Video 3-1 Foundations of Atomic Theory Development of Atomic Models Forces in the Nucleus.
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Transcript of Video 3-1 Foundations of Atomic Theory Development of Atomic Models Forces in the Nucleus.
Video 3-1
• Foundations of Atomic Theory• Development of Atomic Models• Forces in the Nucleus
Chapter 3
Atoms: The Building Blocks of Matter
I. Foundations of Atomic Theory
Several basic laws were introduced after the 1790’s (emphasis on quantitative analysis):
Law of conservation of mass: mass is neither created nor destroyed during ordinary chemical or physical processes.
I. Foundations of Atomic Theory
Law of definite proportions: chemical compounds contain the same elements in exactly the same proportions by mass regardless of the size of the sample.
Ex. NaCl always is composed of 39.34% sodium and 60.66% chlorine by mass.
I. Foundations of Atomic Theory
Law of multiple proportions: if two or more different compounds are composed of the same 2 elements, the ratio of mass of the second element combined with a certain mass of the first is always a ratio of small whole numbers.
I. Foundations of Atomic Theory
Ex. CO and CO2: For the same mass of carbon, the mass of the O in CO to the mass of O in CO2 will be 1:2
If you had 28 g of CO and 44 g of CO2, both would contain 12 g of C. The CO would contain 16 g of O and the CO2 would contain 32 g of O.
I. Foundations of Atomic Theory
Masses in CO Masses in CO2 Ratios
12 g C16 g O
12 g C32 g O
C= 12:12 = 1:1
O = 16: 32 = 1:2
II. Development of Atomic Models:
John Dalton (1808):1. All matter is composed of extremely
small particles called atoms (cannot be subdivided, created, nor destroyed)
2. Atoms of the same element are identical; atoms of different elements are different
II. Development of Atomic Models:
John Dalton (1808):3. Atoms combine in simple whole
number ratios to form compounds4. In chemical reactions, atoms
combine, separate, or are rearranged.
II. Development of Atomic Models:
John Dalton (1808):Which of these were later proven wrong
and why?
1) Atoms can be subdivided 2) Atoms of the same element do
NOT have to be identical
II. Development of Atomic Models:
J.J. Thomson (1897) and Robert Millikan (1909):
Used cathode rays to determine that atoms contained small negatively charged particles called electrons.
Atoms must also contain positive charges to balance the negative electrons
II. Development of Atomic Models:
II. Development of Atomic Models:
Other particles must account for most of the mass of the atom
Millikan determine the size of the charge on the electron (oil drop experiment)
II. Development of Atomic Models:
Ernest Rutherford (1911): What was the structure of the atom? Gold Foil Experiment
Thomson assumed mass and charged particles were evenly distributed throughout the atom (“plum-pudding” model)
II. Development of Atomic Models:
II. Development of Atomic Models:
Ernest Rutherford (1911): Expected most of the particles to pass
with only slight deflection Most particles did, but some showed
wide-angle deflections (some almost came back to the source).
II. Development of Atomic Models:
II. Development of Atomic Models:
Ernest Rutherford (1911): discovery of the NUCLEUS of the
atom small, dense, positively charged center
of the atom number of PROTONS in the nucleus
determines the atom’s identity
II. Development of Atomic Models:
Rutherford Atomic Model (solar system model)
III. Forces in the Nucleus
Repulsive forces should exist between protons in the nucleus (like charges repel).
Why doesn’t the nucleus “fly apart” due to the repulsive electromagnetic force?
III. Forces in the Nucleus
Strong (nuclear) force: attractive force that acts over very
small distances in the nucleus causes proton-proton, proton-
neutron, neutron-neutron attractions
Note: gravitational force is present, but negligible. Why?
Video 3-2
• Atomic Dimensions• Properties of Atoms and Ions• Designating Isotopes• Elements on the Periodic Table• Average Atomic Mass
IV. Atomic Dimensions
How “big” are subatomic particles?
Particle Symbol Relative Charge
Mass Number
Relative Mass (amu)
Actual Mass (kg)
electron e- -1 0 0.0005486 9.109 x
10-31
proton p+ +1 1 1.007276 1.673 x
10-27
neutron n0 0 1 1.008665 1.675 x 10-27
e01
p11
n10
IV. Atomic Dimensions
How “big” are subatomic particles? Atomic radii: 40 to 270 pm Nuclear radii: about 0.001 pm Nuclear density: about 2 x 108 metric
tons/cm3
1 amu (atomic mass unit) = 1.660540 x 10-
27 kg Why?
V. Properties of Atoms and Ions atomic number (Z): number of protons in
an atom mass number: number of protons +
neutrons in an atom (number of nucleons—particles in the nucleus)
isotopes (nuclides): atoms of the same element that have different masses (different number of NEUTRONS)
V. Properties of Atoms and Ions ions: atoms with a charge (protons
electrons) charge = protons – electrons atoms can only turn into ions by gaining or
losing ELECTRONS cation: positive ion anion: negative ion
VI. Designating Isotopes There are two ways to write symbols for an
isotope1. name-(mass number)2. symbol
massnumbereratomicnumb
VI. Designating Isotopes
Examples: Hydrogen has 3 isotopes:
Protium Deuterium Tritium
How many neutrons in each isotope? Note: mass number – atomic number =
number of neutrons
Hydrogen-1 H11H21H31
Hydrogen-2
Hydrogen-3
VII. Elements on the Periodic Table
every periodic table will give you at least 3 pieces of information about elements:
Li
3
6.941
Atomic Number
Symbol
Atomic mass (amu)
VII. Elements on the Periodic Table
What is the basis for the atomic mass unit (amu)?
