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Atoms, Molecules, and Ions Chapter 2 Naming Ionic Compounds • Positive ion is always named first • Negative ion is always named second Tips for memorizing: 1. Learn the “ate” ion 2. The “ite” has one less oxygen 3. The prefix “per” is used with “ate” to indicate one more oxygen 4. The prefix “hypo” is used with “ite” to indicate one less oxygen • Metals with more than one oxidation state (transition metals) must have a roman numeral to indicate the oxidation state • Monatomic anions end in “ide” Fundamental Chemical Laws Law of Conservation of Mass • Mass of reactants equals the mass of the products Law of Definite Proportion • Compounds have an unchanging chemical formula Law of Multiple Proportions • Sometimes two elements can come together in more than one way, forming compounds with similar, though not identical formulas Dalton’s Atomic Theory 1. Each element is composed of Atoms 2. All atoms of a given element are identical; atoms of different elements have different properties 3. Atoms of an element are not changed into different types of atoms by chemical reactions; atoms are neither created nor destroyed in chemical reactions 4. Compounds are formed when atoms of more than one element combine; a given compound always has the same relative number and kinds of atoms. Discovery of Atomic Structure 1. 2. 3. 4. JJ Thomson and the Electron Robert Millikan’s Oil Drop Radioactivity The Nuclear Atom JJ Thomson’s Cathode Rays JJ Thomson and Cathode Rays 1. Determined the charge to mass ratio of the electron 2. Reasoned that all atoms must contain electrons 3. Reasoned that all atoms must contain positive charges Millikan’s Oil Drop Milikan’s Oil Drop 1. Oil drop experiments determined charge of an electron 2. Charge info and Thomson’s charge/mass ratio allowed him to determine the mass of an electron at (9.11 x10-28 g) Behavior of radiation in an Electric Field Radioactivity 1. Gamma (g) rays – high energy light 2. Beta (b) rays – high speed electrons 3. Alpha (a) particles – nuclear particle with a +2 charge Rutherford’s Gold Foil Experiment The Nuclear Atom Rutherford’s Metal Foil Experiment • Most alpha particles pass straight through metal foil • Some particles were greatly deflected Why do you think? What could this prove? • Must have been deflected by a small, central, positive nucleus where the atom’s mass was concentrated Modern View of the Atom Nucleus: Small size, high density Contains: 1. Protons • Positively charged 2. Neutrons • No charge Electrons 1. Negatively charge 2. Source of element’s reactivity 3. Provide most of the atomic volume Subatomic particle relationships • Atomic Number Number of protons • Mass Number Number of protons + number of neutrons • Isotopes Atoms with the same number of protons (same element), but different numbers of neutrons (mass numbers) If I change the number of protons in an atom… • I get a new element If I change the number of electrons in an atom… • I get an ion If I change the number of neutrons in an atoms… • I get an isotope Symbols for the Elements Atomic Weights • Atomic Mass Scale: determined by assigning a mass of exactly 12 amu to an atom of the 12C isotope of carbon. • Average Atomic Masses: most elements exist as mixtures of isotopes so avg. atomic masses is determined by using the masses of its various isotopes and their relative abundances. Mass Spectrometer • Instrument used to measure the precise masses and relative amounts of atomic and molecular ions The Periodic Table • Horizontal Row = “period” or “series” • Vertical Column = “group” or “family” Group 1A – Alkali metals Group 2A – Alkaline earth metals Group 7A – Halogens Group 8A – Noble gases Group B – Transition Metals Metals, Metalloids, Nonmetals?? METALS a) b) c) d) e) Left side of periodic table High electrical conductivity High Luster (very shiny when clean) Ductile (drawn into wires) Solid @ room temperature (except Hg) NON-METALS a) b) c) d) Upper right corner of periodic table Non-lustrous (not shiny) Poor conductors of electricity Some are gases at room temp.