stated that the
atom is mostly empty space with all the
positive charge and almost all of the mass
concentrated in a small region, which he
called the nucleus.
The tiny nucleus is composed of
protons and neutrons.
This is an image of silicon atoms
arranged on a face of a crystal. It is
impossible to "see" atoms this way
using ordinary light. The image was
made by a Scanning Tunneling
Microscope, a device that "feels" the
cloud of electrons that form the outer
surface of atoms.
How small is the nucleus?
If the atom were the size of a football
stadium, the nucleus would be about the
size of a marble …10,000 times smaller!
Properties of Subatomic Particles (Table 3-1, pg. 74)
Symbols
Charge
Electron
e-
-1
0
9.109x10-31
Proton
p+
+1
1
1.673x10-27
Neutron
no
0
1
1.675x10-27
Particle
Mass
Number
Actual Mass (kg)
Atoms of one element differ
from atoms of another
element because they have
different numbers of protons,
neutrons, and electrons.
MORE SUBATOMIC
PARTICLES???!!!
Leptons (elementary particles)
I.
A.
B.
C.
D.
Electron
Mu-meson (muon) – more massive than electrons
Tau-meson (tau) – more massive than electrons
3 types of neutrinos – almost massless
I.
Hadrons (made of quarks)
Mesons of many types
A.
i.
Composed of a quark and an antiquark
Baryons
B.
i.
Composed of 3 quarks of different colors
1.
2.
protons (2 up quarks + 1 down quark)
neutrons (1 up quark and 2 down quarks)
Every particle has an antiparticle. The
antiparticle of the electron is the positron (e+)
Gluons – hold quarks together.
6 “flavors” of quarks
1. up
2. down
3. top (truth)
4. bottom (beauty)
5. strange
6. charm
“Colors” of quarks (charge)
+2/3
-1/3
up, top, charm
down, bottom, strange
Atomic number – the number of protons in
the nucleus of an atom.
The atomic number identifies an element.
1) What is the atomic number of aluminum?
13
2) Which element has 26 protons?
Fe - iron
In a neutral atom, the number of protons
(p+) equals the number of electrons (e-).
3) How many p+ does a neon atom have?
10
4)How many e- does a neon atom have?
10
5)An element with atomic number = 6 has how many
e-?
6
Mass Number – The total number of protons
and neutrons in the nucleus of an atom.
mass number – atomic number = number of
neutrons
6) Oxygen-18 has a mass number of
18. How many protons, neutrons and
electrons does it have?
Atomic # = 8
Therefore, p+ = 8
e- = 8
n0 = 18- 8 =10
Isotopes – Atoms that have the same number
of protons but different numbers of
neutrons.
When writing nuclear chemical symbols for
isotopes, the mass number is written as a
superscript and the atomic number is
written as a subscript before the element
symbol. Ex. 146C. This can also be written
in hyphen notation as carbon-14.
Ex. Hydrogen has three isotopes:
Name:
Hydrogen – 1
Hydrogen – 2
(protium)
Hydrogen – 3
(tritium)
Percent:
99.985%
0.015%
negligible
Symbol:
1H
2H
3H
Atomic #:
1
1
1
Mass #:
1
2
3
p+:
1
1
1
e-:
1
1
1
no:
0
1
2
Atomic Mass is given in units called atomic
mass units.
Atomic Mass Unit(amu) – One amu is
defined as 1/12 the mass of a carbon-12
atom.
An atomic mass unit can also be called a
Dalton.
A carbon-12 atom has 6 p+ and 6 n0, and its
mass is set at 12 amu’s.
But, the atomic masses on the periodic table
are not whole numbers (mainly) because
most elements occur in nature as a mixture
of isotopes. The masses are a weighted
average reflecting the relative abundance of
each isotope.
Average atomic mass – weighted
average mass of the atoms in a naturally
occurring sample of the element.
