Download Chapter 3 Notes 2015

Chapter 3
Atoms: The Building Blocks of Matter
Lesson 1.1: Early Atomic Theory
• Learning Target: I will understand the history and
structure of the atom.
• Success Criteria: I can describe the experiments and
discoveries of Democritus, Dalton, Thomson, and
Rutherford. I can describe Thomson’s and
Rutherford’s model of the atom and the know the
subatomic particles involved.
Democritus (400 B.C.E)
• Greek philosopher
• Hypothesized: Nature has a basic
indivisible particle of which everything
is made of
• Called this particle an atom
• Greek “atomos” = indivisible
John Dalton ( 1808 A.D.)
• English schoolteacher – liked nature and weather
Dalton’s 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 and are different from atoms of other elements
3. Atoms cannot be subdivided, created, or destroyed
4. Atoms of different elements combine in simple whole number
ratios to form chemical compounds
5. In chemical reactions, atoms are combined, separated or
rearranged
Issues with
Dalton’s Atomic Theory
• Atoms can be split into even smaller particles
(nuclear chemistry) and aren’t indivisible
• i.e. nucleus, protons, electrons
• A given element can have different masses
• i.e. isotopes
J. J. Thomson (1897 A.D.)
• Discovered: The 1st subatomic particle: the negatively charged electron,
developed the plum pudding model
• Used a Cathode Ray Experiment
• Cathode Ray Tube – Electric current passed through a
metal disk to another metal disk in a gas at low pressure
(vacuum sealed tube)
• i. e. neon signs and ‘old-fashioned’ television sets
Cathode Ray Experiment
• When a current passed through the
cathode ray tube, the surface of the
tube opposite the cathode glowed
• Glow was hypothesized to be stream of
particles called a cathode ray
• Ray affected by magnetic fields
• Attracted to positive charge
• Deflected from negative charge
• http://www.youtube.com/watch?v=7YH
wMWcxeX8&NR=1
Discovery of the
1st subatomic particle
• Thomson measured the ratio of the charge of the
particles to their mass
• Same ratio no matter what metal or gas was used
• Named this particle an electron
• http://www.youtube.com/watch?v=IdTxGJjA4Jw&feature=related
J. J. Thomson’s
Plum Pudding Model
• Atoms are electrically neutral
• Must have positive charges to balance the negatively charged electrons
• Electrons have a lot less mass than atoms
• Other particles must account for their mass
• Plum Pudding Model
• positively charged sphere with electrons dispersed through it
Ernest Rutherford (1911 A.D.)
• Discovered: The nucleus using the gold foil experiment
• Gold Foil Experiment
• Bombarded thin piece of gold foil with alpha particles
• Expected alpha particles to pass through with minimal deflection
• Surprised when 1 in 8000 deflected back to source
• It was “as if you had fired a 15 inch artillery shell at a piece of tissue paper and it
came back and hit you”
http://www.youtube.com/watch?v=wzALbzTdnc8&feature=related
Rutherford’s
New Model of the Atom
• Discovered the nucleus is a small densely packed
volume of positive charge
• Surrounded by cloud of electrons (more on this in
chapter 4!)
• Size comparison
• Nucleus = marble
• Whole Atom = football field
Lesson 2.2: Basic Laws
• Learning Target: I will understand 3 of the basic laws
of chemistry.
• Success Criteria: I can define the law of conservation
of mass, the law of multiple proportions, and the law
of definite proportions.
1790s – Discovery of Basic Laws
• Law of Conservation of Mass
• Mass is neither created nor destroyed during ordinary chemical
reactions or physical changes
• Law of Definite Proportions
• A chemical compound contains the same elements in exactly
the same proportions by mass regardless of size of sample or
source of compound
• i.e. Every sample of table salt is made of 39.34% Na and
60.66% Cl
• i.e. H2O always has 2 atoms of H and 1 atom of O
Basic Laws Continued
• Law of Multiple Proportions
• 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
• i.e. CO and CO2
• CO = 1.00g of C and 1.33 g of O
• CO2 = 1.00 g of C and 2.66 g of O
• The ratio of the second element is 2.66 to 1.33 or 2 to 1
Lesson 2.3: Inside the Atom
Learning Target: I will understand the structure of the
atom.
