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Chemistry 12
Ch 4:
Atoms and Elements
Page |1
Chapter 4: Atoms and Elements
Bonus Problems: 27, 29, 33, 35, 45, 51, 53, 67, 69, 73, 83, 87, 93, 95, 107, 109
Early Ideas on Matter:
Philosophers (Chinese- yin/yang; Greeks-earth/wind/fire/water) speculated about
the nature of “stuff” without relying on scientific evidence
Leucippus (fifth century BC) and his student Democritus (460-370 BC) first
suggested the material world when broken down to the extreme would consist of
tiny particles called atomos, meaning indivisible.
Alchemists through the middle ages physically experimented with matter aiming to
create gold from base metals and an elixir for everlasting life.
Englishman Robert Boyle (1627-1691) is generally credited as the first to study the
separate science we call chemistry and the first to perform rigorous experiments.
Antoine Lavoisier (1743-1794) discovered the mass of combustion products exactly
equals the mass of the starting reactants.
Law of Mass Conservation (Law of Conservation of Matter); Mass is
neither created nor destroyed in chemical reactions
Joseph Proust (1754-1826) studied copper carbonate, the two tin oxides, and the
two iron sulfides. He made artificial copper carbonate and compared it to natural
copper carbonate, showing that each had the same proportion of weights between
the three elements involved (Cu, C, O). He showed that no intermediate
indeterminate compounds exist between the two tin oxides or the two iron sulfides.
Law of Definite Proportions (Law of Constant Composition); Elements
combine together in specific proportions. All samples of a given compound,
regardless of their source or how they were prepared have the same
proportions of their constituent elements.
These early ideas led to the foundation steps in atomic theory.
Atomic theories explain the behavior of atoms. We will cover Dalton’s Indivisible
atom, J.J. Thomson’s Plum Pudding model, Rutherford’s Nuclear model of the
atom, the Bohr’s Quantum (orbit) model that mathematically only works for one
electron systems and the Orbital Wave Mechanical model. The first three models
are found in Chapter 4 while the last two are found in Chapters 9.
You are to write a paper with references on these five models due Feb 14th.
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Ch 4:
Atoms and Elements
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Dalton’s Atomic Theory (1808):
1.
2.
3.
Elements are composed of tiny, indivisible particles called atoms.
Atoms of a given element are identical in properties, but atoms of one
element are different from the atoms of all other elements.
Compounds form when atoms of two or more different elements combine in
whole number ratios.
Chemical reactions do not create or destroy atoms, they are just rearranged.
Dalton’s atomic theory led to another scientific law…
Law of Multiple Proportions: When two elements form two different
compounds, the masses of element (B) that combine with 1g of element (A) can
be expressed as a ratio of small whole numbers.
Example: CO(1 g C to 1.33 g O) vs CO2 (1 g C to 2.67 g O)
J. J. Thomson (1856-1940);
By the mid-1800’s new experiments
gave data that was inconsistent with an
indivisible atom. Cathode ray tubes
(CRT) contain very low pressures of a
gas and have high voltage passed
through electrodes on either end.
Experiments with CRT gave radiation
that is negatively charged. The same
negative charged substance that
fluoresced (gave off light) was found
using many different gases.
By 1897, JJ Thomson published a paper that concluded the cathode rays are streams
of negatively charged particles, later known as electrons.
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Ch 4:
Atoms and Elements
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Plum Pudding Model;
This experiment led to a divisible neutral atom
which must have both negative and positive
charges. JJ Thomson called his atomic
theory the Plum Pudding Model of the atom.
A
positive sphere like pudding contains particles
(plums) of negatively charged electrons.
Since the atom is neutral, there must be a
positively charged electric field as well.
Thomson assumed there were no positively
charged particles since none showed up in the
experiment. He predicted he mass of the atom comes from the mass of electrons.
In 1909 Robert Millikin;
Robert Millikin measured the charge of an electron (1.6022 x 10-19 Coulombs)
through an oil drop experiment performed numerous times over 5 tedious years.
Using Thomson’s charge to mass ratio (1.7588 x 108 C/g) the electron mass was
accepted as 9.109 x 10-28g, about 2000 times smaller than a single H atom.
