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Chapter 4 “Atomic Structure” Section 4.1 Defining the Atom OBJECTIVES: Describe Democritus’s ideas about atoms. Explain Dalton’s atomic theory. Identify what instrument is used to observe individual atoms. Section 4.1 Defining the Atom _ First to suggest the existence of atoms (from the Greek word “atomos”) He believed that atoms were _ . Dalton’s Atomic Theory 1) are composed of tiny called . John Dalton (1766 – 1844) 2) Atoms of the . --Atoms of any _ are are from those of any . Dalton’s Atomic Theory 3) Atoms of elements in _ ratios to form . 4) In chemical reactions, atoms are – but of _ into atoms . Sizing up the Atom 100,000,000 atoms = 1 cm 1,000,000 atoms = width of hair Can be observed with _ . Section 4.2 Structure of the Nuclear Atom OBJECTIVES: Identify three types of subatomic particles. Describe the structure of atoms, according to the Rutherford atomic model. Section 4.2 Structure of the Nuclear Atom Atoms are divisible into particles: _ _ _ _ Discovery of the Electron used a to discover the _ Mass of the Electron Mass of the electron is 9.11 x 10-28 g The oil drop apparatus determined the of the the mass of a _ atom Conclusions from the Study of the Electron: a) Atoms have must be the , so there to _ of the electrons b) that for have so little _ must account of the _ Conclusions from the Study of the Electron: Eugen Goldstein observed _ (or 1840 times that of an electron) J. Chadwick confirmed the _ _ nearly to a Subatomic Particles Particle Charge Mass (g) Electron (e-) -1 9.11 x 10-28 Proton (p+) _ 1.67 x 10-24 Neutron (no) _ 1.67 x 10-24 Location _ _ Thomson’s Atomic Model J. J. Thomson Thomson - . were like embedded in a charged . Ernest Rutherford’s Gold Foil Experiment - 1911 particles ( fired at a thin Particles that are ) . on the detecting _ Rutherford’s Findings Most A of the particles particles were . Conclusions: a) The is , . , _ The Rutherford Atomic Model Based on his experimental evidence: • Atom is mostly the • the , and is in the center at the . all . The Rutherford Atomic Model • is made of and • • Called the “ . the nucleus. ” Section 4.3 Distinguishing Among Atoms OBJECTIVES: Explain what makes elements and isotopes different from each other. Calculate the number of neutrons in an atom. Calculate the atomic mass of an element. Explain table. why chemists use the periodic Atomic Number Atoms are composed of protons, neutrons, and electrons • How then are atoms of one from element? Atomic Number Elements are contain different in the nucleus (smaller #) # protons = # electrons because they of _ _ Atomic Number: _ # p+ : # e- : _ Atomic Number: _ # p+ : # e- : _ _ _ Atomic Number of an element is the of _ in the nucleus of each atom of that element. Element Carbon (C) # of protons _ Atomic # (Z) _ Phosphorus (P) _ _ Gold (Au) _ _ Mass Number and nucleus of an isotope: Mass # = is the number of in the _ Atomic Number: _ Mass Number: # p+ : # e- : _ _ #n0 : _ _ Atomic Number: _ Mass Number: # p+ : _ # e- : _ #n0 : _ _ Mass Number Practice Atom p+ Oxygen 8 n0 e- Mass # _ _ _ 74 Arsenic _ 41 _ Phosphorus _ _ 15 _ Complete Symbols Contain the symbol of the element, the mass number and the atomic number. Superscript → Subscript → X Symbols Find each of these: a) number of protons _ b) number of 11 _ neutrons 23 c) number of _ electrons _ d) Atomic number e) Mass Number _ Na Symbols If an element has an atomic number of 34 and a mass number of 78, what is the: a) number of protons = b) number of neutrons = c) number of electrons = d) complete symbol Symbols If an element has 91 protons and 140 neutrons what is the a) Atomic number = b) Mass number = c) number of electrons = d) complete symbol Symbols If an element has 78 electrons and 117 neutrons what is the a) Atomic number b) Mass number c) number of protons d) complete symbol Isotopes Dalton was about of the type being . Atoms of the same element numbers of Thus, different These are called . Atomic #: Mass #: # p+: #n0: Atomic #: Mass #: # p+: #n0: _ have . . Atomic #: Mass #: # p+: #n0: Isotopes due to are atoms of the with numbers of , . Naming Isotopes We can also put the mass number after the name of the element: