2 Overview Atomic Theory Atomic Structure Wave nature of light John DaltonLaw of Conservation of MassLaw of Definite ProportionsLaw of Multiple ProportionsErnest RutherfordRobert MillikanJ.J. ThompsonAtomic StructureProtons, neutrons, electronsAtomic numberIsotopesMass numberAverage atomic massWave nature of lightElectromagnetic SpectrumC = λvBohr ModelsPhotoelectric effectAbsorption/emissionE = hc/ λHeisenberg Uncertainty PrincipleConfigurations (orbital, electron, noble gas)Pauli Exclusion PrincipleHund’s RuleParamagnetism/diamagnetismExceptions
3 Chemistry TimeLine 2000 years of Alchemy B.C. 400 B.C. Democritus and Leucippos use the term "atomos”Chemistry TimeLine2000 years of Alchemy1500'sGeorg Bauer: systematic metallurgyParacelsus: medicinal application of minerals1600'sRobert Boyle:The Skeptical Chemist. Quantitative experimentation,identification of elements1700s'Georg Stahl: Phlogiston TheoryJoseph Priestly: Discovery of oxygenAntoine Lavoisier: The role of oxygen in combustion, law of conservation ofmass, first modern chemistry textbook
4 Chemistry timeline1800'sJoseph Proust: The law of definite proportion (composition)John Dalton: The Atomic Theory, The law of multiple proportionsJoseph Gay-Lussac: Combining volumes of gases, existence of diatomic moleculesAmadeo Avogadro: Molar volumes of gasesJons Jakob Berzelius: Relative atomic masses, modern symbols for the elementsDmitri Mendeleyev: The periodic tableJ.J. Thomson: discovery of the electronHenri Becquerel: Discovery of radioactivity1900'sRobert Millikan: Charge and mass of the electronErnest Rutherford: Existence of the nucleus, and its relative sizeMeitner & Fermi: Sustained nuclear fissionErnest Lawrence: The cyclotron and trans-uranium elements
5 The Greeks 400 BC Democritus Matter consists of small particles Called them “atomos”Idea rejected by peersNo scientific proof
6 The Greeks (cont…) Aristotle Idea endured for 2000 years All matter continuous4 elements = earth, water, air, and fireNo scientific proofIdea endured for 2000 years
7 John Dalton - 1808 School Teacher Atomic Theory All matter is composed of extremely small particles called atoms. There are different kinds called elements.Atoms of the same element are identical in size, mass, and other properties; atoms of different elements differ in size, mass, and other properties.Atoms cannot be subdivided, created, or destroyed.Atoms of different elements combine in simple, whole number ratios to form chemical compounds.In chemical reactions, atoms are combined, separated, or rearranged but never destroyed/created.
8 Laws derived from dalton Law of Conservation of MassTotal mass present before chemical reaction is same as mass after chemical reaction2H2O 2H2 + O2 If you have 10 grams of water to start, you will get 1.12 g of hydrogen and 8.88 g of oxygenLaw of Constant Composition (definite proportions)Relative numbers and kinds of atoms are constantWater is 88.8% oxygen and 11.2% hydrogen by mass no matter how much you haveLaw of Multiple ProportionsIf two elements combine to form more than one compound, the masses of the two elements are in the ratio of small whole numbersCO2 versus CO (mass ratio is 2 to 1 for oxygen)
9 J.J. Thomson British Physicist Discovered electron Cathode-ray experimentPlum pudding view of atom
10 Thompson Cathode Ray Experiment Electric current sent through gases in glass tube called cathode-ray tubeSurface of tube opposite the cathode glowed – caused by stream of particlesRay traveled from cathode to anodeCathode rays deflected by magnetic field away from negatively charged object (like a magnet)Cathode rays concluded to have negative charge
11 Robert Millikan - 1909 American Physicist Charge on each electron is sameCharge of electron is x 10-19CCalculated mass of electron as x kgOil drop experiment
12 Millikan Oil Drop Experiment Drops of oil that had picked up extra electrons allowed to fall between two electrically charged platesMeasured how voltage on plates affected rate of fallCalculated charges of drops then deduced charge of a single electron on the drops
