Download Atom - U of L Class Index

Chapter 2 – Atoms and Elements
What is chemistry?
“A branch of science which deals with the elementary
substances or forms of matter, of which all bodies are
composed, the laws that regulate the combination of these
elements in the formation of compound bodies, and 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 ?
Ex) COFFEE
Composition:
i) Organic Compounds:
Proteins
Esters
Acids
Sugars
Caffeine
Pesticides
ii) Inorganic Compounds:
Water
Dissolved Salts
Dissolved minerals
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
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.
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.
Where does chemistry fit?
What is the purpose of modern chemistry?
1) Physical chemistry/chemical physics
Thermodynamics
Kinetics
Spectroscopy
Quantum Mechanics
2) Analytical/Environmental Chemistry
Quantification and identification techniques
Separation Methods
Development of New instrumentation
Forensic/Environmental Chemistry
3) Preparative Chemistry
Synthesis of new organic compounds
Synthesis of new inorganic compounds
Material Science
Pharmaceutical Chemistry
4) Biochemistry/Molecular Biology
Chemical processes of Life
Structure and function of proteins
Exploration of DNA
Rational Drug Design
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
Medieval Times
Religious Society
Communal
Static
God Centered , Hierarchical
No Individual Identity
Cycles of Life and Nature
Preordained unchangeable Order
No Change
“Understanding God’s Creation”
“Moral Teachings From Nature”
Alchemy
Transmutation of lead into gold.
Chemical knowledge of the time
Coal/Peat Fuel
Salt Production
Pigments/Dyes
Poisons/Medicines
Metal work:
- steel
- pewter
- quicksilver
- jewelry
Beer/Wine
Distillation
Fragrances/Extracts
Glass/Ceramics
Enlightenment
Scientific Method
Determinism
Materialism
Mechanistic Thinking
Earth Centered
Individualism
Career Scientist
“ Mechanistic Understanding of the Universe”
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.
Elemental Forms
Molecular Gas
Molecular Liquid
Ex) Bromine is a
liquid composed of
diatomic molecules
Br2.
Ex) Chlorine is a gas
composed of diatomic
molecules Cl2.
H2, N2, O2, F2, and Cl2.
Elemental Forms
Atomic Gases Molecular Solids
Ex)
I2, P4, and S8.
Atomic Liquids
Ex) He, Ne, Ar, Kr,
Xe, and Rn .
Infinite Networks
Ex) Metals
Diamond
Graphite
Ex)
Hg, Ga
Graphite (2-D)
Diamond (3-D)
Compounds
Molecular
Infinite networks – SiO2 Sand, Glass
Ionic
Complex ions are
also possible
such as:
SO32-, NO3-, PO4-
NH4+, H3O+
Properties
Chemical
Is observed by changing a
compound/element into
another compound/element.
Physical
Is observed without changing
a compound/element into
another compound/element
Chemical Reactions
Melting point
Energy of Reaction
Freezing Point
Combustion
Density
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
1896
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
1886using 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
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,
and
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.
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:
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
e.g. 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
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 %
The Mole
It is not practical to work on the scale of individual atoms.
It is necessary to work on the macroscopic scale.
It was found that for 6.0221*1023 atoms for any element
the mass corresponds to the atomic mass in grams.
This number, 6.0221*1023, was named after Amedeo
Avogadro who initialyl proposed the idea.
Ex) 12.00 g of carbon corresponds to 6.02221*1023 carbon atoms
The same is true for molecules.
Ex) CO2 weighs 12.000 + 2*15.999 = 43.998 u
6.02221*1023 molecules of CO2 weighs 43.998 g
“1 mole = 6.0221 × 1023”
Molar Mass
The molar mass of a particle is the mass in grams
of one mole, 6.0221*1023, particles
Ex ) One mole of protons weighs
(6.0221*1023)*(1.67*10-24 g) = 1.01 g
Ex) CO2 has molecular mass of 43.998 u
therefore, it has a molar mass of 43.998 g/mol
Exercise: How many moles of Cl2 are there in 105.7g.
The atomic mass of Cl is 35.46 u
The molecular mass of Cl2 is 2*35.46 u = 70.92 u
Its molar mass is 70.92 g/mol.
# moles Cl2 = mass Cl2/molar mass Cl2
=105.7g/70.92 g/mol = 1.490 moles
Exercise:
What is the mass of carbon dioxide
containing 2.57*1021 atoms of oxygen
CO2 contains two O atoms for every CO2 molecule
# CO2 = (# of O)/2 = 2.57*1021/2 = 1.29*1021
# moles CO2 = (# CO2 molecules )/( 6.0221*1023 molecules/mol)
# moles CO2 = (1.29*1021 molecules)/( 6.0221*1023 molecules/mol)
# moles CO2 = 0.00214 mol
How much does this weigh?
mass CO2 = (# moles CO2)*(molar mass of CO2)
mass CO2 = (0.00214 mol)*(43.998 g/mol) = 0.0916 g
Law of Periodicity
Group
Similar
chemical
properties
“The properties of the elements are
periodic functions of atomic number.”
Period
Repetition of properties
Nonmetals – insulators
not ductile
Metalloids - Semiconductors
Ductile ?
Metals – Conducting, Ductile
Overview of the Elements by Group
Hydrogen (H)
Has properties of groups 1 and
17 but doesn’t belong to either.
Diatomic gas (H2) Unreactive
Group 1: Alkali Metals
(Li, Na, K, Rb, Cs, Fr)
Soft metals that react strongly with
Water and oxygen. (reactivity increases
with atomic mass)
Do not exist in pure form in nature due
to high reactivity
Readily lose 1 electron to make cations
with +1 charge
Group 2:
Alkaline Earth Metals
(Be, Mg, Ca, Sr, Ba, Ra )
Most are metals that react with
water to give X(OH)2 and with
oxygen to give XO.
