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1 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 2 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% 3 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 4 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” 4 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!