13 Ernest RutherfordDiscovered nucleusPlanetary model of the atom
14 Rutherford Gold Foil Experiment Bombarded thin piece gold foil with alpha particles (positively charged particle 4 times mass of hydrogen atom)Expected to pass right through gold foil1 in 8000 particles deflected back toward source“As if you fired 15-inch artillery shell at a piece of tissue paper and it came back and hit you”Concluded most of atom is empty space except for a very small force within atomCalled positive bundle of matter the “nucleus”
15 Modern Atomic Theory Atom consists of proton, neutron, and electron Proton charge = +1Neutron charge = 0 (neutral)Electron charge = -1Protons and Neutrons located in nucleus99.9% of atom’s mass is in nucleusElectrons located outside the nucleus
16 Ag Element Blocks 47 Silver 107.87 Atomic number Name of the element Element SymbolAtomic mass
17 Element Blocks Atomic Number Element Symbol equal to number of protons in an atomElement SymbolFirst letter always capitalizedIf second letter exists, it is lowercase
18 Atomic mass unit (AMU)Too difficult to measure elements in “grams” so we use the atomic mass unitApproximately the mass of 1 proton or 1 neutronRelative to the carbon atom1 amu is 1/12 the mass of the carbon atom
19 isotopesIsotopes are atoms of the same element having different masses due to varying numbers of neutrons.IsotopeProtonsElectronsNeutronsNucleusHydrogen–1(protium)1Hydrogen-2(deuterium)Hydrogen-3(tritium)2
20 Composition of the nucleus Atomic MassAtomic mass is the average of all the naturally isotopes of that element.IsotopeSymbolComposition of the nucleus% in natureCarbon-1212C6 protons6 neutrons98.89%Carbon-1313C7 neutrons1.11%Carbon-1414C8 neutrons<0.01%Carbon =
21 Mass Number Mass Number = Protons + Neutrons Not found on periodic tableIsotopes have different mass numbers (due to neutrons)
22 Symbolizing ElementsC– 12Mass numberMass numberAtomic number
23 Wave-Particle Duality JJ Thomson won the Nobel prize for describing the electron as a particleHis son, George Thomson won the Nobel prize for describing the wave-like nature of the electron.The electron is a particle!The electron is an energy wave!
24 Traveling WavesMuch of what has been learned about atomic structure has come from observing the interaction of visible light and matter.
25 Wave Theory of Electron 1924 De Broglie suggested that electrons have wave properties to account for why their energy was quantized.He reasoned that the electron in the hydrogen atom was fixed in the space around the nucleus.He felt that the electron would best be represented as a standing wave.As a standing wave, each electron’s path must equal a whole number times the wavelength.
26 De BroglieThe electron propagates through space as an energy wave. To understand the atom, one must understand the behavior of electromagnetic waves.Louis deBroglie
27 Waves Wavelength, l Amplitude Frequency, n The distance for a wave to go through a complete cycle.AmplitudeHalf of the vertical distance from the top to the bottom of a wave.Frequency, nThe number of cycles that pass a point each second.
29 Waves Longer wavelength = lower frequency = lower energy Shorter wavelength = higher frequency = higher energy
30 Wavelength Frequency Relationship The SI unit of frequency (n) is the hertz, Hz1 Hz = 1 s-1Wavelength and frequency are relatedc = lnc is the speed of light, x108 m/s
31 Practice ProblemThe wavelength of an argon laser's output is nm. Calculate the frequency of this wavelength of electromagnetic radiation.c = lnConvert nm to m488 nm x (1 m / 109 nm) = 4.88 x 10-7 mThen, substitute into c = λν(4.88 x 10-7 m) (v) = 3.00 x 108 m s-1v = 6.15 x 1014 s-1 = 6.15 x 1014 Hz
32 Electromagnetic Radiation Energy in the form of transverse magnetic and electric waves.Electromagnetic SpectrumContains all forms of electromagnetic radiationVisible spectrumPortion of electromagnetic spectrum that we can see (colors)