(O,N) --others brittle solids (S) Semi-Metals (metalloids) • Divides metals from nonmetals • properties intermediate between metals and nonmetals • Ex: Si, Ge, ... • “stair case” on right Side of periodic table Valence Electrons • Electrons in highest occupied principle energy level • Electrons available for bonding • For Representative Elements: –Number of valence electrons in an atom = the group number of that atom has one valence electron. has two valence electrons. has three valence electrons. And So On…. has seven valence electrons. has eight valence electrons. diagrams that show the valence electrons as dots. Since only the valence electrons are involved in bonding, this is very useful in predicting which atoms will combine chemically. 6 3 4 7 X 2 1 5 8 The X represents the symbol for the element. The dots are placed around the symbol in the order shown above. The Octet Rule • When forming compounds, atoms tend to achieve the electron configuration of a noble gas. (all 8 electrons in outer shell) • Done by losing, gaining, or sharing electrons Molecules and Molecular Compounds Covalent Bonding: • Sharing of electrons • Typically between nonmetals and/or large molecules Diatomic Molecule – a form of covalent bonding in which two identical atoms are bonded together (H2,N2,O2,Cl2,F2,Br2,I2) Mono- 1 • EXAMPLES: Di- 2 Tri- 3 • CO2 Tetra- 4 Penta- 5 Hexa- 6 Hepta- 7 Octa- 8 Nona- 9 Deca- 10 • Carbon Dioxide • P4O6 • Tetraphosphorus Hexoxide Formulas for Nonmetal Compounds (Covalent Compounds) 1. Write correct symbols for the elements. 2. Use prefixes to add the correct subscript. Example: dinitrogen tetroxide N2O4 Representing Molecules • Chemical Formulas: H2O , CH4 • Structural Formulas: – Bonds represented by lines – Ball and Stick – Space Filling • Empirical Formulas: only gives the relative number of atoms of each type in a molecule (smallest possible whole number ratio of atoms) Ions and Ionic Compounds • Cation: Positive ions formed by the loss of electrons • Anion: Negative ions formed by gaining electrons Ionic Bonding 1. Formed by attraction between oppositely charge ions (electrons are transferred) 2. Forms ionic solids (salts) 3. Can be monatomic (one atom) or polyatomic (more than one atom) 4. Typically between a metal and a nonmetal Something to Ponder… • How does a recipe that calls for 4 cups of tomato sauce and 1 teaspoon of salt and the following formula CaCl2 relate to one another? • 1 atom of Calcium for every 2 atoms of Chloride Naming Compounds (Ionic Compounds) 1. Name the cation If it’s a transition metal, give it roman numerals for its correct charge 2. Name the anion • • • • • • • • Examples: NaBr = Sodium Bromide CuO = Copper Oxide ?? Copper (II) Oxide Na2SO4 = Sodium Sulfate Name the following with yourIron Group: (II) Bromide 1. Na2O 9. FeBr2 10. Cu2O 11.CaCl2 2. LiBr 3. SnO2 Tin (IV) Oxide 4. Cu3P2 Copper (II) Phosphide 5. Sr3P2 12.AlP 6. BaO 13.FeCl3 7. PbBr4 Lead (IV) Bromide 14.MgBr2 8. Cr2O3 15.AlCl3 Chromium (III) Oxide Writing Formulas (Ionic Compounds) 1. Write the Symbol and Charge of the cation 2. Write the Symbol and Charge of the anion 3. Add needed subscripts to balance the charges Find least common multiple (hard way) Use crisscross method (easy way) 3+ Fe 2O Find the least common multiple. In this case it is 6. You must have 2 Fe and 3 O to balance the charges. Fe2O3 3+ Fe 2O The easy way to do this is to “crisscross” the charges and make them subscripts. Fe2O3 Note: Always remember to reduce the subscripts to lowest terms. Try a Couple: • • • • • • • Calcium Nitride Ca3N2 Lithium Oxide Li2O Barium Sulfide BaS Why isn’t it Ba2S2 ?? Formulas w/ Polyatomic Ions: 1. Same steps as with monatomic ions 2. must use parenthesis to enclose polyatomic ion if more than one is needed Example: Calcium Nitrate Ca 2+ NO3 1 Ca(NO3)2 Naming Acids • Binary Acids (2 elements – hydrogen + one other) – Prefix “Hydro” + root of 2nd element + “ic” • Oxyacids – If the acid contains an anion whose name ends in “ate”: use the root of anion name and an “ic” ending – If acid contains an anion whose name ends in “ite”: use the root of the anion name and an “ous” ending Electronic Structure of Atoms Chapter 6 Light • Made up of electromagnetic radiation. • Waves of electric and magnetic fields at right angles to each other. Parts of a wave Wavelength l Frequency (n) = number of cycles in one second Measured in hertz 1 hertz = 1 Frequency = n Kinds of EM waves • There are many different EM waves • different l and n • Visible Light is only the part our eyes can detect. (colors of the rainbow) • Greater wavelength means, smaller frequency Gamma Rays MicroX-Rays UV Radio Infrared wave The speed of light, c • in a vacuum is 2.998 x 108 m/s • c = 3.0 x 108 m/s • c = ln Examples What is the