Download Sec 3.1

Sec 3.1
 Science History
 Democritus - ~ 400 BC Greek Philosopher
o 1st particle theory of matter. He called nature’s basic
particle an “atom”, which meant indivisible.
 Aristotle – “Einstein” of the Greeks
o Did not believe in Democritus idea.
o He believed that all matter was continuous.
o Everyone believed him because of his prestige.
Neither option was supported by experimental
evidence, just speculation. For almost 2000 years,
the idea of the “atom” was lost.
Law of Conservation of Mass – we can not create or destroy mass
during ordinary chemical reactions or by physical changes.
Law of Definite Proportions – Compounds are always composed
of a fixed proportion of elements.
ex. NaCl always contains 1 Na and 1 Cl or 39.34% by mass Na
60.66% by mass Cl
Law of Multiple Proportions – the same element may combine to
form more than one compound.
ex. CO (Carbon monoxide) and CO2 (Carbon Dioxide)
O2 (Oxygen) and O3 (Ozone)
Dalton’s Atomic Theory
(In 1808, he explained the last 3 laws)
1. All matter is composed of extremely small particles
2.
3.
4.
5.
called atoms.
Atoms of a given element are identical in size,
mass, and other properties. Atoms of different
elements differ by size, mass, and other properties.
Atoms can not be subdivided, created, or
destroyed.
Atoms of different elements combine in simple
whole # ratios to form chemical compounds.
In chemical reactions, atoms are combined,
separated, or rearranged.
** By relating atoms to the measureable property of mass,
Dalton turned Democritus’ idea into a scientific theory.
Parts of Dalton’s Atomic Theory no longer hold true, but it has been
modified to explain new observations - Modern Atomic Theory
ex. Atoms are divisible into smaller particles (protons, neutrons, e-)
but the conservation of mass still holds true.
ex. Atoms of the same element can have different mass ( C12& C14 )
Sec 3-2
Atom – the smallest particle of an element that retains the chemical
properties of that element.
The proton, neutron, and electron are called subatomic particles.
Joseph Thomson (1897) – discovered the e- when he passed electric
current through a various gases in a cathode ray tube.
Robert Millikan (1909) – showed that the mass of e- was 1/1837
the mass of the hydrogen atom or 9.109 x 10-31 kg and that e- carry
a charge.
Because of these discoveries, 2 inferences were made about the atomic
structure.
1.
Because atoms are electrically neutral, they must contain a (+)
charge to balance the (-) charge of the e- .
2.
Because e- have so much less mass than atoms, atoms must contain
other particles that account for most of the mass.
Ernest Rutherford (1911) – discovered the atomic nucleus. He
bombarded a thin foil with alpha particles.
** If the atom were equal to the size of a football field, the nucleus
would be the size of a marble. (Mostly empty space)
All nuclei (except the simplest type of H) contains protons (p+) and
neutrons (N).
Mass of p+ = 1.673 x 10 -27 kg or 1836 times the mass of e-.
Mass of n = 1.675 x 10 -27 kg
All About Atoms - What are atoms?
Sec 3-3
Atoms of different elements have different # of p+.
Atomic Number = tells us the # of p+ in the nucleus of each atom of
that elements. Also tells us the # of e-.
Ex. How many protons does Ca have?
How many electrons does Ca have?
Ex. What is the element that has an atomic #7?
Nitrogen
The identity of an atom is determined by the # of p+, not eor N
Isotopes – atoms of the same element that have different masses
(because they differ by the number of neutrons).
Isotopes of hydrogen are Protium, Deuterium, and Tritium.
Most elements consists of mixtures of isotopes: Tin has 10 isotopes.
Mass # = p+ + n
We don’t include the e- because their mass is negligible.
p+
n
Mass #
Protium
1
0
1
Deuterium
1
1
2
Tritium
1
2
3
Normally, we name the isotope by its mass #.
ex. Hydrogen – 1 or H - 1
Uranium – 235 or U - 235
How many protons, neutrons, and electrons does Na – 23 have?
23 = p+ + n
23 = 11 + n
n = 12
Another way to write Na – 23 is
How many p+, n, and e- are in
?
How many p+, n, and e- are in
? And what is the name of x?
***Do not confuse the mass # with the average atomic mass on the
periodic table.
The masses for atoms are extremely small.
ex. O atom = 2.657 x 10-29 g
Therefore we use the relative atomic mass, which is relative to the
Carbon – 12 atom.
C-12 has been assigned a mass of exactly 12 atomic mass units or
amu’s. (not kg or grams)
Using this standard, the p+ and n have ≈ masses = 1 amu.
Mass # ≈ atomic mass in amu’s.
Ave atomic mass – is a weighted average of all the isotopes that
occur naturally.
Ex. Your grades are weighted by:
20% Homework
40% Quizzes
40% Tests
Ex. C-12 (98.9%)
C-13 (1.1%)
C-14 (very very small %)
It takes an enormous amount of H2O molecules to get enough for us to
see. Too many for us to count.
However, in chemistry, we need to keep track of all of these molecules,
so chemists relate the # of molecules to its mass. We call this a
“mole”.
1 mole of any substance = 6.022 x 10 23
Avagadro’s Number
Avagadro’s Number was found by experiment. Exactly 12 g of C-12
atoms contained 6.022 x 1023 C-12 atoms or we say 12g/mole.
That is 602 sextillion – an extremely huge #
1 oxygen atom has a mass of 15.999 amu
1 mole of oxygen atoms has a mass of 15.999g or can be written as
15.999 g/mole.
Gram to Mole Conversions
. What is the mass in grams of 3.50 mol
of Cu?
1
2. What is the mass of 2.25 mol of Fe?
3. A chemist produced 11.9 g of Al. How
many moles of Al were produced?
4. How many moles of Cu are in 3.22 g Cu?
Practice problems on pg 83:
top of the page 2-4
bottom of the page 1 & 2
6. How many atoms of Ca are in 4 moles of Ca?
7. How many moles of K are in 4.52 x 1023 atoms of K?
8. How many grams of Ag are in 3.01 x 1023
atoms of Ag?
9. How many grams of C are in 2.25 x 1022
atoms of C?
10. How many atoms are in 2.5 g of Fe?
11. What mass of gold, contains the same
# of atoms as 9.0 g of Aluminum?