Reactivity increases with
atomic mass.
Beryllium does not react with
water (Highly toxic)
Do not exist in pure form in
nature due to high reactivity
Readily lose 2 electrons to make
cations with +2 charge
3
Groups
3-12:
Transition
Metals
4
5
6
7
8
9
10
11
12
Scandium
Titanium
Vanadium
Chromium Manganese Iron
Cobalt
Nickel
Copper
Zinc
44.9559
47.88
50.9415
51.9961
58.9332
58.693
63.546
65.39
Sc
Ti
V
21
22
Yttrium
Zirconium Niobium
88.9059
91.224
Y
39
23
Zr
55.847
Mn
25
Fe
26
Co
27
Molybdenum Technetium Ruthenium Rhodium
95.94
(98)
101.07
102.906
Mo
Nb
Tc
Ru
Ni
Cu
Zn
28
29
30
Palladium
Silver
Cadmium
106.42
107.868
112.411
Rh
Pd
Ag
Cd
40
41
42
43
44
45
46
47
48
Hafnium
Tantalum
Tungsten
Rhenium
Osmium
Iridium
Platinum
Gold
Mercury
178.49
180.948
183.85
186.207
190.2
192.22
195.08
196.967
200.59
La-Lu
Hf
Ta
W
Re
Os
Ir
Pt
72
73
74
75
76
77
Rutherfordium
Dubnium
Seaborgium
(263)
Bohrium
Hassium
Meitnerium Darmstadium
(262)
(265)
(266)
(261)
Ac-Lr
Cr
24
92.9064
54.9380
Rf
104
(262)
Db
105
Sg
106
Bh
107
Hs
108
78
Mt
109
(281)
Au
Hg
79
80
(1994)
(272)
(1996)
(285)
111
112
Dt
110
Group 3 metals lose 3 e- to make +3 cations otherwise act like
Group 2 metals
Groups 4-11 are the ‘true’ transition metals in that they lose e- to
form coloured compounds in which the metal atom has a positive
charge
Groups 11 and 12 mimic the behaviour of Groups 1 and 2
respectively but are less reactive
The metals in the middle of the
transition groups are hardest
Ex) group 6 Cr, Mo, W
The metals at the edges are
the softest (Groups 3 and 12)
Ex) Zn and Ga
Readily lose electrons to make
cations
Mercury is the only liquid metal;
the rest are solids
Copper and gold are the only
coloured metals
Gold is the most malleable metal
Silver is the best conductor
Group 13 (B, Al, Ga, In, Tl)
Most are metals;
Boron is a metalloid
All are solids,
Gallium -low m.pt.
Aluminum
- industrially important
- third most abundant
- produces its own
protective layer
Lose 3 electrons to make
cations with +3 charge
Form compounds in a 1:3
ratio with halogens (e.g. BCl3)
Group 14
(C, Si, Ge, Sn, Pb)
Form compounds in a 1:4 ratio
with halogens (e.g. CCl4)
C is a nonmetal;
Si and Ge are metalloids;
Sn and Pb are metals
Si is the 2-nd most abundant
Element which does occur
in pure form but as silicates
(compounds made of silicon
and oxygen) which form rocks,
sand, glass, etc.
CARBON
Carbon exists in several different Allotropes:
1) graphite
2) diamonds
3) fullerenes – Many types
Carbon is the backbone atom of
organic and biological molecules
Group 15: Pnictogens (N, P, As, Sb, Bi)
N and P are nonmetals
As and Sb are metalloids
Bi is somewhat metallic
N is a highly stable diatomic
gas (N2)and the most abundant
element in the atmosphere
P in three allotropes
White, P4 – Fire Bombs
Red and black, polymers
- used in match heads)
Form compounds in a 1:3 ratio
with hydrogen (e.g. NH3)
Group 16: Chalcogens (O, S, Se, Te, Po)
O, S & Se are nonmetals
Te is a metalloid
Po is a metal
O is the most abundant in the earth’s
crust & the second most in the atm.
O exists in two allotropes : O2 and O3
both are very reactive gases
S exists in many allotropes:
S2, S6, S8, etc.
Form compounds in a 1:2 ratio
with hydrogen (e.g. H2O)
Gain 2 electrons to make anions with
-2 charge
Group 17: Halogens (F, Cl, Br, I, At)
Nonmetals that exist as diatomic molecules
(except for astatine which is too unstable to study)
F & Cl are gases
Br is a liquid
Iodine is a solid
colourful
F2 is yellow
Cl2 is yellow-green
Br2 is red-brown
I2 is dark purple
l2
Br2
Cl2
F most reactive known
Form compounds in a 1:1 ratio with hydrogen (e.g. HF)
Gain 1 electron to make anions with -1 charge
Group 18: Noble Gases (He, Ne, Ar, Kr, Xe, Rn)
Inert- unreactive gaseous nonmetals
Exist primarily in
elemental from
Compounds have been
made containing Xe
Ex) XeF2
Helium
low density – Balloons
Low BP - Coolant
Xenon
Used as a probe to study
structure in porous material
Glow when a current passes
through them:
Ex) Neon Lights
Concepts
Dalton’s atomic theory of matter
Elemental Forms
Chemical & physical properties
Subatomic particles (protons, neutrons, electrons)
Models of the atom
Isotopes,
calculating average atomic mass and percent abundance
Atomic number and mass number
Elements (names and symbols)
Avogadro’s number and the mole
Periodic table (groups and periods)