wavelength of light with a frequency 5.89 x 1014 Hz? l = c v = 3.0 x 108 m/s 5.89 x 1014 Hz = 5.09 x 10-7 m = 509 nm (green light) What is the frequency of blue light with a wavelength of 484 nm? v= c l = 3.0 x 108 m/s 484 x 109 m = 6.20 x 1014 Hz Planck and the Quantum Theory • Energy is gained or lost in whole number multiples (n) of the quantity hv. (Similar to energy required to go up stairs vs. up a ramp) • Frequency = v • Planck’s constant = h = 6.63 x 10-34 J-s • Planck found discovered that Energy is transferred to matter in “energy packets” called a quantum (hv) DE = nhn Einstein, the Photoelectric Effect, and Photons • EM radiation is quantized a stream of particles -- “photons” • Ephoton = hn = hc/l • Combine this with E = mc2 • You get the apparent mass of a photon. m = h / (lc) Is light a Wave or does it consist of particles? • Both… • Macroscopically like a wave, • But consists of a collection of photons that we only see at the atomic level. • called The Wave-Particle Duality (Like describing an entire beach and then beginning to examine the grains of sand.) Examples • Calculate the energy of one photon of yellow light whose wavelength is 589nm 1. Find the frequency • 5.09 x 1014 s-1 2. Then use Plank’s equation to find E • 3.37 x 10-19 J Matter as a wave • Using the velocity (v) instead of the frequency (n) we get: • De Broglie’s equation l = h/mv • Can calculate the wavelength of an object. Line Spectra • Spectrum = the range of frequencies present in light • Continuous Spectrum = contains all wavelengths of light. (white light… can be broken down into “rainbow”) • Line Spectrum = contains only specific wavelengths of light. Hydrogen spectrum • Emission spectrum because these are the colors it gives off or emits. • Called a bright line emission spectrum. • There are just a few discrete lines showing 656 nm 434 nm 410 nm 486 nm Bright Line Spectra • Excited electrons return to lower NRG states • NRG is emitted in the form of a photon of definite wavelength. • Definite change in energy corresponds to: – Definite frequency – Definite wavelength • Use DE = hn = hc / l • Only certain energies are possible within any atom. Niels Bohr • Developed the Quantum Model • Described the atom like a solar system • Electrons attracted to (+) nucleus because of their (-) charge • Electrons didn’t fall into nucleus because they were moving around Bohr’s atom • Found only certain NRGs were allowed; called them NRG levels. • Putting NRG into atom moves electron away from the nucleus (ground state excited state) • When e- returns to ground state, it gives off light of a certain NRG The Bohr Atom n=4 n=3 n=2 n=1 Available NRG levels E = -2.178 x 10-18 J (Z2 / n2 ) • n = quantum number (NRG level) • Z = nuclear charge (+1 for Hydrogen) • J = energy in joules • The more negative the NRG is, the more stable the atom will be. change in Energy • When the electron moves from one energy level to another: • DE = Efinal - Einitial DE = -2.178 x 10-18J [(1/ nf2)– (1/ ni2)] l = hc / DE Shortcomings of Bohr Model • Only works for Hydrogen atoms • Electrons don’t move in circular orbits • The quantization of energy is right, but not because they are circling like planets • Questions Bohr couldn’t answer: Why are e- confined to only certain energy levels? Why don’t e- eventually spiral and crash into the nucleus? The Quantum Mechanical Model • New approach that viewed electron as a standing wave of NRG • Standing waves don’t propagate through space • Standing waves are fixed at both ends (similar to vibrations of a stringed instrument) What’s possible? • You can only have a standing wave if you have complete waves. • There are only certain allowed waves. • In the atom there are certain allowed waves called electrons. • 1925 Erwin Schroedinger described the wave function of the electron. “The Schroedinger Equation” • Much math but what is important are the solutions. Schroedinger’s Equation 2x2 22 • • • • • • 2y2 22 2z2 22 82m h2 (E V) = 0 The wave function, is a F(x, y, z) Solutions to the equation are called orbitals. These are not Bohr orbits. Each solution is tied to a certain energy. These are the energy levels. Many strange and seemingly impossible behaviors occur when the electron is treated as a wave! Orbitals • Orbitals are not circular orbits for electrons • Orbitals are areas of probability for locating electrons There is a limit to what we can know… • about how the electron