Download Chemistry 1 Lectures

ed. Brad Collins
Atoms and Elements
Chapter 2
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Sunday, August 18, 13
Early Ideas of the atom
• 5th Century BC - Democritus
• All matter consists of very small,
indivisible particles.
• Called them atmos
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Alchemists
• Interested in transmutation of elements
• Mystical aspects - Philosopher's Stone
• Developed many techniques still in use:
Geber, ca. 721-815
• Distillation
• Purification by crystallization
• Some of the first users of the Scientific Method
• Abu Musa Jābir ibn Hayyān (Geber)
• Paracelsus (Philippus Theophrastus Aureolus
Bombastus von Hohenheim)
• "The dose makes the poison"
• 'Father' of toxicology
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Paracelsus, 1492-1541
Law of Conservation of Mass
• Roots in Ancient Greek and Indian (Jain) philosophy
•
•
Empedocles ca.450 BCE "Nothing comes from nothing"
•
Mahavira ca.500 BCE 'The universe and its constituents cannot be destroyed or created'.
Nasīr al-Dīn al-Tūsī ca.1250 AD
•
"A body of matter cannot disappear completely. It only changes its form, condition,
composition, color and other properties and turns into a different complex or elementary
matter".
•
Mikhail Lomonosov (1748) showed conservation so mass experimentally
•
Antoine Lavoisier (1774) "Matter cannot be created or destroyed in a chemical reaction".
•
•
Sunday, August 18, 13
Mercury (II) oxide (solid) + heat ---> Mercury (liquid) + Oxygen (gas)
50.0 g
46.3 g
3.7 g
Law of Definite Proportions
• 1799 - Joseph Proust
• Different samples of the same compound always contain its constituent
elements in the same proportion by mass.
• Lead sulfide combines in a ratio of 6.5 g lead to 1 g sulfur.
• Doesn't matter how much of either you start with: 10 g of lead will react
with 1 g of sulfur, leaving 3.5 g of unreacted lead behind.
• Works because atoms of a particular element have the same mass.
• Lead (Pb) 207.2, Sulfur (S) 32.06
• Lead sulfide contains one lead atom and one sulfur atom,
• 207.2:32.06 (6.46:1)
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Law of Multiple Proportions
• Elements can combine in more than one fixed mass ratio to form
different compounds
• Example: carbon and oxygen
• Can combine in a 3:4 ratio and 3:8 ratio
• 3 g carbon + 4 g oxygen make 7 g carbon monoxide
• 3 g carbon + 8 g oxygen make 11 g carbon dioxide
• Note that ratio of the oxygen masses is 2:1
• Law of multiple proportions says that ratio of elements in multiple
compounds will be simple whole numbers
Sunday, August 18, 13
Dalton’s Atomic Theory (1808)
1. Elements are composed of extremely small particles called
atoms.
2. All atoms of a given element are identical, having the same
size, mass and chemical properties. The atoms of one
element are different from the atoms of all other elements.
3. Compounds are composed of atoms of more than one
element. The relative number of atoms of each element in a
given compound is always the same (fixed ratio).
4. Chemical reactions involve only the rearrangement of the
atoms. Atoms are not created or destroyed in the chemical
reactions.
The idea that atoms are neither created nor destroyed in
chemical reactions is known as the Law of the Conservation
of Mass.
2.1
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Application of Dalton’s Laws
• Dalton's Laws helped to explain several phenomena
associated with chemical compounds.
• The Law of Definite Proportions:
• When atoms combine to form compounds they combine
in a fixed ratio.
• Atoms of a given element all have the same weight.
Water, 2H + 1O
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Nitrogen, 2N
Ammonia, 3H + 1N
Application of Dalton’s Laws
• The Law of Multiple Proportions
• Definition: If two elements can combine to form more than
one compound, the masses of one element that combine
with a fixed mass of the other element are in ratios of small
whole numbers.
• Why?
• Different compounds made up of the same elements differ in the
number of atoms of each kind that combine.
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Law of Multiple Proportions
Mass
16
12
Mass
32
12
2.1
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Law of Conservation of Mass
16 X
+
8Y
8 X2Y
2.1
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Atomic Structure
Definition: Atom - the basic unit of an element
that can enter into chemical combination.
Dalton believed atoms were indivisible.
Experiments beginning ca. 1850 showed atoms
have internal structure.
Radiation: The emission and transmission of
energy through space in the form of waves.
2.3
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(Uranium compound)
Radioactivity: the spontaneous emission of particles and/or radiation
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2.3
J.J. Thomson, measured mass/charge of e(1906 Nobel Prize in Physics)
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2.3
Cathode Ray Tube
2.3
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Measured mass of e(1923 Nobel Prize in Physics)
Miliken’s experiment: e- charge = -1.602 x 10-22 C
J.J. Thomson’s charge/mass of e- = -1.76 x 108 C/g
e- mass = 9.10 x 10-28 g
2.3
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2.3
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(1908 Nobel Prize in Chemistry)
α particle velocity ~ 1.4 x 107 m/s
(~5% speed of light)
1.
2.
3.
4.
Atoms positive charge is concentrated in the nucleus
Proton (p) has opposite (+) charge of electron (–) = +1.602 x 10–22 C
Mass of p is 1.67 x 10-24 g, which is 1840 x mass of e–
Relative charge of proton is +1; electron is –1
2.3
Sunday, August 18, 13
Rutherford’s Atomic Model
Nucleus
Atom’s Positive Charge
Electron Cloud
Atom’s Mass
Question: If all the protons are contained in the nucleus, why
don’t they repel each other?
