Transcript Rock On! Climbing Club - School District of La Crosse

Chapter 4 notes
• I. GREEK THEORY
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A. DEMOCRTUS- The universe is made of indivisible
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pieces of matter called atoms
1. atom- Greek for unable to divide
2. Not able to explain all observations though
and was dismissed
II DALTON’S THEORY
A. Every element is made of tiny unique particles
called atoms and cannot be subdivided
B. atoms of the same element are exactly alike
C. atoms of different elements can join to form
molecules
• III ATOMIC THEORY
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A. Nucleus- The center of the atom with a positive charge
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1. Protons- positively charged particle in the nucleus with an
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atomic mass of 1AMU( ATOMIC MASS UNIT)
2. Neutrons- The neutral charged particle of the nucleus
having the mass of 1 amu
B. Electron cloud- The region around the nucleus with a
negative electric charge
1. contains electrons
2. Rutherford Gold Foil Experiment
• B.Ruherford’s gold foil experiment
• 1.Hypothesis- the atom is solid and the
particles will bounce off
• 2.Observations- most of the particles
went straight through, except once in a
while one would bounce back
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3.New conclusions- the atom is mostly
empty except it has a dense core
• IV. MODELS OF THE ATOM
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A. BOHR’S MODEL- An explanation of electron
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movement about the nucleus
1. Suggested energy levels
a. energy levels- any of the possible energies
an electron can have in the atom. Level 1= 2 e:
level 2=8e: level 3= 18e: level 4=32: level 5=50
1. electrons gain energy to move to a higher
level and give up energy when they drop back into a
lower level( ground state)
1. No longer assumed the electron moved in definite
orbitals( Heisenberg uncertainty principle)
• B. NEW VIEW OF BOHR’S MODEL
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1. Electrons behaves more like waves
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2. spinning fan image- Cannot locate
the location of the fan blade at any time, just
know it’s in the blur
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a. same applies to an electron,- The
best we can do is give the probable location
of the electron.
• C. ENERGY LEVELS OF ELECTRONS
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1. Orbital- the region where there is a high probability of finding the
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electron.
a. s orbital- the simplest has a spherical shape
1. 1 possible orientation in space
2. surrounds the nucleus
b. P orbital1) Dumbbell shaped
2) 3 orientations in space ( X.Y,Z)
c. d and f orbitals- much more complex
1) 5 possible d orbitals
2) 7 possible f orbitals
2. Electrons occupy the lowest energy levels in the atom, and
within the levels they occupy orbitals with the lowest amount of energy.
s- ORBITAL- SPHERICAL
P- ORBITAL
D- orbital
f- ORBITAL
Sub-level organization
• Atomic orbitals
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1s
2s
3s
4s
5s
6s
7s
Level 1
2p Level 2
3p 3d Level 3
4p 4d 4f Level 4
5p 5d 5f 5g Level 5
6p 6d 6f 6g Level 6
7p 7d 7f Level 7
• 2. Electrons occupy the lowest energy levels in
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the atom, and within the levels they occupy
orbitals with the lowest amount of energy.
a. s has lowest then p followed by d and f
3. Valence electrons- The outer level of
electrons
a. The combining ability of the atom
depends on the valence
• IV TOUR OF THE PERIODIC TABLE
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A. Organization of the periodic table
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1. Developed by Dmitri Mendeleev 1869
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a. arranged the elements according to repeating
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properties
b. He left blank spaces for elements yet to be
discovered.
2. Groups similar elements together
a. makes it easier to predict properties
b. arranged based on the number of protons in the
nucleus
c. Periodic Law-properties of elements tend to change
in a regular pattern when elements are arranged in order of
increasing atomic number or the number of protons in their atom.
• B. Using the periodic table
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1. Periods- horizontal rows
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a. increasing number of protons
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b. each period represents the energy levels of
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electrons.
2. Groups( family)- the vertical columns
a. tells the number of valance electrons
1. valence electrons- determines the
chemical properties of the element
• C Ionization- The process of adding or
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removing electrons from an atom,
thus forming a charged atom .
1. remove an electron the atom will have
more protons than electrons, and the atom is
positive( metals)- CATIONS
2. add an electron to the atom. The atom
will have more electrons than protons, thus it
will be negative( nonmetals) ANIONS
• D. HOW ATOMS DIFFER IN STRUCTURE
• 1. ATOMIC NUMBER(Z)- This tells the number of
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protons and electrons in an electrically neutral atom.
2. Mass number- This is the total number of protons
and neutrons in an atom.
a. isotope- An atom with the same atomic
number, but different atomic mass
1. Hydrogen- deuterium, tritium, protonal
b. calculating the number of neutrons- The
mass number- atomic number
• 3. Atomic mass unit- equal to 1/12 the
mass of a carbon-12 atom
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4. Average atomic mass- This is an
umber which represents the average of
the atomic masses including every
isotope.
PERIODIC TABLE OF THE
ELEMENTS
• V.Broad classification• 1. metals characteristics- left side of P.T.
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a. shiny
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b. ductile- pulled into wire
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c. malleable- pounded into sheets
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d. excellent conductors of heat and
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electricity
e. Donate electrons- become + ion
• 2. Nonmetals characteristics- right side of
P.T.