is moving or how it gets from one energy level to another. • about both the position and the momentum of an object. • The Heisenberg Uncertainty Principle “we cannot know the exact location and exact momentum of an electron at the same time.” Quantum Mechanical Model and Quantum Numbers • Note: A quantum mechanical orbital is not the same as a Bohr orbit because the motion of the electron in an atom cannot be precisely measured or tracked. (Heisenberg uncertainty Principle) • There are 4 quantum numbers to describe the “location” of an electron. (sort of like how a zip code works) Principal Quantum Number (n) • Indicates probable distance from the nucleus (old Bohr orbitals) • Gives the size and energy of the orbital • Has integer values >0 • According to the periodic table, what would the highest principal quantum number be? Angular Momentum Quantum (l ) • Gives the shape of the orbital (more detail to come) • Integral values from 0 to (n-1) for each principal quantum number (n) Value of l 0 1 2 3 4 Letter used for shape* s p d f g *letters s, p, d, f come from the words sharp, principal, di and fundamental, which were used to describe certain fea of spectra before quantum mechanics was developed. Magnetic Quantum Number (ml ) • Relates to the orientation of the orbital in space relative to the other orbitals. (It tells you if the orbital will be on the x, y or z axis.) • Integral values from l to –l including 0. n l 1 2 3 4 0 0 1 0 1 Orbital designation 1s 2s 2p 3s 3p ml 0 0 -1, 0, 1 0 -1, 0, 1 # of orbitals 1 1 3 1 3 2 0 1 2 3 3d 4s 4p 4d 4f -2, -1, 0, 1, 2 0 -1, 0, 1 -2, -1, 0, 1, 2 -3, -2, -1, 0, 1, 2, 3 5 1 3 5 7 Important Observations 1. The shell w/ quantum #n will have exactly n subshells. 2. Each subshell has a specific number of orbitals. Each orbital corresponds to a different allowed value of ml. For a given value of l, there are 2l + 1 allowed values of ml. 3. The total number of orbitals in a shell is n2. The resulting number of orbitals for the shells – 1, 4, 9, 16 – is related to a pattern seen in the periodic table… We see the number of elements in the table – 2, 8, 18, 32 – equal twice these numbers… S orbitals n=1 n=2 n=3 P orbitals At another energy level the solutions are “dumbell” shaped. There are 3 possible solutions for this energy leve P Orbitals All 3 p orbitals may exist at the same time. d orbitals At another energy we get “flower” shaped orbitals for a solution. All 5 may exist at the same time F orbitals And finally, at another energy, 7 f orbitals are the solution. Orbital Energies • All orbitals with the same value of n have the same energy • The lowest energy state is called the “ground state” • When the atom absorbs energy, electrons may move to higher energy orbitals – “excited state” Electron Spin Quantum Number (ms ) • An individual orbital can hold only 2 electrons • Electrons must have opposite spins (why important?) • Spin can have two values +½ or –½ Pauli Exclusion Principle “in a given atom, no two electrons can have the same set of four quantum numbers” What this means for the atom? • Each atomic sub-orbital may contain a maximum of 2 electrons • Those electrons must have opposite spins Increasing energy 7s 6s 5s 4s 3s 2s 1s 7p 6p 6d 5p 4d 4p 5d 3d 3p 2p Helium with 2 electrons 5f 4f Increasing energy 7s 6s 5s 4s 3s 2s 1s 7p 6p 6d 5p 4d 4p 5d 3d 3p 2p Li with 3 electrons 5f 4f Increasing energy 7s 6s 5s 4s 3s 2s 1s 7p 6p 6d 5p 4d 4p 5d 3d 3p 2p Boron with 5 electrons 5f 4f 2 more important rules: • Aufbau Principle – electrons enter orbitals of lowest energy first. • Hund’s Rule -- When electrons occupy orbitals of equal energy, one electron enters each orbital before they pair. For Example: 2s 2p After the s sublevel gets two electrons, three electrons enter the p orbitals before they pair. Increasing energy 7s 6s 5s 4s 3s 2s 1s 7p 6p 6d 5p 4d 4p 3p 2p 5d 3d 5f 4f Electron Configuratoin p s d f 3 QUESTIONS TO ASK • What Row? –(principle energy level) • What section? –(type of sub-orbital) • What seat? –(how many electrons in that sub-orbital) Example 1: Write the electron configuration for nitrogen. 