2.3
Sunday, August 18, 13
An ion is an atom, or group of atoms, that has a net
positive or negative charge.
cation – ion with a positive charge
If a neutral atom loses one or more electrons
it becomes a cation.
Na
11 protons
11 electrons
Na+
11 protons
10 electrons
anion – ion with a negative charge
If a neutral atom gains one or more electrons
it becomes an anion.
Cl
17 protons
17 electrons
Cl-
17 protons
18 electrons
2.3
Sunday, August 18, 13
Neutrons
Rutherford hypothesized that a neutral third particle must exist.
Observations:
H atoms have 1 p; He atoms have 2 p
Mass He/Mass H should = 2
BUT: measured mass He/mass H = 4
Walter Bothe and Herbert Becker (1930):
Bombarded beryllium with alpha particles gave off electrically neutral radiation
Gamma ray?
Frederic Joliot-Curie (1932):
Experiments with Bothe-Becker radiation showed that if it fell on paraffin, it could eject protons with very
high energy.
Could not be a gamma ray - photons have no mass
James Chadwick (1932):
Showed that Bothe-Becker radiation could eject protons when it fell on substances other than paraffin.
Energies of the protons ejected could be accounted for, if mass of neutral particle = mass of proton
2.3
Sunday, August 18, 13
Neutrons
Neutrons (n) are neutral (charge = 0)
n mass ~ p mass = 1.67 x 10-24 g = 1 amu*
e– mass = 5.45 x 10–4 amu
*amu = Atomic Mass Unit or Dalton (Da)
2.3
3.1
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Modern Model of
the Atom
atomic radius ~ 100 pm = 1 x 10-10 m
nuclear radius ~ 5 x 10-3 pm = 5 x 10-15 m
“If the atom is the Houston Astrodome, then the
nucleus is a marble on the 50-yard line.”
2.3
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Summary: Subatomic Particle Properties
Particle
Mass (g)
Mass (amu)
Charge (C)
Charge Unit
Electron
9.10939 x 10–28
5.45 x 10–4
–1.6022 x 10–19
–1
Proton
1.67262 x 10–24
1
+1.6022 x 10–19
+1
Neutron
1.67493 x 10–24
1
0
0
2.3
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Atomic number (Z) = number of protons in nucleus
Mass number (A) = number of protons + number of neutrons
= atomic number (Z) + number of neutrons
Isotopes are atoms of the same element (X) with different
numbers of neutrons in their nuclei
Mass Number
A
X
Z
Atomic Number
1
1H
235
92
2
1H
U
Element Symbol
(D)
238
92
3
1H
(T)
U
2.3
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2.3
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Do You Understand Isotopes?
How many protons, neutrons, and electrons are in
14
C
?
6
6 protons, 8 (14 - 6) neutrons, 6 electrons
How many protons, neutrons, and electrons are
11
in 6C ?
6 protons, 5 (11 - 6) neutrons, 6 electrons
2.3
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Development of the Periodic Table
• Dmitri Mendeleev (1834-1908)
• Saw a pattern of chemical
properties in the known elements.
• Na similar to Li
• Mg similar to Be
• Made a set of cards that could be
arranged based on this pattern
(periodicity)
• Predicted that gaps in his table
would be occupied by as yet
undiscovered elements
• Gallium discovered in 1869,
matched Mendeleev’s predicted
properties
2.4
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Group
Period
2.4
Sunday, August 18, 13
Other Forms of the Periodic
Table
2.4
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Atomic Mass
• The weighted average of masses of the naturally occurring isotopes
of an element.
• Carbon-12 is defined as 12.00 amu
• All other elements masses are relative to carbon-12
• Chlorine has 2 naturally occurring isotopes:
• chlorine-35 (mass 34.97 amu) 75.77%
• chlorine-37 (mass 36.97 amu) 24.23%
• Atomic mass of chlorine = (34.97)(0.7577) + (36.97)(0.2433) = 35.45
2.4
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Periodic Table
B Group:
A Group:
A Group:
Main group elements
Representative elements
Transition elements
Main group elements
Transition metals
Representative elements
Inner Transition Elements:
Sometimes called ‘Rare earth elements’
2.4
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Periodic Table
Alkaline earth metals
Halogens
Noble Gases
Alkali metals
2.4
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Periodic Table
Metals:
Good conductors of heat and electricity
Malleable, ductile
High luster solids (except mercury)
Lose electrons easily
Nonmetals:
Poor conductors of heat and electricity
Brittle
May be gases, liquids or solids
Tend to gain electrons
Metaloids:
Properties between metals and nonmetals e.g. fair conductor, but brittle.
2.4
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Elements
•
118 Known elements
•
Elements are represented by a 1 or 2 character symbol.
•
•
Sometimes based on English names:
•
Hydrogen (H), Carbon (C)
•
Helium (He), Cobalt (Co)
Sometimes based on another language (e.g., Latin)
•
Copper (Cu), Latin, Cuprum
•
Tin (Sn) Latin, Stannum
•
Mercury (Hg), Greek, Hydragyrum
•
Tungsten (W), Sweedish, Wolfram
•
Most elements are monatomic (exist as individual atoms)
•
Some elements are diatomic (2-atom pairs): H2, N2, O2
•
A few elements are polyatomic (groups of 3 or more atoms): sulfur (S8), phosphorus (P4)
2.4
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