carbon-12 • carbon-14 • uranium-235 • Mass: Mass: Mass: are atoms of the same element having different masses, due to varying numbers of neutrons. Isotope Proton s Hydrogen–1 Electrons Neutrons (protium) Hydrogen-2 (deuterium) 1 1 0 1 1 1 Hydrogen-3 (tritium) 1 1 2 What’s the only thing that changes? Nucleus Atomic Mass How heavy is an atom of oxygen? - there are different of oxygen atoms. We want the . Based on _ of of that element in nature. Measuring Atomic Mass Measure atomic mass with the _ the mass of a Defined as atom. Each isotope has its own atomic thus we determine the , from . To calculate the average: the atomic of each isotope by it’s (expressed as a then the results. Expressed as C-12 = . _ ), Atomic Masses Atomic mass is the average of all the naturally occurring isotopes of that element. Isotope Symbol Carbon-12 12C Carbon-13 13C Carbon-14 14C Composition of the nucleus 6 protons 6 neutrons 6 protons 7 neutrons 6 protons 8 neutrons Carbon = 12.011 % in nature 98.89% 1.11% <0.01% Atomic Mass Example B-10 = 19.8% B-11 = 80.2% At. Mass = The Periodic Table: A Preview elements in which the are based on a set of Allows easy - arrangement of _ into . of the of different elements The Periodic Table: A Preview row (there are 7 of them) column Also called a _ Elements in a have similar and physical . Identified with number and “A” or “B” Draw an arrow and label a period and a group. Chapter 5 “Electrons in Atoms” Section 5.1 Models of the Atom OBJECTIVES: • Identify the inadequacies in the Rutherford atomic model. • Identify the new proposal in the Bohr model of the atom. • Describe the energies and positions of electrons according to the quantum mechanical model. Describe how the shapes of orbitals related to different sublevels differ. Ernest Rutherford’s Model Discovered _ “nucleus” Electrons surround and nucleus Like planets around the sun Atom is mostly _ Ernest Rutherford’s Model Did not explain of elements Better description of the was needed Niels Bohr’s Model Electrons move in circular , paths, or at different levels. Amount of fixed energy, ( separates one level from another. Electrons can from one level to another. ,) Bohr’s model analogous to the rungs of a ladder Electrons exist energy levels, just like you can’t stand between rungs on a ladder Bohr’s model – amount of energy required to an electron from one energy level to another The away from the , the more the electron has. The Quantum Mechanical Model In 1926, Erwin Schrodinger derived the Determines the of an States the of certain distance from the an electron a . The Quantum Mechanical Model The atom is found inside a blurry “ An area where there is a of finding an electron. Think of fan blades ” Atomic Orbitals regions where there is a of finding an electron. - arranged in sections: letters Each corresponds to a different . – Principal Quantum Number “ ”- it denotes the in which the electron is located. ~ Maximum number of electrons that can fit in an energy level is: _ How many e- in level 2? 3? level # of shapes (orbitals) Maximum electrons Starts at energy level s 1 2 1 p 3 d 5 6 10 2 f 7 14 4 3 Number of sublevels due to number of different shapes of orbitals _ _ _ _ By Energy Level First Energy Level Has orbital only electrons 1s2 Second Energy Level Has orbitals Electrons: 2 in s, 6 in p 8 total electrons By Energy Level Third energy level Has _ orbitals Electrons: 2 in s, 6 in p, and 10 in d _ 18 total electrons Fourth energy level Has _ orbitals Electrons: 2 in s, 6 in p, 10 in d, and 14 in f _ 32 total electrons Electron Configuration Sublevel # of Orbitals Available # of Electrons Available s p d 1 3 5 2 6 10 f 7 14 d= f= Section 5.2 Electron Arrangement in Atoms OBJECTIVES: • Describe how to write the electron configuration for an atom. • Explain why the actual electron configurations for some elements differ from those predicted by the aufbau principle. 