This caused the question: What is the major contributor of an atom’s mass.
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Ch 4:
Atoms and Elements
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Ernest Rutherford (1871-1937):
In 1910 Ernest Rutherford created an experiment to test Thomson’s Plum
Pudding model, the gold foil experiment. The results showed that most of the
heavy positive alpha particles passed right through a thin gold foil. Surprisingly, a
small portion of alpha particles were deflected or even sent back. If Thomson’s
atomic model was correct, this would be similar to a rifle shot through tissue paper,
and no bullet could be deflected.
This led to Rutherford’s Nuclear Model of the atom. The nuclear model has all
the positive charge (protons) densely set in the center (nucleus) and the particles of
electrons spread out in a cloud around the nucleus.
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Neutrons: It did not make sense to have all the positive particles (protons) so close
together in a nucleus, they would repel each other. Additionally, some
mass was missing. One of Rutherford’s students, James Chadwick (18911974), proposed there are neutrons, neutral particles within the nucleus
similar to protons. Neutrons were isolated later in 1932.
Atomic Structure:
What we have so far…
Particle
Charge
Electron
-1
Proton
+1
Neutron
0
Mass (amu)
0.0005486 amu
1.0073 amu
1.0087 amu
Mass (g)
9.109 x 10-28 g
1.673 x 10-24 g
1.675 x 10-24 g
1 amu = 1.66054 x 10-24g
Solve for the inverse of this number:
amu = 1 g
-10
Atoms are extremely tiny with diameters around 1-5 x 10 m:
1 angstrom = 1 x 10-10m
All atoms are surrounded by a cloud of negatively charged electrons (-1).
The nucleus contains almost all the mass of an atom. It is positively charged
and contains protons (+1 ) and neutrons (0 charge)
The nucleus is tiny with a diameter 10,000 times smaller than the atom.
Example: a marble (nucleus) in the center of a large football stadium (atom).
Neutral atoms have the same number of electrons and protons.
Each chemical element (X) has a unique number of protons (atomic
number, Z). The number of protons defines the element.
Ions have more or less electrons than protons.
Cations lose electrons, are positive (metals)
Anions gain electrons, are negative (nonmetals)
The number of neutrons may vary creating various isotopes. Isotopes are
chemically identical.
The protons plus neutrons is the Mass Number (A)
Nuclide symbols indicate particular isotopes and ions. AZ X
Chemistry 12
Ch 4:
Atoms and Elements
Example 1:
Fill in the nuclide symbols chart.
Nuclide symbol,
name
A
ZX
carbon-12
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protons neutrons electrons atomic mass
6
6
carbon-14
6
6
74
Sulfide ion
74
16
Potassium ion
184
18
18
39
Periodic Table: Patterns and the Periodic Law
Development:
1869 Dmitri Mendeleev (Russia) and Lothar Meyer (Germany) classified
known elements (about 65 known at that time) and noted similar physical and
chemical properties were found periodically when arranged by increasing
atomic weight and grouped together by chemical reactivity. Several holes
led to predictions of elements and their properties that were not yet
discovered “eka-aluminum” (Ga) and “eka-silicon”(Ge).
Periodic Law – when the elements are arranged in order of
increasing atomic mass, certain sets of properties recur
periodically
Ordered elements by atomic mass
Put elements with similar properties in the same column
Used pattern to predict properties of undiscovered elements
Where atomic mass order did not fit other properties, he re-ordered
by other properties
Example: Te & I
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1915 Henry Moseley developed the concept of atomic numbers. He
improved the periodic table by ordering the elements by increasing atomic
number. More “holes” were found, which led to the discovery of more
elements and the family of noble gases.
The periodic table gives us a great amount of information in an organized manner.
Vertical columns are called groups or families. If you are aware of the properties of a
couple elements in a group, you can make a good guess at the properties of the other
elements in the same group. Periods are the horizontal rows in the periodic table. Many
patterns can be seen or predicted following periods and groups.
Atomic Weights:
The atomic mass scale is arbitrarily defined by international agreement and is
based a standard isotope carbon-12, defining its mass to be exactly 12 amu.
