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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 Democritus First to suggest the existence of atoms (from the Greek word “atomos”) He believed that atoms were indivisible and indestructible. Dalton’s Atomic Theory 1) All elements are composed of tiny indivisible particles called atoms. John Dalton (1766 – 1844) 2) Atoms of the same element are identical. --Atoms of any one element are different from those of any other element. Dalton’s Atomic Theory 3) Atoms of different elements combine in whole-number ratios to form compounds. 4) In chemical reactions, atoms are combined, separated, or rearranged – but never changed into atoms of another element. Sizing up the Atom 100,000,000 atoms = 1 cm 1,000,000 atoms = width of hair Can be observed with scanning tunneling (electron) microscopes 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 three subatomic particles: Electrons Protons Neutrons Discovery of the Electron J.J. Thomson used a cathode ray tube to discover the negatively charged electron Mass of the Electron Mass of the electron is 9.11 x 10-28 g The oil drop apparatus Robert Millikan determined the mass of the electron: 1/1840 the mass of a hydrogen atom Conclusions from the Study of the Electron: a) Atoms have no charge, so there must be positive particles to balance the negative charge of the electrons b) Electrons have so little mass that other particles must account for most of the mass Conclusions from the Study of the Electron: Eugen Goldstein observed positive proton Mass of 1 (or 1840 times that of an electron) James Chadwick confirmed the neutral neutron Mass nearly equal to a proton Subatomic Particles Particle Charge Mass (g) Location Electron (e-) -1 9.11 x 10-28 Electron cloud Proton (p+) +1 1.67 x 10-24 Nucleus Neutron (no) 0 1.67 x 10-24 Nucleus Thomson’s Atomic Model J. J. Thomson Thomson - plum pudding model. Electrons were like plums embedded in a positively charged pudding. Ernest Rutherford’s Gold Foil Experiment - 1911 Alpha particles (helium nuclei) fired at a thin gold foil. Particles that hit on the detecting screen are recorded Rutherford’s Findings Most of the particles passed right through A few particles were deflected. Conclusions: a) The nucleus is small, dense, and, positively charged The Rutherford Atomic Model Based on his experimental evidence: • Atom is mostly empty space. • All the positive charge, and almost all the mass is in the center at the nucleus. The Rutherford Atomic Model • Nucleus is made of protons and neutrons • Electrons surround the nucleus. • Called the “nuclear model” 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 identical protons, neutrons, and electrons • How then are atoms of one element different from another element? Atomic Number Elements are different because they contain different numbers of PROTONS Atomic number - number of protons in the nucleus (smaller #) # protons = # electrons Atomic Number: 35 # p+ : 35 # e- : 35 Atomic Number: 53 # p+ : 53 # e- : 53 Atomic Number Atomic number (Z) of an element is the number of protons in the nucleus of each atom of that element. Element # of protons Atomic # (Z) Carbon (C) 6 Phosphorus (P) 15 15 Gold (Au) 79 79 6 Mass Number Mass number is the number of protons and neutrons in the nucleus of an isotope: Mass # = p+ + n0 Atomic Number: 35 Mass Number: 79.9 # p+ : 35 # e- : 35 #n0 : 45 Atomic Number: 53 Mass Number: 127 # p+ : 53 # e- : 53 #n0 : 74 Mass Number Practice Atom p+ n0 e- Mass # Oxygen 8 8 8 Arsenic 33 Phosphorus 15 41 16 16 33 74 15 31 Complete Symbols Contain the symbol of the element, the mass number and the atomic number. Atomic Superscript → number Subscript → X Mass number Symbols Find each of these: a) number of protons 11 b) number of 11 12 neutrons 23 c) number of 11 electrons d) Atomic number 11 e) Mass Number 23 Na Symbols If an element has an atomic number of 34 and a mass number of 78, what is the: a) number of protons = 34 b) number of neutrons = 43 c) number of electrons = 34 d) complete symbol 34 X 78 Symbols If an element has 91 protons and 140 neutrons what is the a) Atomic number = 91 b) Mass number = 131 c) number of electrons = 91 d) complete symbol Symbols If an element has 78 electrons and 117 neutrons what is the 78 a) Atomic number 195 b) Mass number c) number of protons 78 d) complete symbol Isotopes Dalton was wrong about all elements of the same type being identical Atoms of the same element can have different numbers of neutrons. Thus, different mass numbers. These are called isotopes. Atomic #: 6 Mass #: 12 # p+: 6 #n0: 6 Atomic #: 6 Mass #: 13 # p+: 6 #n0: 7 Atomic #: 6 Mass #: 14 # p+: 6 #n0: 8 Isotopes Isotopes are atoms of the same element with different masses, due to varying numbers of neutrons. Naming Isotopes We can also put the mass number after the name of the element: carbon-12 • carbon-14 • uranium-235 • Mass: 12 Mass: 14 Mass: 235 Isotopes 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 Nucleus What’s the only thing that changes? # of neutrons Atomic Mass How heavy is an atom of oxygen? Depends - there are different masses of oxygen atoms. We want the average atomic mass. Based on abundance (percentage) of each variety of that element in nature. Measuring Atomic Mass Measure atomic mass with the Atomic Mass Unit (amu) Defined as one-twelfth the mass of a carbon-12 atom. Each isotope has its own atomic mass, thus we determine the average from percent abundance. To calculate the average: Multiply the atomic mass of each isotope by it’s abundance (expressed as a decimal), then add the results. Expressed as amu. C-12 = 12 amu. 