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a. dull
b. brittle
c. poor conductors
d. Accept electrons= negative ion
• 3. Metalliods- along the zig-zag
• a. properties of metals and nonmetals
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b. semiconductors
• VI SURVEY OF THE PERIODIC TABLE
• A.ALKALI METALS (Group 1)
• 1 soft high reactive metals.
• 2.When mixed with water they form a slippery solution
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that can remove grease.
3.Alkali metals are bases, and are highly reactive
because they have only one electron in the outermost
shell….why?
4. Alkali metals are good conductors of electricity.
Sodium is used in street lamps and fog lights.
5. Lithium is used as a medication to treat
depression.
• B.ALKALINE EARTH METALS (Group 2)
• 1 are harder, denser, stronger, and have higher melting
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points that the group 1 metals.
2.These metals are less reactive than group 1
because they have 2 electrons in the outer shell.
3. In order for them to have a full outer shell they
must loose two electrons.
4.Magnesium is used in construction because of its
strength and lightness.
5.Calcium is the best know of the alkali earth
metals.
• C. TRANSITION ELEMENTS (Group 3-12) are all
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metals.
1.Transition metals are not as reactive as
metals in groups 1 & 2.
2.They are harder denser and have higher
melting points than groups 1 and 2.
3.Transition elements do not have similar
electron configurations as do groups 1 & 2.
4.There are two more periods of transition
elements located at the bottom of the periodic
table.
• D. Boron family- group 13
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1. Boron- used in control rods
of nuclear reactors, and detergents
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a. 3 valence electrons
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2. aluminum- Used as a light
weight metal in industry.
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• E. Carbon family
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1. 4 valence electrons
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2. CARBON
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a. Basis of life on this planet
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b, Bonds with itself and many other
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elements
3. Silicon and germanium- the heart of the
computer industry
4. Tin- Acts as a metal - alloys to make
useful metals
5. Lead- Usually the end product of
radioactive decay
a. very dense
• F. NITROGEN FAMILY
• 1. 5 valence electrons
• 2. nitrogen and phosphorus- Used in
fertilizers and explosives
• 3. Arsenic- a poison used in insecticides.
• 4. Bismuth- Used in automatic sprinkler
systems because it has a low melting
point
• G. OXYGEN FAMILY
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1)6 valence electrons- wants 2 to be full, fairly
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reactive
2) Oxygen- reactive
a. Used in respiration to release ATP
3) Sulfur- some in ringed forms
a, Vulcanization of rubber- makes rubber
harden
b. Selinium- poisonous, but necessary in small
amounts
1. dairy industry- cows
c. Polonium- radioactive
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H. HALOGENS- SALT PRODUCERS
1. 7 valence electrons... very reactive, only want 1 electron
2. fluorine- very reactive and poisonous
a. Used to etch glass- HF acid
3. Chlorine- Very poisonous
a. used in bleaches and detergents- sodium hyperchloride
4. Bromine- purple liquid at room temperature
5. iodine- Gray crystals
a. sublimes when heated
b. necessary for the proper functioning of the thyroid gland
6. astatine- Radioactive
• I. NOBLE GASES- VERY BORING
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A) 8 electrons in the valence level, so all levels are
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full.
B) He is lighter than air so used in airships
a. helium comes from the Greek work
helios which means sun
1. discovered by studying the spectrum of
sunlight
2) the rest of them are used in lighted signs
or as gases to extend the life of light bulbs
• VII. USING MOLES TO COUNT ATOMS
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A. ITS LIKE A PACKAGING TERM
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1. dozen=12
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2. gallon= 128 fluid ounces
• B. CHEMIST DEAL WITH A LARGE
NUMBER OF SMALL PARTICLES.
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1. Mole-The SI unit which describes
the amount of a substance
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a. 1 Mole of anything is 6.022x1023
of anything
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1) 6.022x1023 is Avogado’s
constant
• 2. Molar mass(FWIG)- The mass of 1 mole of
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any substance, or the mass of 6.02x1023 of
anything
a. The average atomic mass is the weight of
1 mole of the substance
1. 16g of oxygen = 1 mole
2. 18g of H2 O =1Mole
3.molar mass of NaCl is?
b. The molecular weight of a compound also
equals 1 mole
• C Using conversion factors for calculating moles
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1. Must make sure all the units cancel away, and end up with the
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desired units.
PRACTICE
1. HOW MANY MOLES OF WATER ARE IN 72 g OF WATER
72gx 1m/18g=4M
2. HOW MANY PARTICLES ARE IN 72 g OF WATER
4Mx 6.02x1023 particles/1mole= 24.08x1023 particles
3. HOW MANY MOLES ARE IN 174g OF NaCl
174gx 1m/58g=3M
4.THE NUMBER OF PARTICLES IN NaCl
4M x 6.02x1023 particles
• 30g of HCl=______atom
• 30gx 1m/36g x 6.02x1023atoms/1m
• 98g of H2 SO4=_______atoms
• 98gx 1m/98g x 6.02x1023 atoms/m
• 300 g of Ba3 (PO4)2=______atoms
• 300g x 1m/601g x 6.02 x1023 /m
• 6M of NaCl=_______________atoms
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6M x 6.02 x1023 atoms/m
• .9M of H2 O=___________atoms
More practice
• 1) 20g Na=__________M
• 2) 45g F=___________M
• 3) 126g O=__________M
• 4) 30g Na2 O=________M
• 5) 3M Ba=___________g
• 6) 6M CH4=_________g