Chemistry UNIT 3. Name: Date: Chemistry Unit 3 Atomic Theory and structure of an Atom.
Transcript of Chemistry UNIT 3. Name: Date: Chemistry Unit 3 Atomic Theory and structure of an Atom.
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Chemistry UNIT 3
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Name:Date:
Chemistry Unit 3
Atomic Theory and structure of an Atom
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Definitions
Model: A familiar idea used to explain unfamilar facts observed in nature.
Theory: An explanation of observable facts and phenomena– To remain valid, models and theories must:
Explain all known facts Enable scientists to make correct predictions
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History of an Atom
Democritus– Proposed the existence of an atom– Word comes from the Greek word atomis which
means not to cut or indivisible
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Aristotle– Rejected the idea of the atom– Said matter could be cut continually
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Dalton’s theory proposed that atoms:– Are building blocks of matter– Are indivisible– Of the same element are identical– Of different elements are different– Unite in small, whole number ratios to form
compounds
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J.J. Thomson– Credited with the discovery of electron; a blow to
Dalton’s indivisible atom– Proposed the plum pudding model of the atom:
negatively charged electrons embedded in a ball of positive charge
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Rutherford’s Gold Foil experiment:– Aimed alpha particles at gold foil– Most passed through– A few particles were deflected– Some particles bounced back
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Rutherford’s Experiment
Most of the atom is empty space Dense positively charged core Planetary model
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Bohr’s Model of the Atom
Nucleons- particles in the nucleus of atom– Protons– Neutrons
Atomic number- number of protons in the nucleus of an atom
Neutral atom- same number of protons (+) and same number of electrons(-)
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Isotopes
Isotopes- atoms of an element that have different numbers of neutrons
Hydrogen-1– _______ proton and ______ neutrons
Hydrogen-2– _______ proton and ______ neutrons
Hydrogen-3– _______ proton and ______ neutrons
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Mass number
Total number of protons and neutrons in an atom– Carbon-14– Neon-20
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Opening
What is the difference between
C-12 and C-14?
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Particle Chart
Particle Charge Mass Location
Proton Positive 1 amu nucleus
Neutron Neutral 1 amu nucleus
Electron Negative 0 Electron cloud
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Atomic mass
Average of the masses of all the element’s isotopes
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Subatomic particles
# of protons = atomic number # of electrons = atomic number # of neutrons = mass number – atomic
number
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Examples
Iron Fe-56 Oxygen-17 He-4 Calcium-40
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Bohr’s Energy Levels
Electrons in certain energy levels Low energy levels are closer to nucleus High energy levels are further from
nucleus Ground state- all electrons are in lowest
energy level possible
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Excited Atom
Atom has absorbed energyExcited state is unstableAtom soon emits same amount of
energy absorbedEnergy is seen as visible light
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Wave Description of Light
Wavelength (): distance between corresponding points on adjacent waves
Frequency (f): the number of waves passing a given point in a given time
c = 3.0 X 108 m/s, speed of light c = f
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Ex.
What is the frequency of light if the wavelength is 6.0 X 10-7m?
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Particle Description of Light
Energy exists as particles called quanta or photons.
E = hf
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The Modern View of Light
Light has a dual nature Light may behave as a wave Light may behave as a stream of particles
called quanta or photons.
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Spectroscopy
Spectral lines represent energy releases as electron returns to lower energy state
Spectral lines identify an element Called Bright line spectrum of an element
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Orbital
Region of space where an electron is likely to be found
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Quantum Numbers
n, l, m, s Used to describe an electron in an atom
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n
Principle quantum number Represents the main energy level of electron Is always a whole number Max. # of electrons in an energy level is 2n2
– What is the maximum number of electrons that can be in the 5th main energy level?
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l
The 2nd quantum number Describes the orbital shape within an energy
level Number of orbital shapes possible in energy
level = n
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Orbital shapes
Designated s, p, d, f Level 1: s Level 2: s,p Level 3: s, p, d Level 4: s, p, d, f
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How many electrons can each sublevel hold?
s = 1 orbital X 2 electrons= 2 electrons p = 3 orbitals X 2 electrons = 6 electrons d = 5 orbitals X 2 electrons = 10 electrons f = 7 orbitals X 2 electrons = 14 electrons
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m
The 3rd quantum number Describes orientation of orbital in space x, y, z axis
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s
The 4th quantum number Describes spin of electron in orbital Hund’s Rule- orbitals of equal energy are
each occupied by one electron before any orbital is occupied by a second electron
Pauli Exclusion Principle: No two electrons can have the same four quantum numbers.
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Diagonal Rule
See worksheet