Download mass - U of L Class Index

What is chemistry?
“A branch of science which deals with: i) the elementary
substances or forms of matter, of which all bodies are
composed; ii) the laws that regulate the combination of
these elements in the formation of compound bodies; and
iii) the phenomena that accompany their exposure to
diverse physical conditions”.
Composition
Preparation
Reaction
What is it made of ?
How is it made?
How does it interact with others
or react with its surroundings ?
COFFEE
Composition:
i) Organic Compounds:
Proteins
Esters
Acids
Sugars
Caffeine
Pesticides
ii) Inorganic Compounds:
Water
Dissolved Salts
Dissolved minerals
COFFEE
Preparation
Grown
• Biochemical processes make the organic & biological
compounds
Roasted
• Heat combined with air burns off undesired compounds &
converts some to those that give flavor
• Caffeine is burned off if roasted too long
• Decaffeination
COFFEE
Preparation
Ground
• Pulverization of the bean to increase the surface are to aid
extraction process
• Makes it more vulnerable to oxidation affecting taste & shelf
life
Extraction
• Hot water poured over powder, where all water soluble
compounds dissolve.
• The liquid is separated from the bean
residue by a filtration.
COFFEE
Reaction with Surroundings
Caffeine
• Stimulant – increases heart rate by promoting adrenaline
production
• Diuretic – stimulates urine production
Burns
• When it sits the element exposed to the air the organic
compounds oxidizes causing a bitter taste.
Atomic Theory
Greeks
Atom ( A – not, tomos – to cut)
Plato
Aristotle
- Revelation of truth through logic
- Cosmic order
- Hierarchy of being
Atomic Theory
Greeks
Five perfect shapes
Tetrahedron
Cube
Octahedron
Dodecahedron
Icosahedron
Five elements
Fire
Water
Wind
Earth
Ether
Technology
Steam Engines
Organs
Jewelry
Reinforced Concrete
Enlightenment
Scientific Method
“ Mechanistic Understanding of the Universe”
Determinism
Materialism
Individualism
Mechanistic Thinking
Earth Centered
Career Scientist
Times of Change/Discovery
- French and American Revolution.
- Industrial Revolution
- Rapid exploration of chemistry began:
New Elements
Natural Products
Synthetic Methods
Lavoisier
Joseph Proust
John Dalton
1785 “Conservation of mass”
1794 “Law of Definite Proportions”
1808 “Atomic Theory of Matter”
1. All matter consists of solid and indivisible atoms.
2. All of the atoms of a given chemical element are
identical in mass and in all other properties.
3. Different elements have different kinds of atoms;
these atoms differ in mass from element to element.
4. Atoms are indestructible & retain their identity in all
chemical reactions.
5. The formation of a compound from its elements occurs
through the combination of atoms of unlike elements in
small whole-number ratios.
Modifications Required to Daltons Theory
1. Atoms can be further divided into subatomic particles.
Ex) Protons, neutrons, electrons
2. Different isotopes of an element have different masses
Ex) Carbon-12 12.000 u
Carbon-13 13.003 u
Carbon-14 14.003 u
3. Valid, However some have very similar masses.
Ex) Nitrogen-14 14.003 u.
Carbon-14
14.003 u.
4. In nuclear reactions, atoms do not retain their identity.
Ex) Radium-226 → Radon-222 + a-4
5. Valid, however, Dalton was unaware that not all elements
are made up of single atoms.
Modern Atomic Theory
In the late 19-th and early 20-th century the
basic principles of modern atomic theory were
laid down
Electron
J.J. Thomson
1896
R. A. Millikan
1909
Henri Becquerel
1996
Marie and Paul Currie
1899
Proton/Nucleus
Ernest Rutherford
1919
Neutron
J. Chadwick
1932
Radioactivity
Electrons
Hole drilled
in tube. Gass
entering tube
glows
Cathode Ray Tube
Cathode: negative electrode
Anode: positive electrode
Current flows when tube is evacuated
Cathode Rays
Electron charge-to-mass ratio
J.J. Thomson – 1897 -
cathode rays are negatively
charged particles
CRT with electric
and magnetic
fields applied at
right angles
Beam deflects to
positively charged
plate
Magnetic field applied to deflected beam
Changes in the deflection behaviour allowed the mass to
charge ratio of the electron to be determined at 1.7588202 C/kg
Oil Drop Experiment
R Millikan and H A Fletcher (1909)
Accurate measurement of
the electron charge.