Average atomic mass = (Percent abundance 1)(Isotopic mass 1) +
(Percent abundance 2)(Isotopic mass 2)+ etc.
*when doing weighted averages, change percents to decimal
form: ex. 0.55 instead of 55%.*
The average atomic mass of an element
depends on both the mass and the
relative abundance of each of the
element’s isotopes.
7) Calculate the atomic mass of bromine. The two
isotopes of bromine have atomic masses and
relative abundance of 78.92 amu (50.69%) and
80.92 (49.31%).
(78.92 amu) X (0.5069) = 40.00 amu
(80.92 amu) X (0.4931) = +39.90 amu
Atomic mass Br = 79.90 amu
8) Naturally occurring iron consists of four isotopes with the
abundances indicated here. From the masses and relative
abundances of these isotopes, calculate the atomic weight
of naturally occurring iron.
Isotope
Isotopic Mass (amu)
Iron-54
53.9396
Iron-56
55.9349
Iron-57
56.9354
Iron-58
57.9333
(53.9396 amu) x (0.0582) = 3.1393
(55.9349 amu) x (0.9166) = 51.2699
(56.9354 amu) x (0.0219) = 1.2469
(57.9333 amu) x (0.0033) =+ 0.1912
Atomic mass of Fe 55.8473
% Abundance
5.82
91.66
2.19
0.33
amu
amu
amu
amu
amu
9) The atomic weight of gallium is 69.72 amu. The masses
of the naturally occurring isotopes are 68.9257 amu and
for 6931Ga and 70.9249 amu for 7131Ga. Calculate the
percent abundance of each isotope.
Let x = fraction of 6931Ga. Then (1- x) = fraction of7131Ga.
x(68.9257 amu) + (1 – x)(70.9249 amu) = 69.72 amu
68.9257x + 70.9249 – 70.9249x = 69.72
-1.9992x = -1.20
x = 0.600
Multiply by 100 to get percentage
x = 0.600 = fraction of 6931Ga; therefore, 60.0% 6931Ga
(1-x) = 0.400 = fraction of 7131Ga; therefore, 40.0%7131Ga
10) The atomic weight of rubidium is 85.4678 amu. The two
naturally occurring isotopes of rubidium have the
following masses: rubidium-85, 84.9118 amu; rubidium87, 86.9092 amu. Calculate the percent of rubidium-85 in
naturally occurring rubidium.
Let x = fraction of rubidium-85. Then (1- x) = fraction of
rubidium-87.
x(84.9118 amu) + (1 – x)(86.9092 amu) = 85.4678 amu
84.9118x + 86.9092 – 86.9092x = 85.4678
-1.9974x = -1.4414
x = 0.722
72.2% rubidium-85
Another reason that masses of atoms are not
whole numbers is called mass defect. When
protons and neutrons come together to form
the nucleus, some of their mass is changed
into energy (binding energy) to hold the
nucleus together. This “lost” mass is called
mass defect. The energy can be calculated
by adding together the masses of the protons,
neutrons, and electrons and subtracting the
actual mass of the atom. The missing mass
goes into Einstein’s special relativity formula,
E = mc2.
The Periodic Table:
A Russian chemist, Dmitri Mendeleev (1834-1907),
constructed the first periodic table by listing the
elements in order of increasing atomic mass and
arranging them according to similarities in their
properties. He was able to predict the physical and
chemical properties of missing elements.
In 1913, Henry Moseley, a British
physicist, perfected the periodic table
by arranging the elements in order of
increasing atomic number instead of
mass. This is the arrangement of our
modern periodic table.
Period – a horizontal row on
the periodic table. There are 7
periods.
Periodic Law – When the
elements are arranged in order
of increasing atomic number,
there is a periodic repetition of
their physical and chemical
properties.
Group or Family – a vertical column
on the periodic table. Elements in the
same vertical column have similar
properties. Each group is identified by
a number and a letter A or B. Some
versions of the periodic table use
numbers 1-18 for the groups rather
than having A and B.
