Download Atomic Theory of Matter: matter is made up of fundamental particles

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Atomic Theory of Matter: matter is made up of _______________________ particles
Democritus (~400 B.C.), Greek philosopher: “_________________”
 Believed that atoms were indestructible and ________________________.
___________________ Atomic Theory (1st Scientific Theory, 1803)
Dalton learned about atoms by studying the ratios in which elements combine in chemical reactions.
1. All matter is made up of _________________.
2. Atoms are _______________________ and ___________________________. [Not true: subatomic particles]
3. a) Atoms of the same element are ________________________. [Not entirely true: isotopes]
b) Atoms of one element are _______________________________ from atoms of another element.
4. Atoms can physically mix together or can ________________________ combine in
_________________-number ratios to form compounds.
5. Chemical reactions occur when atoms are ______________________, ________________, or ___________________. However,
atoms of one element are never changed into atoms of another element as a result of a chemical reaction.
Thomson’s Plum Pudding Model (1897)
 Atoms are __________________________ and can be broken down into _____________________ particles.
 Discovered the ___________________ and the ____________________.
 Envisioned the atom as tiny particles of negative charge embedded in a ball of positive charge.
 Electron
o Passed an electric current through gases in a ____________________________ producing a glowing beam
(cathode ray).
o If electrically charged plates are placed near the cathode tube, the cathode ray will be attracted
towards the positive plate and repelled from the negative plate.
o Therefore, the ray must be made of ____________________ charged particles.
 Proton
o Inferred the existence of positively charged particles since matter is ______________! (You do not get a
shock every time you touch matter!)
Milikan’s Oil Drop Experiment (1916)
 Determined the charge and mass of an ___________________.
Chadwick (1932)
 ___________________
o ____________________: atoms of an element that are chemically alike but differ in mass
 same number of protons and electrons, _______________ number of __________________!
Particle
Proton
Neutron
Electron
Symbol
Relative charge
Relative Mass
Rutherford’s Gold Foil Experiment (1911)
 A sheet of gold foil was bombarded with _________________ ___________________.
 Most of the particles passed through the foil.
=> ______________________________________________________.
 A few particles were deflected.
=> ___________________________________________________________.
 Nucleus: small, dense core containing protons and neutrons.
o Electrons surround the nucleus.
Actual Mass (g)
1.67 x 10-24
1.67 x 10-24
9.11 x 10-28
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Bohr’s Planetary Model (1913)
 Bohr studied how atoms emit or absorb ___________________.
 Electrons move in circular paths (orbits) around the positive nucleus.
 Electrons overcome the attraction of the nucleus because they are in motion.
 Given __________ ________________, electrons will move to an orbit ______________ ___________ from the nucleus.
 Electrons can only occupy certain orbits (__________________________).
To move from one energy level to another, an electron must absorb or release just the right amount of energy.
 Energy Levels: regions of ______________ in which ________________ can move about the nucleus of an atom.
Quantum Mechanical Model (Wave Model or Electron Cloud Model)
 Schrödinger (1926)
 Determines the allowed energies an electron can have.
 Determines how likely it is to find the electron in various locations.
 Cannot pinpoint the ___________________ paths of electrons (only probability).
 Probability of finding an electron is represented by an electron cloud.
 ___________________________________________: It is impossible to know the velocity and position of a particle at the
same time.
Subatomic Particles
Mass Number: # of _______________ + # of _________________
 Most of the mass of an atom is concentrated in the nucleus.
Atomic Number: # of ________________
 Every element has a different number of protons.
 In a neutral atom, # ________________ = # ___________________.
Example:
9
4
Be
Mass Number = 9 Atomic Number = 4
Isotopes: Atoms that have the same number of protons but a different number of neutrons.
 Isotopes have different mass numbers because they have a different # of neutrons.


Example: Carbon – 12, Carbon – 14
(ALL carbon atoms have 6 protons and 6 electrons)
Isotope
Chlorine-35
Chlorine-37
Atomic #
Mass #
#p+
#e-
#n0
% of Naturally Occurring Chlorine
75%
25%
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Average Atomic Mass: sum of the masses of all the protons and neutrons
 Measured in grams or atomic mass units (amu).
 An _______________________________(amu) is defined as one twelfth of the mass of a carbon-12 atom.

Average atomic mass (on the ______________ _______________): the weighted average of all the naturally
occurring isotopes of that element.
Mass of a proton = 1.001 amu
Mass of a neutron = 1.001 amu
Mass of electron = 0.0005486 amu
Average Atomic Mass = ________________________________________________________
Example: Calculate the avg. atomic mass of oxygen if its abundance in nature is 99.76% 16O, 0.04% 17O, and 0.20%
18O.
Ion: formed with an atom ______________ or _______________ an ___________________ and becomes ___________________.
 Gained an electron: ________________ charge
o Has more electrons than protons
Na1+: 11 protons __________= ______ electrons
 Lost an electron: ______________ charge
O2-:
8 protons ___________= ______ electrons
o Has more protons than electrons
To find the number of electrons, start with the number of protons and do the opposite of what the charge says.
Electrons
Atomic Orbital: region of _______________ in which there is a high ____________________ of finding an _____________.
 Denoted by letters: _____, ______, _______, and ______
 Each atomic orbital corresponds to a specific shape at a specific energy level.
Sublevel
Energy Level
Number of Orbitals
Number of
Sublevels
Type of Sublevel
Maximum Number of Electrons
Maximum Number of
Electrons
1
2
3
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Electron Configuration
The most stable arrangements of electrons in sublevels and orbitals.
General Rules:
Aufbau Principle: electrons fill the ________________ energy level first.
 “Lazy Tenant Rule”
Pauli Exclusion Principle: each orbital can hold ______ electrons with ____________ spins.
Hund’s Rule: within a sublevel, place _______ electron per ___________ before pairing them.
 “Empty Bus Seat Rule”
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Notation:
Orbital Diagram:
Electron Configuration:
O: 1s2 2s2 2p4 (long-hand); O: [He] 2s2 2p4 (short-hand)
Periodic Patterns:
Period Number (Row #): represents the ____________ __________________ (subtract 1 for d; 2 for f)
A/B Group Number (Column #): represents total number of ________________ electrons.
Valence Electrons: electrons in the outermost energy level
o Determines the chemical and physical properties of elements.
Column within sublevel block: represents number of electrons in sublevel.
Example: Hydrogen
1s1
Determining Short-Hand Configuration:
Core e-: Go up one row and over to the Noble Gas.
Valence e-: On the next row, fill in the # of e- in each sublevel.
Example: Germanium
Stability:
Electrons are most stable with a full energy level or a half-filled energy level.
Exceptions in Filling Order
 Transition elements sometimes fill unpredictably because the energy levels 4s and 3d are so close in
energy.
Example:
Copper: 1s22s22p63s23p64s13d10
Example: Chromium: 1s22s22p63s23p64s13d5
Copper gains stability with a full d-sublevel
Chromium gains stability with a half-filled sublevel
Ion Electron Configuration:
Write the electron configuration for the closest ______________ ________.
Example: Oxygen ion: O2- (gained 2 electrons: 8 + 2=10) ≅Ne (10 electrons) = [He]2s22p6
Ion electron configurations will always have a filled p-sublevel!