1 amu = exactly 1/12 the mass of a carbon-12 atom (6 protons, 6 neutrons, 6 electrons)
All other atomic masses are based on comparisons to C-12 (exactly 12 amu).
VII. Elements on the Periodic Table
Example:C-13 has a mass that is 1.083613 times
heavier than C-12. The mass of C-13 is
(1.083613) x 12 amu = 13.003356 amu
VIII. Average Atomic Mass Carbon has 3 isotopes (nuclides):
C-12 (12 amu) C-13 (13.003 amu) C-14 (14.003 amu)
Their average mass should be(12 + 13.003 + 14.003) / 3= 13.002 amu
VIII. Average Atomic Mass The atomic mass of carbon (periodic table)
is 12.011 amu. WHY?
VIII. Average Atomic Mass Atomic Mass of an element: weighted
average of all the atoms in a naturally occurring sample of that element (NOTE: not every atom in that sample has the same mass)
Ex. How would you determine the average age of the students in this class?
VIII. Average Atomic Mass atomic mass = sum of (mass of each
isotope x percent abundance)
VIII. Average Atomic Mass Example:
C-12 (12 amu) 98.90% C-13 (13.003 amu) 1.10% C-14 (14.003 amu) trace
atomic mass of C = (12 amu)(0.9890) + (13.003 amu)(0.0110) + (14.003 amu)(0)
= 12.011 amu
VIII. Average Atomic Mass NOTE: the atomic mass of most elements
will usually give you an idea of the most common isotope of that element (mass number that is closest to the atomic mass)
Video 3-3
• Counting Atoms and Stuff
IX. Counting Atoms and Stuff 1 mole = 6.0221367 x 1023 things Ex. 1 mole of eggs contains 6.0221367 x
1023 eggs
6.0221367 x 1023 = Avogadro’s number (N)
1 mol = 6.022 x 1023 particles
IX. Counting Atoms and Stuff 1 mole = 6.0221367 x 1023 things 1 amu = 1.660540 x 10-24 g
Suppose you had a sample of one mole of particles. Each particle weighed exactly 1 amu. How many amu would the sample weigh? How many grams would the sample weigh?
IX. Counting Atoms and Stuff 1 mole of amu = (6.0221367 x 1023)
(1.660540 x 10-24 g)= 1.00000 gram (exactly)
What is the significance of this?
How much will 1 mole of C-12 atoms weigh (in grams)?
12 grams (exactly)
IX. Counting Atoms and Stuff
Molar Mass: mass of one mole of a substance units: grams/mole equal to the ATOMIC MASS of the element for compounds, equal to the SUM of the
masses of all the elements in the compound (multiply each elements’ atomic mass by the subscript)
IX. Counting Atoms and Stuff
Example: Find the molar masses of the following:
NaCl CO2
Ca(C2H3O2)2
Na = 22.99 Cl = 35.45 C = 12.01O = 16.00 H = 1.008 Ca = 40.08
58.44 grams/mol
44.01 grams/mol
158.168 grams/mol
IX. Counting Atoms and Stuff
1 mole of X = atomic mass of X (grams)
1 mole X = 6.022 x 1023 atoms
These equalities will let you do DIMENSIONAL ANALYSIS Convert grams to moles and moles to grams Convert moles to atoms and atoms to moles
NEVER EVER put 6.022 x 1023 in front of GRAMS
IX. Counting Atoms and Stuff
1 molecule C6H12O6 has 6 C atoms, 12 H atoms, and 6 O atoms.
1 mole of C6H12O6 has ___ moles of C atoms, ____ moles of H atoms, and ___ moles of O atoms.
612 6
IX. Counting Atoms and Stuff
1 mole XaYb = a moles X = b moles Y
Ex.1 mole Na2C2O4
= 2 moles Na = 2 moles C = 4 moles O
IX. Counting Atoms and Stuff 1 mole X= 6.0221367 x 1023 atoms of X 1 mole of X = atomic mass of X (grams) a moles X = 1 mole XaYb = b moles Y
IX. Counting Atoms and Stuff
Moles X Moles XaYb
atoms X
grams Xgrams XaYb
molecules XaYb
1 mole of X = atomic mass of X
1 mole X= 6.0221367 x 1023 atoms
IX. Counting Atoms and Stuff Problem-solving: MAP the problem first (determine what you
are starting with, where you want to end up, and the path to follow).
Example:What is the mass of 3.60 moles of Cu?
IX. Counting Atoms and Stuff
Moles X Moles XaYb
atoms X
grams Xgrams XaYb
molecules XaYb
1 mole of X = atomic mass of X
1 mole X= 6.0221367 x 1023 atoms
START
End
IX. Counting Atoms and Stuff Map out the problem first (determine what
you are starting with, where you want to end up, and the path to follow).
Example:What is the mass of 3.60 moles of Cu?
Moles of Cu grams of Cu
3.60 moles Cu xmoles Cu
grams Cu
1
63.546= 229 grams Cu
IX. Counting Atoms and Stuff
Map the following problems FIRST, then solve: How many moles are in 11.9 grams? How many atoms are in 3.60 x 10-10 grams
of gold? How many grams of sodium are in 2.34
moles of Na2CO3? How many grams of Fe are in 13.86 grams
of Fe2O3?
IX. Counting Atoms and Stuff
Moles X Moles XaYb
atoms X
grams Xgrams XaYb
molecules XaYb
1 mole of X = atomic mass of X
1 mole X= 6.0221367 x 1023 atoms