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name: ____________________
date: _________
The Atom
History
 by 400 BC, Democritus (Greek) had postulated there must be some basic unit of matter that
could not be divided any further. This basic unit = atom, (Greek – indivisible).
 disagreement over this idea for the next 2200 years, until the invention of the chemical
balance – the tool needed to study composition of pure substances quantitatively
 Watershed event – 1803 – John Dalton postulated his atomic theory:
1) All matter is composed of extremely small particles called atoms.
2) Atoms of a given element are identical in size, mass, and other properties; atoms of
different elements differ in size, mass, and other properties
3) Atoms cannot be subdivided, created, or destroyed
4) Atoms of different elements can combine in simple, whole-numbered ratios to form
chemical compounds.
5) In chemical reactions, atoms are combined, separated, or rearranged.
 we now know there are exceptions to Dalton’s theory (atoms can be split, and elements can
have atoms with different masses, called isotopes)
 still, Dalton’s theory = cornerstone in thinking about chemistry
Atoms and mass
 So, if atoms are indivisible and atoms of different elements can combine in chemical
reactions, then mass must be conserved
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+
+
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name: ____________________
date: _________
The Experiments
1) CRT experiments (Cathode Ray Tube – Sir Joseph John Thompson)  discovery that the
electron is negatively-charged
2) Millikan’s Oil-Drop Experiment  determination of actual mass and charge of the electron
 (show overhead and explain)
3) Rutherford’s gold foil experiments  discovery of the atomic nucleus
 fast-moving positively-charged particles shot through gold foil, with only slight
deflections
 but, 1 in 8000 ricocheted back toward the source
 must be repelled by a very strong force in the middle of the atom that takes up a very
small amount of space  it must also have a positive charge (b/c like charges repel)
 nucleus – the positively charged, dense central portion of the atom that contains nearly
all of its mass, but takes up only an insignificant fraction of its volume
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nucleus
electron path
Atomic Arrangement
 within the nucleus are protons (+) and neutrons (0)
 protons – found in nucleus; positive charge equal to negative charge of an electron
 about 1840x heavier than electrons
 neutrons – electrically neutral particles found in atomic nuclei (just slightly heavier than
protons)
 electrons – negatively-charged subatomic particles that orbit the nucleus of an atom in clouds
proton
neutron
mass
1
1
charge
1
0
electron
1/1840
-1
 atomic number – the number of protons in the nucleus of each atom of that element
 # of protons = # of electrons (since all atoms are electrically neutral)
 mass number – the total number of protons + neutrons in the nucleus of an isotope
 # of neutrons = mass # - atomic #
name: ____________________
date: _________
 isotopes – atoms of the same element that have different masses
 diff. masses due to diff. # of neutrons
 protium (99.985%), deuterium (0.015%), and tritium (very rare, radioactive) – isotopes of
hydrogen
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N
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N
N
Protium
Deuterium
 2 ways to name isotopes:
mass#
1) nuclear symbol: atomic
# symbol =
A
Z
Tritium
E
238
92
U
ex:
2) hyphen notation: element–mass #
ex: uranium–238 or U–238
 so, the three isotopes of hydrogen could be named:
 protium: 11 H or hydrogen – 1
 deuterium:
 tritium:
3
1
2
1
H or hydrogen – 2
H or hydrogen – 3
Mass and the Atom (Part II)
 atomic mass unit (u) – 1/12 the mass of a carbon-12 atom
 relative atomic mass – mass of an atom expressed in atomic mass units (it’s relative to
carbon-12)
 non-integer atomic masses (e.g. – average atomic masses) result from the existence of a
number of different isotopes
 average atomic mass – the weighted average of the masses of the naturally occurring
isotopes of an element
35
37
Cl & 17
Cl
 for example, there are 2 main isotopes of chlorine: 17
35
37
Cl , with a mass of 35u and 25% is 17
Cl , with a mass of 37u; so
 75% of all chlorine is 17
the average atomic mass for chlorine is (.75 x 35u) + (.25 x 37u) = 35.5u
 take carbon as a second example:
carbon-12 = 12.0u
98.90%, in nature
carbon-13 = 13.003355u
1.10%, in nature
 average atomic mass for carbon = 12.011u
name: ____________________
date: _________
The Mole

mole – the amount of a substance that contains the same number of particles as the number of
atoms in 12 g of carbon-12
 = SI unit for amount of a substance
 the mole is a counting unit, like a dozen

Avogadro’s number – the # of particles in one mole of a pure substance
 Avogadro’s # = 6.022 x 1023
 this can mean 6.022 x 1023 atoms, molecules, or formula units
 1 mole of carbon-12 = 6.022 x 1023 atoms of carbon-12 = 12 g of carbon-12

molar mass – the mass (in grams) of one mole of an element
 mass of one mole, in grams = mass in atomic mass units (u)

these 3 concepts can be used in factor label problems to convert from grams to moles to
atoms
 introduce Y chart
 note: 1 mol of any gas = 22.4 L of that gas
problems:
 2.0 mol He  g He (8.0 g) note: sig figs
 2.25 mol Fe  g Fe (126 g) note: sig figs
 11.9 g Al  mol Al (.441 mol Al)
 3.01 x 1023 atoms Ag  mol Ag (0.500 mol Ag)
 2.75 mol Al  atoms Al (1.66 x 1024 atoms)
 4.00 g S  atoms S (7.50 x 1022 atoms)
 What mass of Au contains the same # of atoms as 9.0g Al? (67g Au)