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Chapter 3 Atoms; The Building Blocks of Matter
I: The Atom: From Philosophical Idea to Scientific Theory
Continuous nature of matter or particle nature of matter
Democritus v Aristotle
atom
A: Foundations of Atomic Theory
In late 1700's accepted that
a) element cannot be broken down further by ordinary chemical means
b) elements combine to form compounds
Some questioned if elements combined in the same ratio for a particular compound.
Chemical Reaction - definition
1790's - new quantitative approach to matter
measured masses of elements and compounds which led to
Law of Conservation of Mass - statement
Law of Definite Composition (Proportions) - statement
Law of Multiple Proportions - statement
Two compounds, same elements in each:
| amt of carbon |
amt of oxygen |
||
| carbon |
1 gram |
1.33 grams |
|
| carbon |
1 gram |
2.66 grams |
Ration of oxygen in the two compounds is 1.33:2.66 or 1:2
B: Dalton's Atomic Theory
1808 John Dalton
Explained the, Law of Multiple Proportions, and Law of Definite Composition.
Postulates of his atomic theory.
Dalton's theory and Law of Conservation of Mass --
see figure 3-2 page 67
Dalton's theory and Law of Definite Proportion --
see figure 3-3 page 67
C: Modern Atomic Theory
Some of Dalton's postulates have been proven to be incorrect.
His third postulate indicated that atoms cannot be subdivided.
His second postulate indicates atoms of the same element have identical mass. Isotopes
His theory has been modified -- not discarded.
Homework: Section 3-1
II: The Structure of the Atom
Atoms consist of two regions: nucleus, containing protons and neutrons, and the electron cloud, containing electrons.
Subatomic particles include: protons, neutrons, electrons.
A: Discovery of the Electron
from experiments dealing with electricity and matter;
passing electricity through gases at low pressure;
cathode ray tubes -- figure 3-4 page 70
1. Cathode Rays and Electrons
surface opposite the cathode glowed;
first thought: due to particles that make up "cathode ray";
ray traveled from cathode to anode;
testing the hypothesis they found:
a) an object placed between the cathode and anode cast a shadow on the glass;
b) a paddle wheel placed on rails between the electrodes rolled along the rails from the cathode toward the anode (figure 3-5 page 71).
Conclusion: ray did exist; ray had sufficient mass to set the paddle wheel in motion.
Further experiments showed:
a) cathode rays were deflected by a magnetic field in the same manner as a wire carrying electric current, which was known to have a negative charge
b) deflected by negative charged objects
Conclusion: particles in ray had a negative charge
Joseph Thomson in 1897: back up conclusions and also measured the charge to mass ratio; ratio was always the same, regardless of the metal used to make the cathode or the nature of the gas inside the cathode-ray tube.
Conclusion: all cathode ray particles are identical and negatively charged.
2. Charge and Mass of the Electron
From Thomson's work it was determined that the electron had
a) a very large charge compared to its
b) very small mass.
In 1909, Robert Millikan determined the mass of the electron (approximately 1/2000 the mass of hydrogen); later determined to be 9.109 x 10-31 kg.
Millikan also determined the negative charge of the electron. Concluded that electrons are the same no matter where they come from and that all matter contains electrons.
The cathode-ray experiments indicated that atoms are divisible.
Two important ideas about atomic structure come out of this work:
a) since atoms are electrically neutral, they must contain a positive charge to balance the negative electrons;
b) because electrons have so much less mass than atoms, atoms must contain other particles that account for the bulk of the mass of the atom.
B: Discovery of the Atomic Nucleus
1911 Ernest Rutherford: gold foil bombarded with alpha particles (Helium nuclei or He+2)
Figure 3-6 page 72
Observed: majority of particles passed right through the foil; 1 in 8000 particles were redirected back toward the source of the alpha particles; some particles were slightly deflected from their straight line path.
Conclusions:
** those particles that were redirected back toward the source experienced a
powerful force and that force must occupy only a small portion of the atom of
gold since so few particles were redirected back;
**force must be due to a densely packed bundle of matter with a positive charge;
**those particles that were slightly deflected came into proximity with this
bundle of matter and since the particles have a positive charge and two positives
repel, they were deflected.
Further pieces to the puzzle:
a) volume of the nucleus is very small compared to the size of the atom; (if nucleus were size of marble, atom would be the size of a football field);
b) electrons surrounded the nucleus like planets around the sun; could not explain motion of electrons around the nucleus.
C: Composition of the Atomic Nucleus
Two particles in all nuclei, except for some hydrogen:
| Particle | Charge | Mass | Relative Mass | Mass number |
| proton | +1 | 1.673 x 10-27kg | 1.007 276 AMU | 1 |
| neutron | none | 1.675 x 10-27 kg | 1.008 665 | 1 |
| electron | -1 | 9.109 x 10-31 | 0.000 5486 | 0 |
Nuclei of different elements have differing number of protons; each element has a characteristic number of protons; the number of protons in the nucleus of the atom determines the identity of that element; that number is called the atomic number.
Table 3-1 page 74
1: Forces in the Nucleus
like charges repel (neg/neg or pos/POs)
unlike charges attract (POs/neg or neg/POs)
two protons close to each other, such as in the nucleus, have a strong attraction to each other
similar attraction exists between neutrons in the nucleus
nuclear forces - definition
C: The Sizes of Atoms
electrons located in what is thought to be an electron cloud
nucleus in center of atom
radius of atom - definition
expressed in picometers -- 1 pm = 10-12 m = 10-10 cm)
Homework: Section 3-2
III: Counting Atoms
Hard to measure atoms individually.