7p 7s 6s 6p 5p 6d 5f 5d 4f 4d Increasing energy 5s 4p 3d 4s 3p 3s 2p 2s aufbau diagram - page 133 1s Aufbau is German for “building up” Electron Configurations… …are the way are in various orbitals around the nuclei of atoms. Three rules tell us how: Rule #1 - Aufbau Principle Electrons must the orbital with the energy _ Example: Oxygen 1s22s22p4 1s 2s 2p 1s 2s 2p Rule #2 - Pauli Exclusion Principle Orbitals can only have The 2 electrons must have electrons max spins Example: Oxygen 1s22s22p4 1s 2s 2p 1s 2s 2p Rule #3 - Hund’s Rule Orbitals of energy are each by one electron any pairing occurs Example: Oxygen 1s22s22p4 1s 2s 2p 1s 2s 2p 7p 7s 6s 6d 5f 6p 5d 4f 5p 4d Increasing energy 5s 4p 3d 4s 3p 2p 2s 1s Elec. Conf. of P? The first two electrons go into the 1s orbital Notice the opposite direction of the spins only 13 more to go... 3s 7p 7s 6s 6d 5f 6p 5d 4f 5p 4d Increasing energy 5s 4p 3d 4s 3p 3s 2p 2s 1s The next electrons go into the 2s orbital only 11 more... 7p 7s 6s 6p 5p 6d 5f 5d 4f 4d Increasing energy 5s 4p 3d 4s 3p 3s 2p 2s 1s • The next electrons go into the 2p orbital • only 5 more... 7p 7s 6s 6p 5p 6d 5f 5d 4f 4d Increasing energy 5s 4p 3d 4s 3p 3s 2p 2s 1s • The next electrons go into the 3s orbital • only 3 more... 7p 7s 6s 6p 5p 6d 5f 5d 4f 4d Increasing energy 5s 4p 3d 4s 3p 3s 2p 2s 1s Orbital notation • The last three electrons go into the 3p orbitals. They each go into separate shapes (Hund’s) • 3 unpaired electrons = 1s22s22p63s23p3 Orbitals fill in an order energy to energy. Adding electrons can change the energy of the orbital. are the absolute best situation. Orbitals fill in an order However, • • orbitals have a lower energy, and are next . Makes them more . Changes the filling order Practice Problems Write electron configurations for the following atoms 1. Li 5. P 2. N 6. Si 3. Be 7. Mg 4. C 8. Al Electron Configurations can be written in terms of noble gases To save space, configurations can be written in terms of noble gases Elec. Conf. for S? Look at noble gas before it. Example 1: Ne = 1s22s22p6 S = 1s22s22p63s23p4 Or S= [Ne] 3s23p4 Elec. Conf. for Mn? Look at noble gas before it. Example 2: Ar = 1s22s22p63s23p6 Mn = 1s22s22p63s23p64s23d5 Mn = [Ar] 4s23d5 Write the electron configurations for these elements: Titanium - 22 electrons Vanadium - 23 electrons Chromium - 24 electrons But this is not what happens!! Chromium is actually: 1s22s22p63s23p6 Why? This gives us two orbitals (the others are all still full) Half full is slightly in energy. The same principal applies to copper. Copper’s electron configuration Copper has 29 electrons so we expect: 1s22s22p63s23p6 But the actual configuration is: 1s22s22p63s23p6 This change gives one more filled orbital and one that is half filled. Remember these exceptions: Irregular configurations of Cr and Cu Chromium steals a 4s electron to make its 3d sublevel HALF FULL Copper steals a 4s electron to FILL its 3d sublevel s = spin When an electron moves, it generates a . s describes the direction an They must spin in Spin= up down There are two values of s: directions Section 5.3 Physics and the Quantum Mechanical Model OBJECTIVES: • Describe the relationship between the wavelength and frequency of light. • Identify the source of atomic emission spectra. • Explain how the frequencies of emitted light are related to changes in electron energies. • Distinguish between quantum mechanics and classical mechanics. - Page 139 “R O Frequency Increases Wavelength Longer Y G B I V” Parts of a wave Crest Origin ____________ Trough _______ Energy ______ Energy Radio Micro Infrared Ultra- XGamma waves waves . violet Rays Rays __________ __________ Frequency Frequency ___________ ___________ Wavelength Wavelength Visible Light Wavelength = Frequency = ENERGY Wavelength = Frequency = ENERGY _ _ _ _ _ _ Wavelength Table Atomic Spectrum • • – the discrete representing the frequencies of emitted by an element to each element, like fingerprints! Very useful for identifying elements When these electrons return to their energy levels, they lose energy by emitting Changing the energy Let’s look at a hydrogen atom, with only , and in the first energy level. Changing the energy Heat, electricity, or light can move the electron up to different energy levels. The electron is now said to be “ ” Changing the energy As the electron back to the ground state, it gives the energy back as Changing the energy Fall down in Each step has a different steps Ultraviolet The more energy is and the higher the Visible Infrared they fall, .