Representative Elements Group 1A, 2A, 3A, 4A, 5A,
6A, 7A, 8A or 0.
Alkali metals – Group 1A (extreme left
column)
Alkaline earth metals – Group 2A
Halogens – Group 7A
Noble (or inert) gases – Group 0 (or 8A)
Group B elements:
Transition metals – Group B elements
in the middle of the periodic table
Inner transition metals (or rare earth
elements) – the bottom two rows of
the periodic table
Classification of elements:
Metals - found on left-hand side of
periodic table
*high electrical conductivity
*high luster
*ductile and malleable
*most elements (~80%) are metals
*all except Hg (mercury) are solids at
room temperature
Nonmetals -found on the right-hand
side of the periodic table
*non-lustrous
*poor conductors of electricity
*may be gases, solids or liquid
(bromine is the only liquid at room
temperature)
Metalloids (or semi-metals) – along
the stair-step between metals and
nonmetals (but not aluminum).
*properties are intermediate
between metals and nonmetals
Nuclear Chemistry
Types of Radiation
1. alpha radiation
-helium nuclei emitted from a radioactive source
-2 protons and 2 neutrons
- +2 charge
- symbol is 
-don't travel far and aren't very penetrating
-stopped by a sheet of paper
-can't penetrate skin cells but dangerous if
ingested
-very common with very heavy radioactive
nuclides
2. Beta radiation - fast moving electron formed by
the decomposition of a neutron of an atom. The
neutron breaks into a proton and an electron. The
proton stays in the nucleus and the electron is
ejected. (net effect: neutron changes into proton)
1
1
0
0 n 1 H  -1e
-much smaller than alpha particles
-symbol is 
-charge is -1
-much more penetrating than alpha particles
-stopped by aluminum foil or thin pieces of wood.
14
14
0
C

N

e
6
7
-1
-Nuclides that have too high of a
neutron/proton ratio tend to undergo
beta particle emission.
3. Gamma radiation
- electromagnetic radiation (high energy)
emitted from a nucleus as it changes from
an excited state to a ground energy state
-often emitted along with alpha or beta
radiation
-symbol is 
-have no mass and no charge
-high energy photon
238
4
234
92 U 2 He  90Th
 2
-the emission of gamma rays is one way
that a nucleus with excess energy (in an
excited nuclear state) can relax to its
ground state
-extremely penetrating, very dangerous
-stopped somewhat by several feet of
concrete or several inches of lead
Fission- splitting a heavy nucleus into two
nuclei with smaller mass numbers.
- used for nuclear energy
Ex.
1
235
142
91
1
0 n  92 U 56 Ba 36 Kr 30 n
-production of neutrons causes a chain reaction
(which must be controlled)
-1 kg of uranium-235 is equivalent to 20,000
tons of dynamite
*Fission in a nuclear reactor is carefully
controlled. Much of the energy is heat. This
energy is used to produce steam and
subsequently, electricity.
*A coolant (usually water) is needed.
*The water (or carbon) also acts as a moderator.
It slows the neutrons down so that they can be
captured by the U-235 fuel.
*Control rods made of cadmium are present to
absorb excess neutrons to slow down the
reaction. They can be raised or lowered into
the reactor core.
Fusion- combining two light nuclei to
form a heavier, more stable nucleus
-stars produce their energy this way
-currently, high temperatures are
necessary in order to initiate fusion
-possible future energy source
Methods of detection:
Geiger counter- uses a gas-filled metal tube to
detect radiation
-primarily detects beta
Scintillation counter-uses a specially coated screen todetect radiation
-can detect all types
Film badges-detect beta and gamma
Radioisotopes in Research and Medicine
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Tracer: Iodine-131 is used to check for thyroid
problems
Tracer: Radioactive barium is used to check for
digestive system problems
Radiation sources: Cobalt-60 and Cesium-137
(among many others) are used as radiation sources
for cancer treatment