Balanced the force of
gravity with an opposing
electric force
The balancing force
between droplets had
common factor
He surmised that the
charge of a single electron
e = 1.60217646 10-19 C
Applying the charge/mass ratio,
mass of e = 9.1093819 10-31 kg
“Canal Rays” and Protons
e-
e- e
-
+
+
+
Cathode
+
Anode
E Goldstein (1850-1930)
discovered canal rays in
1886 using a “reverse
cathode ray” tube
Those that pass through
the hole (“canal”) can be
analyzed for charge-mass
ratio, which are much
smaller than electron, but
largest for hydrogen
Electrons emitted from the
cathode hit gas molecules
causing ionization into (more)
electrons and leaving positively
charged “ions” which travel to
the cathode - cations.
E. Rutherford determined that
the hydrogen cation is a
fundamental particle, and
named it the proton
Radioactivity
Three types
of radiation:
alpha, a ,
beta, b,
&
gamma, g.
Paul and Marie Currie isolated the radioactive elements Radium
and Polonium. They postulated that their spontaneously emitted
radiation was the result of nuclear disintegration.
Three fundamental types of nuclear radiation were identified
by how they respond to electric fields by E. Rutherford.
Radioactivity: properties
From their charge-mass ratios and other experiments
of these rays were characterized and identified
Alpha particles: He2+ nuclei m = 4 amu q =+2)
Beta particles: electron (e-) (identical to cathode rays)
Gamma rays: high-energy light, with wavelengths shorter
than X-rays
Rutherford experiment
Using alpha particles, he bombarded a very thin foil of gold
and observed deflections using a circular fluorescent screen
The nuclear atom
He tried to prove the plum pudding model of the atom
propose by Thomson, which is composed of electrons
imbedded in a sphere of uniform positive charge.
Rutherford said of the alpha particles
deflected almost straight back.
Deflection angle and frequency were carefully
measured, which led to the conclusions:
1. Most of gold foil is empty space
2. There are small centers of highly-positive
charge
3. Centers have high mass to resist
displacement
4. Size of atom estimated from distance
between centers to be ~10-10 m diameter.
5. Size of centers estimated to be ~10-15 m
diameter
Centers were called the nucleus.
Electrons occupy the volume of the atom outside
the nucleus
Constituents of the atom
In 1920 Rutherford predicted the existence of the neutral
particle with mass equal to that of a proton and electron.
In 1932 Chadwick verified experimentally the existence
of the neutron
Relative mass of carbon defined t be 12 u
The mass spectrometer
Mass spectrometer is a variation on the CRT, developed
by J.J. Thomson, which allows the determination of m/z
ratios of cations.
Cations of differing m/z ratio’s can be selected by adjusting
the magnetic field strength
Average atomic mass
Isotopes are atoms of the same element that differ in
mass due to differences in the number of neutrons
35Cl
has 17 protons
and 18 neutrons
37Cl
has 17 protons
and 20 neutrons
The atomic mass of Chlorine is a weighted
average between the two isotopes as:
Similar to
computing an
average grade
of a class
Atomic Mass = Mass(Cl-35) *frac.(Cl-35) + Mass(Cl-37) *frac.(Cl-37)
= (34.968)*(0.7537) + (36.956)*(0.2463) = 35.46 u
Defining an Element
The atomic mass unit (u) is defined as one twelfth of the
mass of a carbon atom containing six protons, six neutrons
and six electrons:
1 u = 1.661 × 10-24 g
The mass of an atom in u will be approximately equal to the
combined number of protons and neutrons it contains.
Mass number (A) = # protons + # neutrons
mass number
symbol
atomic number
If # p’s = #e’s neutral
12
6
C
If # p’s > # e’s cation
If # p’s < # e’s anion
Atomic number (Z) = # protons
The atomic # determines the identity of the element (optional).
Exercise
Gallium has two naturally occurring isotopes and an
average atomic mass of 69.723 u:
69
71
68.926 u
70.925 u
G
G
Atomic masses
measured
using Mass
Spectrometry
Calculate the percent abundance of each isotope of gallium.
At. Mass = M(69G)*frac(69G) + M(71G)*frac(71G)
frac(69G) + frac(71G) =1
frac(69G) =1- frac(71G) =1-x
At. Mass = M(69G)*(1-x) + M(71G)*x
69.723 = (68.926)*(1-x) + (70.925)*x= 68.926+1.999*x
x =(69.723-68.926)/1.999 = 0.3987 = 39.87 %