Chem 115: Week 2
Dr. Babb
Atomic Theory: History of the Atom
•
Atomic Theory:
1.
Law of Conservation of Mass -During a chemical reaction, mass is
conserved. (A. Lavoisier 1743-1794)
Ex.
2.
experimental observations that led scientists to
postulate the existence of the atom (smallest bit of
an element).
71.8 g iron oxide decomposed to 55.8 g iron and 16.0 g oxygen.
Law of Definite Proportions - In a given compound, the elements are
always combined in the same ratio by
mass. (Joseph Proust 1754-1826)
Ex.
1.668 g molybdenum disulfide decomposed to 1.000 g Mo and 0.6680 g S.
Mass Ratio Mo/S =
5.000 g molybdenum disulfide decomposed to 2.996 g Mo and 2.004 g S.
Mass Ratio Mo/S =
Dalton’s Atomic Theory of Matter
John Dalton (1766-1844) proposed this theory to explain the experimental
observations given by the laws of conservation of mass and definite
proportions.
Postulates:
•
Matter consists of tiny particles called atoms.
•
Atoms are indestructible. In chemical reactions, the atoms are rearranged
but they do not themselves break apart.
•
Atoms of the same element are identical in mass and other properties.
•
Atoms of different elements differ in mass and other properties.
•
Chemical combination of elements to form compounds occurs. However, in
a given compound the atoms of each element are present in a fixed number
ratio.
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Chem 115: Week 2
Dr. Babb
Law of Multiple Proportions
When two elements can form more than one compound, the mass ratios of the
elements in the two compounds occur in small whole number ratios.
Example:
Consider two different compounds of iron and sulfur.
A.
Iron(II) disulfide (pyrite or “fool’s gold”):
5.00 g pyrite decomposed to 2.67 g S and 2.33 g Fe.
Mass Ratio S/Fe =
B.
Iron(II) sulfide:
5.00 g iron(II) sulfide decomposed to 1.82 g S and 3.18 g Fe.
Mass Ratio S/Fe =
C.
Law of Multiple Proportions:
Mass Ratio S/Fe in pyrite:iron(II) disulfide =
In that Dalton’s Atomic Theory predicted the law of multiple proportions,
this helped to give validity and force acceptance of the theory.
Atomic Mass: mass of one atom of an element
Atoms are too light to weigh individually. However, scientists could
determine the relative mass of one atom of an element to another (i.e.
relative atomic mass).
Relative Atomic Mass:
A.
5.000 g molybdenum disulfide decomposes to 2.996 g Mo and 2.004 g S.
Thus, mass ratio Mo/S =
B.
Also known that in each molybdenum disulfide unit there is one atom of
Mo to two atoms of S (MoS2).
Thus, atom ratio of Mo/S =
C.
From mass ratio and atom ratio can find relative atomic mass of Mo:S.
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Chem 115: Week 2
Dr. Babb
Carbon-12 (12C) Atomic Mass Scale
Relative atomic masses were not useful until a standard reference point was
established. Atomic masses of all elements were referenced to the atomic
mass of the most abundant isotope of carbon (12C).
Atomic Mass Reference: Carbon-12 or 12C
1 atom 12C = 12 amu (exactly)
OR
1 amu = 1/12 the mass of an atom of 12C
Designation was arbitrary but gave atomic masses close to whole numbers for most elements.
Example:
If the relative mass of Mo:12C is 7.995, what is the atomic mass of
Mo on the 12C atomic mass scale?
Example:
If the relative mass of Fe:S is 1.74 and the relative mass of Fe:12C is
4.65, what is the atomic mass of S?
Example:
The relative mass of an unknown element to Zn is 0.245, while the
relative mass of Zn:12C is 5.45. What is the atomic mass of the
unknown element and what is the identity of the unknown element?
Atomic masses on 12C atomic mass scale are shown as non-integer
numbers below the elements on the periodic table.
But…..
Why is atomic mass of carbon given as 12.011 amu instead of as 12 amu?
Atomic masses shown on periodic table are average atomic masses taking
into account the different isotopes of each element and their percent abundances.
Isotopes are atoms of the same element but with a different mass. These isotopes
occur in different percentages in nature (percent abundances or isotopic
abundances).
Thus, the third postulate of Dalton’s Atomic Theory (Atoms of the same element
are identical in mass) is NOT strictly true.
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Chem 115: Week 2
Dr. Babb
Calculation of Average Atomic Masses
Example:
It is found that carbon consists of two naturally occurring isotopes
(12C and 13C) with atomic masses and % abundances given below. Calculate the
average atomic mass of carbon.
Isotope Atomic Mass
% Abundance
12C
12 amu
98.89%
13C
13.0034 amu
1.11%
Example:
Naturally occurring boron consists of two isotopes, boron-10 and
boron-11. Given the data shown below, calculate the average atomic mass of boron.
Isotope Atomic Mass
% Abundance
10B
10.0129 amu
19.8%
11B
11.0093 amu
Calculation of % Abundances From
Average Atomic Mass
Example:
A sample of naturally occurring gallium has an average atomic mass
of 69.7 and consists of two isotopes, gallium-69 and gallium-71. Given the information
shown below, calculate the % isotopic abundances of the two isotopes.
Isotope Atomic Mass
69Ga
68.9 amu
71Ga
70.9 amu
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Chem 115: Week 2
Dr. Babb
Subatomic Particles: Particles Within the Atom
Three key experiments helped to elucidate the structure of the atom and the
nature of the subatomic particles.