Success Criteria: I can describe the location, mass, and
charge of each subatomic particle.
Recap: Structure of the Atom
• Atom (today’s definition) = Smallest particle of an element
that retains the chemical properties of that element
• Two regions:
• Nucleus
• Very dense, small center of the atoms
• Protons and neutrons
• Electron Cloud
• Region occupied by electrons
Inside the Nucleus
• 2 types of particles inside the nucleus:
• Protons = positively charged (+1)
• Neutrons = neutral charge (0)
• Mass in the nucleus
• Protons = 1 amu
• Neutrons = 1 amu
• 1 amu (atomic mass unit) = 1.66 x 10-24 g
Where are the Electrons?
• In the Electron Cloud
• A cloud of negative charge
outside of the nucleus (more in
chapter 4!)
• Electrons = Negatively
charged particles (-1)
• Mass of electrons = 0 amu
• Actually it is 9.109 x 10-31 g, but
we round to zero for simplicity
Properties of Subatomic Particles
Particle
Symbol
Charge
Mass (amu)
Electron
e-, 0e
-1
0
Proton
p + , 1H
+1
1
Neutron
n◦, 1n
0
1
Characteristics of Atoms
• Atomic Number = number of protons
• Identifies the element
• # of protons is what give that element its characteristic properties
• Elements with different protons are NOT THE SAME
ELEMENT!!!
Lesson 3.4: Atoms, Ions, and
Isotopes
Learning Target: I will understand what atoms, ions,
and isotopes are, and their relationship to subatomic
particles.
Success Criteria: I can count the number of subatomic
particles of each type in atoms, isotopes, and ions.
Atoms
• Atoms have a neutral charge (charge=0)
• total positive charge equals the total negative charge
• # protons (+1 each) = # electrons (-1 each)
Ions
• Atoms with a charge
• Negative – more electrons than protons
• Positive – more protons than electrons
• Charge = #protons - # electrons
• Example: Magnesium atom with 12 protons and 10 electrons has a
charge of +2
Isotopes
• Atoms of the same element (i.e.
same # of protons) that have
differing number of neutrons
• Isotopes of the same element:
• have different masses
• do not differ significantly in
chemical behavior because they
have the same # of protons
• Are named by their mass number
Mass Number
Mass number = #protons + # neutrons
Element
Atomic
Number
# of
Protons
# of
Neutrons
Mass
Number
Carbon
6
6
6
12
Oxygen
8
8
8
16
Nitrogen
7
7
8
15
Lesson 2.5: Average Atomic Mass
Learning Target: I will understand the relationship of
isotopes to an atoms average atomic mass.
Success Criteria: I can calculate the average atomic mass
of an element given the mass and percent abundance
of each of it’s isotopes.
Average Atomic Mass
• Every element has isotopes
• The periodic table takes into account all naturally occurring
isotopes of an element and averages them
• Average Atomic Mass is listed on the periodic table
• UNIT is amu = atomic mass unit
How to Calculate Average Atomic Mass:
Mass of isotope #1 x
abundance in nature
(decimal)
+
Mass of isotope #2 x
abundance in nature
(decimal)
+…
=
Average
Atomic Mass
Example of Calculating the Average
Atomic Mass – Hydrogen
• There are two naturally occurring isotopes of hydrogen
• Hydrogen with 1 proton and zero neutrons
• Hydrogen with 1 proton and one neutron
• Differentiating between the two isotopes (symbol – mass number)
Hydrogen Isotopes
Element – mass #
Atomic Mass
Naturally occurring
abundance %
1 proton
0 neutrons
H-1
1.007825 amu
99.9885
1 proton
1 neutron
H-2
2.014102 amu
0.0115
Calculation:
(1.007825 x 0.999885) + (2.014102 x 0.000115) = 1.01 amu