Weighted average atomic masses take into consideration the natural abundance
of all the isotopes of an atom.
Masses and isotopic abundances are measured by Mass Spectroscopy.
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Mass Spectrum quantifies the results…
The mass spectrum for
zirconium
Isotopes:
The 5 peaks in the mass spectrum
shows that there are 5 isotopes of
zirconium - with relative isotopic masses of
90, 91, 92, 94 and 96 on the 12C scale.
The abundance of the isotopes
In this case, the 5 isotopes (with their relative percentage abundances) are:
zirconium-90 51.5
zirconium-91 11.2
zirconium-92 17.1
zirconium-94 17.4
zirconium-96
2.8
(This simple example rounds off much more than I generally accept.)
Working out the relative atomic mass
Using the equation
Weighted Atomic Mass = (0.515 x 90)+(0.112 x 91)+(0.171 x 92)+(0.174 x 94)+(0.028 x 96)
= 91.3 is the relative atomic mass of zirconium.
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Mass: Simple vs. Weighted Average:
Simple average: add all the numbers and divide by the count
Solve for the simple average...
Given:
12.0 g, 16.0 g, 17.0 g
Weighted average: Mass =
Weighted average takes into consideration the fractional abundance of each
number. Fractional abundance is the decimal form of the percent abundance.
All fractional abundance values add up to a total of one (1.00) so there is no
reason to divide by the count.
Solve for the weighted average…
Given:
12.0 g (80.0%), 16.0 g (15.0%), 17.0 g (5.0 %)
Naturally occurring weighted masses for elements are found on the periodic table:
Atomic mass = isotopic mass x fractional abundance
Atomic mass = (massA x fract. abund.A) + (massB x fract. abund.B) + (….
All the fractional abundance values add up to a total of one (1.00)
Example 2:
Use the weighted average to solve the average atomic mass found in nature for Si
Given the following information on its naturally occurring isotopes…
Keep appropriate significant figures.
28
Si:
27.977 amu
92.21%
29
Si:
28.976 amu
4.70%
30
Si:
29.974 amu
3.09%
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Example 3:
There are two naturally occurring isotopes of chlorine. Calculate the percent abundance of
each isotope given the following information on the masses and given that the naturally
occurring weighted atomic mass of chlorine is 35.453 amu…
35
Cl:
34.9689 amu
(1-x)
37
Cl:
36.9658 amu
(x)
Periodic Table:
Organization:
family/group
period
metals
nonmetals
metalloids/semiconductors,
Groups:
Main Group
Transition Metals
Inner Transition Metals or Actinides and Lanthanides
Alkali Metals
Alkaline Earth Metals
Halogens
Noble Gas
Coinage metals
Others
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Preview of Chapter 5:
Ions:
Valence electrons (duet/octet rule)
Cations
Naming:
Cations with known oxidation state of metal
Group 1A (+1), 2A (+2) , Al and Ga (+3), Zn and Cd (+2), Ag (+1)
Name of ion is identical to the name of the atom for cations
Variable oxidation state of metal
Transition metals and metals below the nonmetal on the right have a
variable oxidation state that must be indicated by Roman Numerals in
parenthesis (this method is what I expect you to learn.
Fe+3, iron (III); Fe+2, iron (II); Cu+1, copper (I); Sn+4, tin (IV)
An alternative method differentiates from the higher oxidation number
and lower oxidation number using the old form of the name and ic or
ous as an ending respectively. (you are not responsible for knowing the
ic and ous ending of metal cations) Fe+3, ferric Fe+2, ferrous; Cu+2,
cupric; Cu+1, cuprous; Sn+4, stannic; Sn+2, stannous
Anions
Naming:
Group VA (-3); VIA (-2), VIIA (-1)
Name of the element root followed by ide.
N-3, nitride; S-2, sulfide, Br-1, bromide
Counting Atoms by Moles:
Avogadro’s number: 6.022 x 1023 particles = 1 mole
Converting atoms to moles
Converting moles to atoms
Molar Mass:
Solving for molar mass of molecules and compounds
O2
H2O
CoBr3