7N 2 2 3 1s 2s 2p Example Write the electron 2: configuration for Fe. 26Fe 2 2 6 2 6 2 6 Condensed Electron Configurations • Put the symbol for the Noble gas from the previous principal energy level, then add the electron configuration after that point. • Example 1 for Nitrogen: [He] 2s22p3 • Example 2 for Iron: • [Ar] 4s23d6 Nuclear Chemistry Chapter 21 Nuclear Radiation Natural Radioactivity Nuclear Equations Producing Radioactive Isotopes Half-Life Nuclear Fission and Fusion Subatomic Particles • Protons- plus charge In the nucleus • Neutrons- neutral • Electrons - negative charge Outside the nucleus Radiation • Radiation comes from the nucleus of an atom. • Unstable nucleus emits a particle or energy a alpha b beta g gamma Composition of Stable Nuclei Notice the nucleus packs more neutrons in the nucleus (compared to protons) as the size of the nucleus increases. 200Hg 80 Number of Neutrons ---> All elements from #83 and beyond are naturally radioactive. 120Sn 50 90Zr 40 20 Number of Protons -----> 83 Radiation Protection • Shielding alpha – paper, clothing beta – lab coat, gloves gamma- lead, thick concrete • Limit time exposed • Keep distance from source Radiation Protection Radioactive particles and rays vary greatly in penetrating power. Alpha Particle a • Same as a helium nucleus 4 2 He or a • Two protons • Two neutrons • +2 charge • Relative penetrating power = 1 Alpha decay Try this. • Write the nuclear reaction for alpha decay of 185Au. 185Au 79 181 Ir + 4 77 Atomic mass decreases by 4, Atomic No. decreases by 2 because it lost 2 protons and 2 neutrons (a) a 2 Beta Particle b An electron emitted from the nucleus 0 e or b 1 A neutron in the nucleus breaks down 1 1 n P 0 1 0 + e -1 Beta Particle b • Charge of -1 • Relative Penetrating Power = 100 • Essentially they are fast moving electrons Beta decay 234Th 90 234Pa 91 + 0e 1 beta particle b Atomic mass stays the same, at. no. goes up by 1 because a neutron changes into a proton. Try this Write the nuclear equation for the beta emitter Co 60. 60Co 27 60Ni 28 + 0 e 1 Gamma (g) Radiation • Pure radiation - high energy wave (photons) • No mass, NOT shown in the equation • Almost all radioactive changes involve some amount of gamma rays being released • Charge = 0 • Mass = 0 • Relative Penetrating Power = 10,000 Gamma radiation No change in atomic or mass number 11B 5 11B 5 + 0g 0 boron atom in a high-energy state Unnecessary to write a nuclear reaction for this. Two other common types of Radioactive Decay • Positron - same mass as an electron, but opposite charge 0 e +1 - has a very short life: annihilated as soon as it collides with an electron Write a Nuclear reaction for C-11 positron emission: 11 11 0 C B + e 6 5 +1 At. No. decreases by 1 because a proton is converted to a neutron within the nucleus. • Electron Capture - an electron from the electron cloud is captured by the nucleus 0 e 1 Write a nuclear reaction for Rb-81 undergoing electron capture. 81 Rb 37 + 0 e 1 81 Kr 36 At. no. decreases by 1 because the effect is same as a positron, a proton is converted into a neutron. Balancing Nuclear Equations • In the reactants and products Atomic numbers must balance, Mass numbers must balance, and Charges must balance Half-Life of a Radioisotope The time for the radiation level to fall (decay) to one-half its initial value decay curve initial 1 half-life 8 mg 2 4 mg 3 2 mg 1 mg Examples of Half-Life Isotope C-15 Ra-224 Ra-223 I-125 C-14 U-235 Half life 2.4 sec 3.6 days 12 days 60 days 5700 years 710 000 000 years Learning Check The half life of I-123 is 13 hr. How much of a 64 mg sample of I-123 is left after 26 hours? 64/2 = 32 mg remaining in 13 hr. 32/2 = 16 mg remaining after 26 hr. Answer: 16 mg Producing Radioactive Isotopes Bombardment of atoms produces radioisotopes = 60 59Co + 27 = 60 1n 56Mn 0 25 = 27 cobalt neutron atom particle + 4H e 2 = 27 manganese radioisotope alpha Learning Check What radioactive isotope is produced in the following bombardment of boron? 10B 5 + 4He 2 ? + 0 1n Solution What radioactive isotope is produced in the following bombardment of boron? 10B 5 + 4He 2 13N 7 nitrogen radioisotope + 1n 0 Nuclear Fission Fission large nuclei break up 235U 92 + 1n 0 139Ba 56 + 94Kr 36 + 3 1n Energy + 0 Fission Nuclear Fusion Fusion small nuclei combine 2H 1 + 3H 4He 1 2 + 1n + 0 Occurs in the sun and other stars Energy Learning Check Indicate if each of the following are (1)Fission &/or (2) Fusion A. B. C. D. Nucleus splits Large amounts of energy released Small nuclei form larger nuclei Hydrogen nuclei react Energy Solution Indicate if each of the following are (1)Fission (2) Fusion A. 1 Nucleus splits B. 1 + 2 Large amounts of energy released C. 2 Small nuclei form larger nuclei D. 2 Hydrogen nuclei react