Download Chapter 3 Notes

Accelerated Chemistry
Chapter 3 Notes
Mr. Seidel
(Student’s edition)
Chapter 3 problem set: 2, 7, 10, 13, 17ace, 18, 20, 21, 22ab, 23ab
3.1
The Atom: From Idea to Theory
Historical Background- In approximately 400 BC, Democritus (Greek) coins the term
“atom” (means
). Before that matter was thought to be one continuous piece
- called the continuous theory of matter. Democritus creates the discontinuous theory of
matter. His theory gets buried for thousands of years
18th century - experimental evidence appears to support the idea of atoms.
Law of Conservation of Mass - Antoine Lavoisier (French) - 1770’s
Law of Definite Proportions - 1799 - Joseph Proust (French) - “The proportions of
masses of chemicals in reactions is always the same.”
example
Law of multiple proportions - 1803 - John Dalton - English school teacher
“The mass of one element combines with masses of other elements in simple,
whole # ratios.”
example
Dalton’s Atomic Theory - Dalton put together the laws of conservation of mass,
definite proportion, and multiple proportion to create his own atomic theory.
1. All matter is made up of atoms which are indivisible
*2 Atoms of the same element are identical
*3. Atoms cannot be created or destroyed
4. Atoms combine in simple whole number ratios
5. Chemical reactions are the result of the separating or combining of atoms
* #2 is wrong because of the existence of
* #3 is wrong as an exception is
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)
3.2
The Structure of the Atom
Updating Atomic Theory
1870’s - English physicist William Crookes - studied the behavior of gases in
vacuum tubes (Crookes tubes - forerunner of picture tubes in TVs). Crookes’
theory was that some kind of radiation or particles were traveling from the
cathode across the tube. He named them
.
20 years later, J.J. Thomson (English) repeated those experiments and devised
new ones. Thomson used a variety of materials, so he figured cathode ray
particles must be fundamental to all atoms. 1897 - discovery of the ___________.
Charge and Mass of the electron - Thomson and Milliken (oil drop experiment)
worked together to discover the charge and mass of the electron
charge =
mass =
this is the smallest charge ever detected
this weight is pretty insignificant
1909 - Gold Foil Experiment (Rutherford - New Zealand)
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atom -
atoms are made up of __________________,_______________,______________
Subatomic
particle
Proton
Neutron
Electron
Location
Charge
Relative Mass
Determines…
Not all atoms of the same element are identical…
Isotope –
. Some isotopes occur naturally - most are produced artificially.
The isotopes of Hydrogen
Name
Drawing
Relative Mass
Relative
Abundance
3.3Weighing and Counting Atoms
We look to the periodic table to give us information about the number of particles are in
atoms and also to help us count atoms in a sample.
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atomic # (Z) - (always a whole number, smaller number on the periodic table) = # of
protons in the nucleus - also indicates the # of electrons if the element is not charged
atomic mass – the average mass of all of the isotopes of an element – is a number with a
decimal – is always the larger number on the periodic table.
mass number (A) - sum of the protons and neutrons in a nucleus
this number is rounded from atomic mass due to the fact that there are isotopes
# neutrons =
example - # of neutrons in Li =
Ion – a charged atom. Atoms become charged by gaining electrons (become a negative
charge) or losing electrons (become a positive charge)
Lots of practice:
Element
Atomic
(ion/isotope) Number
Atomic
Mass
Atomic
Mass
Number
Proton
Electron
Neutron
He
Be
F
F-1
K
K+1
35
Cl
37
Cl
25
Mg+2
Modern standard of atomic mass and determining masses from weighted averages
Mass of Cl thought to be 35.5 X that of hydrogen
today we know this isn’t true- it’s the weighted average of 2 isotopes
75%
35
Cl
and
25%
37
Cl
average atomic mass = (%)(mass of 1st isotope)+(%)(mass of 2nd isotope)......
sample problem - find the average atomic mass of B
B11 = 80.20%
B10 = 19.80%
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sample problem - find the %’s of 2 isotopes of Carbon given the following information:
average atomic mass = 12.0111 isotope 1 = 12 C , isotope 2 = 13 C
History lesson - originally H was the basis of all atomic masses and was given the
mass of 1.0. Later, chemists changed the standard to oxygen being 16.000 (which left H
= 1.008). In 1961, chemists agreed that 12C is the standard upon which all other masses
are based.
1/12 of the mass of 1 atom of 12C = 1 amu
The Mole, Avogadro’s number, and Molar Mass
Atoms are tiny, so we count them in “bunches” - a mole is a “bunch of atoms”
mole (definition) - The amount of a compound or element that contains 6.02 x 1023
particles of that substance.
1 mole = 1 molar mass = 6.02 x 1023 particles
Molar Mass - the sum of the atomic masses of all atoms in a formula
Round to the nearest tenth! (measured in amu or grams)
ex - H2
H2O
Ca(OH)2
molar mass is a term that can be used for atoms, molecules (covalent compounds or
elements) and formula units (ionic compounds)
Official names may also be:
Formula mass (ionic compounds)
Molecular mass (covalent compounds and diatomic elements)
Atomic weight, Atomic mass, grams formula weight, etc.
examples:
1 mole Na =
1 mole O2 =
1 mole HCl =
1 mole NaCl =
atoms =
molecules =
molecules =
formula units =
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g
g
g
g
Mole Relationships
For starters in chemistry, we have to be able to convert between moles, grams, and
molecules/atoms of substance (also liters when we work with gases)
use the “mole map”
_____ mole
_____ atom/molecule
_____ gram
_____ liter (at STP)
2 steppers
convert 13.8 g Li to moles
convert 2.0 moles Ne to g
convert 3.0 moles of Be to atoms
convert 44.8 L of O2 to moles
3 and 4 steppers
convert 1.2 x 1024 atoms of Magnesium to grams
convert 128 g of O2 to molecules of O2
convert 128 g of O2 to atoms of oxygen
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