J.J. Thomson’s Experiments with Cathode Ray Tubes (1897)
•
Voltage applied across metal plates. Cathode ray obtained independent of metal used to make
plates.
•
Cathode ray traveled from negative to positive.
•
Magnet deflected cathode ray. Used amount of deflection and magnetic field strength to calculate
charge to mass ratio of cathode ray.
Atom consists of parts one of which
is the electron.
Electron is negatively charged.
Charge to mass ratio of e-=1.76x108 C/g.
R. Milliken’s Experiments with Oil Droplets (1909)
• Milliken watched how fast oil droplets fell: calculated mass of oil droplets.
• Placed negative charge on oil droplets using X-rays. Applied voltage to plates
and suspended oil droplets in midair.
• From mass of oil droplets and voltage needed to suspend, calculated charge
on each oil droplet. Charge always some whole # multiple of 1.60x10-19 C.
Fundamental charge on electron: 1.60x10-19 C.
From e- charge and Thomson’s charge to mass ratio found
mass of electron: 9.09x10-28 g.
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Chem 115: Week 2
Dr. Babb
Rutherford’s Experiments w/Alpha Particles (positive and 7000xheavier than e-)
• Most of alpha particles went straight through metal foil.
• 1/20,000 alpha particle deflected at large angles (repelled by something positive).
• 1/20,000 alpha particle deflected straight back toward source (hitting something
massive).
Most of atom is empty space.
In center of atom there is a massive, positively charged core called the nucleus.
Overall Picture of the Atom
•
Atom consists of mostly empty space (wherein reside the electrons) with a
small, dense, positively charged core at the center called the nucleus.
Ex.
•
Place a pea at center of astrodome (gives relative size of the nucleus to that of the atom).
Three Subatomic Particles
Particle
Rel. Charge
Electron
-1
Proton
+1
Neutron
0
Mass
9.109x10-28 g (5.486x10-4 amu)
1.673x10-24 g (1.007 amu)
1.675x10-24 g (1.009 amu)
•
Protons and neutrons (nucleons) reside within the nucleus.
•
In neutral atom, # protons = # electrons.
•
Isotopes: atoms of the same element but with different mass
– same # protons (and same atomic number)
– different # neutrons
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Chem 115: Week 2
Dr. Babb
Subatomic Numbers
1.
Atomic Number (Z): gives the # protons in the nucleus
a.
Given as integer # above the element on periodic table.
Ex.
C
Z=?
Ca
Z=?
b.
2.
It is the # protons (or atomic #) that specifies the element.
Ex.
An element has 15 protons in the nucleus. What element is present?
Mass Number (A):
nucleus.
gives the # nucleons (#protons plus neutrons) in the
a. Not given on the periodic table.
b. Mass number can be calculated once # neutrons is known.
A = # p + #n or….
A = Z + #n or…..
#n = A – Z
3.
Examples: Find # of subatomic particles present in boron-10 and boron-11.
Ions: Cations and Anions
Ions are charged species and are formed by gain or loss of
electrons.
Cations: positively charged ions
Anions: negatively charged ions
•
•
•
•
•
•
formed by loss of electrons
# protons > # electrons
In ionic compounds, metals tend to
form cations.
Formed by gain of electrons.
# electrons > # protons
In ionic compounds, non-metals tend
to form anions.
Example: How many of each type of subatomic particle (# e-, # p, and # n) are
present in the following?
27Al+3
32S-2
137Cs+1
An element with A=14,
Z=6, and charge=-4.
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Chem 115: Week 2
Dr. Babb
Periodic Table (Mendeleev 1869)
Arranged such that elements with similar properties fall within the
same group.
Periodic Table Terminology
•
•
Groups:
Periods:
elements that lie within the same column (labeled IA, IIIB, etc.)
elements that lie within the same row (Period 1 7)
•
Metals:
all elements to lower left of dividing line (not including H).
» Have similar properties: conduct heat & electricity, have luster, are malleable & ductile,
and most are solids at RT.
•
Non-metals: elements to upper right of dividing line (includes H).
» Have similar properties: poor conductors of heat & electricity, not malleable or ductile
(are brittle), and can be s, l, or g at RT.
•
Metalloids:
elements bordering dividing line (B, Si, Ge, As, Sb, Te, Po, At)
» Have intermediate properties: conduct electricity but not as well as metals.
•
•
•
Representative/Main Group Elements: elements in A-Groups
Transition Elements:
elements in B-Groups
Inner Transition Elements:
elements in lanthanide and
actinide series (elements in rows
below main body of periodic table).
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Chem 115: Week 2
Dr. Babb
•
Alkali Metals: metals in Group IA (Li, Na, K, Rb, Cs)
» Have similar properties: form +1 cations in ionic cmpds. & oxides are strongly basic.
•
Alkaline Earth Metals: metals in Group IIA (Be, Mg, Ca, Sr, Ba)
» Have similar properties: form +2 cations & oxides are strongly basic but less soluble.
•
Halogens:
non-metals in Group VIIA (F, Cl, Br, I, At)
» Have similar properties: form -1 anions, exist as diatomic molecules in elemental form
at RT (F2, Cl2, Br2, I2), & chlorides have a “salty” taste.
•
Noble Gases:non-metals in Group VIIIA (He, Ne, Ar, Kr, Xe, Rn)
» Have similar properties: not very reactive and do not readily form compounds, & exist as
monatomic gases at RT (He(g), Ne(g), Ar(g), etc.)
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