Find out about basic properties of atoms. Use this information to count the number of atoms of an element in a sample whose mass you know.
Mole is unit of measure we will use here and for the remainder of the year.
A: Atomic Number
Atoms of different elements have different number of protons.
Atoms of the same element have the same number of protons -- always.
Atomic Number - definition
Located in each block of periodic table.
Elements on table listed in order of increasing atomic number, left to right.
Atomic number identifies the element.
B: Isotopes
Isotopes - definition
table 3-2 page 76
Most of the elements consist of mixtures of isotopes.
C: Mass Number
mass number - definition
table 3-2 page 76
D: Designating Isotopes
Isotopes of hydrogen have names; most isotopes of an element do not.
Two notations for writing isotopes:
a) hyphen notation: name of the element followed by a hyphen followed by the mass number; e.g. hydrogen-3 for tritium
b) nuclear symbol:
235
U
92
where the superscript represents the mass number and the subscript represents the atomic number.
Mass number minus the atomic number = number of neutrons
Nuclide - definition
Table 3-3 page 77 five different nuclides
Sample problem 3-1 page 77
E: Relative Atomic Masses
masses of atoms measured in grams are very small quantities
use relative atomic masses
relative implies a standard
standard is carbon-12
units for relative masses are atomic mass units (AMU)
AMU - definition
standard of carbon-12 has a relative mass of exactly 12 AMU
all relative atomic masses are determined by comparing the mass of the nuclide with carbon-12.
table 3-4 page 80
isotopes have different physical properties because they have different masses; they have similar chemical properties because they have the same number of electrons which have the same energy
for subatomic particles the relative masses are:
table 3-1 page 74
mass number v relative atomic mass
mass number is the number of protons + neutrons
atomic mass (relative) is the mass of protons + neutrons + electrons relative to carbon-12
F: Average Atomic Masses of Elements
most elements are a mixture of isotopes
Average atomic mass - definition
Calculating a weighted average:
25 marbles x 2.00 g = 50 g
75 marbles x 3.00 g = 225 g
total mass = 275 g
average: 275g / 100 marbles = 2.75 g/marble
Alternate calculation:
mass of each marble x decimal fraction representing its percentage in the mixture
from above 25 marbles is 25% (25/100 x 100) and
75 marbles is 75% (75/100 x 100)
25 % is .25
75% is .75
(2.00 g x 0.25) + (3.00 x 0.75) = 2.75g
1. Calculating Average Atomic Mass
Naturally occurring copper consists of 69.17% copper-63, atomic mass of 62.929 598 AMU, and 30.83% copper-65, atomic mass of 64.927 793 AMU Calculate the average atomic mass.
Given: Copper-63, 69.17%, 62.929 598 AMU
Copper-65, 30.83%, 64.927 793
average atomic mass = ?
(percent) x (atomic mass) + (percent) x (atomic mass)... = average atomic mass
(.6917) x (62.929 598 AMU) + (.3083) x (64.927 793 AMU) = 63.55 AMU
Author rounds atomic mass to two decimal places before using it in a calculation.
G: Relating Mass to Numbers of Atoms
mole & Avogadro's Number & molar mass: relate mass in grams to numbers of atoms
1. The Mole
SI unit for amount of substance
mole - definition
The mole is a counting unit as is the dozen, or the gross.
2. Avogadro's Number
The number of particles in a mole has been experimentally determined.
6.02 x 1023
Thus exactly 12 g of carbon-12 contains 6.02 x 1023 atoms of carbon.
The Avogadro's Number - definition
3. Molar Mass
Molar mass - definition
Numerically equal to the atomic mass for elements.
one mole = molar mass = Avogadro Number of particles
4. Gram/Mole Conversions
molar mass is used as a conversion factor in calculations
Figure 3-11 page 82
Sample Problem 3-2 page 82
Sample Problem 3-3 page 83
5. Conversions with Avogadro's Number
amount in moles ---> number of atoms
number of atoms ---> amount in moles
grams ---> moles
moles ---> grams
Sample problem 3-4 page 84
Sample Problem 3-5 page 84
Homework: Section 3-3
A chemical reaction is the transformation of a substance or substances into one or more new substances.
Law of Conservation of Mass states that mass is neither destroyed nor created during ordinary chemical or physical reactions.
Law of Definite Composition states that a chemical compound contains the same elements in exactly the same proportions by mass regardless of the size of the sample or source of the compound.
Law of Multiple Proportions states that if two or more different compounds are composed of the same two elements, then the ratio of the masses of the second element combined with a certain mass of the first element is always a ratio of small whole numbers.
Nuclear Forces are short-range proton-neutron, proton-proton, and neutron-neutron forces holding the nuclear particles together.
The radius of atom is distance from center of nucleus to outer portion of electron cloud.
It is measured by taking 1/2 the distance between two adjacent nuclei of the same element.
The Atomic Number (Z) of an element is the number of protons in the nucleus of each atom of that element.
Isotopes are atoms of the same element that have the same atomic number but different masses.
Mass number is the total number of protons and neutrons in the nucleus of an isotope.
Nuclide is a general term for any isotope of any element.
One atomic mass unit (AMU) is exactly 1/12 the mass of a carbon-12 atom, or 1.660 540 x 10-27kg.
Average Atomic Mass is the weighted average of the atomic masses of the naturally occurring isotopes of an element.
A mole is the amount of a substance that contains as many particles as there are atoms in exactly 12 g of carbon-12.
Avogadro's Number is the number of particles in exactly one mole of a pure substance.
The mass of one mole of a pure substance is called the molar mass of that substance.