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 = (10.0)(.198) + (11.0)(.802) = 10.8 amu The Periodic Table: A Preview Periodic table - arrangement of elements in which the elements are separated into groups based on a set of repeating properties. Allows easy comparison of the properties of different elements The Periodic Table: A Preview Period - horizontal row (there are 7 of them) Group - vertical column Also called a family Elements in a group have similar chemical and physical properties Identified with number and “A” or “B” Draw an arrow and label a period and a group. Group Period 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 dense positive “nucleus” Electrons surround and orbit nucleus Like planets around the sun Atom is mostly empty space Ernest Rutherford’s Model Did not explain chemical properties of elements Better description of the electron behavior was needed Niels Bohr’s Model Electrons move in specific circular paths, or orbits, at different levels. Amount of fixed energy, (energy levels,) separates one level from another. Electrons can jump from one level to another. Bohr’s model Energy level analogous to the rungs of a ladder Electrons cannot exist between energy levels, just like you can’t stand between rungs on a ladder Bohr’s model Quantum – amount of energy required to move an electron from one energy level to another The further away from the nucleus, the more energy the electron has. The Quantum Mechanical Model In 1926, Erwin Schrodinger derived the quantum mechanical model Determines the energy of an electron States the probability of finding an electron a certain distance from the nucleus. The Quantum Mechanical Model The atom is found inside a blurry “electron cloud” An area where there is a chance of finding an electron. Think of fan blades Atomic Orbitals Atomic orbitals – regions where there is a high probability of finding an electron. Sublevels- arranged in sections: letters s, p, d, and f Each sublevel corresponds to a different shape. Principal Quantum Number “n”- it denotes the energy level in which the electron is located. ~ 1, 2, 3, etc. Maximum number of electrons that can fit in an energy level is: 2n2 How many e- in level 2? 3? # 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 S-orbital P-orbitals D-orbitals F-orbitals By Energy Level First Energy Level Has only s orbital only 2 electrons 1s2 Second Energy Level Has s and p orbitals Electrons: 2 in s, 6 in p 2s22p6 8 total electrons By Energy Level Third energy level Has s, p, and d orbitals Electrons: 2 in s, 6 in p, and 10 in d 3s23p63d10 18 total electrons Fourth energy level Has s, p, d, and f orbitals Electrons: 2 in s, 6 in p, 10 in d, and 14 in f 4s24p64d104f14 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 = s level – 1 f = s level – 2 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 electrons are arranged in various orbitals around the nuclei of atoms. Three rules tell us how: Rule #1 - Aufbau Principle Electrons must occupy the orbital with the lowest energy first Example: Oxygen 1s22s22p4 1s 2s 2p 1s 2s 2p Rule #2 - Pauli Exclusion Principle Orbitals can only have two electrons max The 2 electrons must have opposite spins Example: Oxygen 1s22s22p4 1s 2s 2p 1s 2s 2p Rule #3 - Hund’s Rule Orbitals of equal energy are each occupied by one electron before 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 Lowest energy to higher energy. Adding electrons can change the energy of the orbital. Full orbitals are the absolute best situation. Orbitals fill in an order However, • • half filled orbitals have a lower energy, and are next best. Makes them more stable. Changes the filling order Practice Problems Write electron configurations for the following atoms 1. 2. 3. 4. Li N Be C 1s22s1 1s22s22p3 1s22s2 1s22s22p2 5. 6. 7. 8. P Si Mg Al 1s22s22p63s23p3 1s22s22p63s23p2 1s22s22p63s2 1s22s22p63s1 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 1s22s22p63s23p64s23d2 1s22s22p63s23p64s23d3 Chromium - 24 electrons 2 2 6 2 6 2 4 1s 2s 2p 3s 3p 4s 3d (expected) But this is not what happens!! Chromium is actually: 1s22s22p63s23p64s13d5 Why? This gives us two half filled orbitals (the others are all still full) Half full is slightly lower in energy. The same principal applies to copper. Copper’s electron configuration Copper has 29 electrons so we expect: 1s22s22p63s23p64s23d9 But the actual configuration is: 1s22s22p63s23p64s13d10 This change gives one more filled orbital and one that is half filled. Remember these exceptions: d4, d9 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 magnetic field. s describes the direction an electron spins They must spin in opposite directions Spin= up down There are two values of s: +1/2 and -1/2 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 Wavelength Origin Amplitude Trough Low Energy High Energy Radio Micro Infrared waves waves . Low Frequency Long Wavelength Ultra- Xviolet Rays Visible Light Gamma Rays High Frequency Short Wavelength Long Wavelength = Low Frequency = Low ENERGY Short Wavelength = High Frequency = High ENERGY Wavelength Table Atomic Spectrum • • Atomic Emission Spectrum – the discrete lines representing the frequencies of light emitted by an element Unique to each element, like fingerprints! Very useful for identifying elements When these electrons return to their lower energy levels, they lose energy by emitting light Changing the energy Let’s look at a hydrogen atom, with only one electron, 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 “excited” Changing the energy As the electron falls back to the ground state, it gives the energy back as light Changing the energy Fall down in specific steps Each step has a different energy Ultraviolet Visible The further they fall, more energy is released and the higher the frequency. Infrared