34 Separation of Light‘White’ light is actually a blend of all visible wavelengths. They can separated using a prism.
35 Line SpectraNeils Bohr studied the spectra produced when atoms were excited in a gas discharge tube.
36 Line SpectraEach element produces its own set of characteristic lines
37 Bohr ModelBohr proposed a model of how electrons moved around the nucleus.He wanted to explain why electrons did not fall in to the nucleus.He also wanted to account for spectral lines being observed.He proposed that the energy of the electron was quantized - only occurred as specific energy levels.
38 BoHr ModelIn the Bohr model, electrons can only exist at specific energy levels (orbit).Each energy level was assigned a principal quantum number, n.Energy
39 Bohr Model The Bohr model is a ‘planetary’ type model. Each principal quantum represents a new ‘orbit’ or layer.The nucleus is at the center of the model.
40 Transitions Electron transitions involve jumps of definite amounts Of energy.
41 Absorption EmissionAbsorption – Electromagnetic radiation is absorbed by an atom causing electrons to jump to a higher energy state (excited state).Emission – Energy is released by an atom as particle of light (photon) as electrons fall back to the lower energy state (ground state).Depending on frequency of photon, different colored light may be seen
42 Particle PropertiesAlthough electromagnetic radiation has definite wave properties, it also exhibits particle properties.Photoelectric effect.First observed by Hertz and then later explained by Einstein.When light falls on Group IA metals, electrons are emitted (photoelectrons).As the light gets brighter, more electrons are emitted.The energy of the emitted electrons depends on the frequency of the light.
43 Photoelectric EffectThe energy of a photon is proportional to the frequency.(Photon energy) E= hnThe energy is inversely proportional to the wavelength (remember c = λν so v = c/λ ).E = hc /lh is Plank’s constant, x J . Sc is the speed of light, x108 m/s
44 Photon Energy ExampleDetermine the energy, in kJ/mol of a photon of blue-green light with a wavelength of 486 nm.E === x Jh cl(6.626 x J.s)(2.998 x 108 m.s-1)(4.86 x 10-7 m)
45 De Broglie Equation l = h mv l = wavelength, meters h = Plank’s constantm = mass, kgv = frequency, m/sl =hmv
46 De Broglie EquationUsing De Broglie’s equation, we can calculate the wavelength of an electron.l =hmv6.6 x kg m2 s-1(9.1 x kg)(2.2 x 106 m s-1)l == 3.3 x mThe speed of an electron had already been reportedby Bohr as 2.2 x 106 m s-1.
47 Heisenberg Uncertainty PRinciple In order to observe an electron, one would need to hit it with photons having a very short wavelength.Short wavelength photons would have a high frequency and a great deal of energy.If one were to hit an electron, it would cause the motion and the speed of the electron to change.According to Heisenberg, it is impossible to know both the position and the speed of an object precisely.
48 Quantum ModelSchrödinger developed an equation to describe the behavior and energies of electrons in atoms.His equation is similar to one used to describe electromagnetic waves.Each electron can be described in terms of its quantum numbers.
49 Quantum Numbers Energy level Orbital shape Orientation Spin Each electron in an atom has a unique set of 4 “numbers” which describe itEnergy levelOrbital shapeOrientationSpin
50 Main Energy level Principal quantum number, n n = 1, 2, 3, …… Tells the size of an orbital and largely determines its energy.n = 1, 2, 3, ……
51 Shape of the orbital Angular momentum Orbitals The number of subshells that a principal level contains. It tells the shape of the orbitals.s p d fOrbitalsAn orbital is a region within an energy level where there is a probability of finding an electronOrbital shapes are defined as the surface that contains 90% of the total electron probability.
53 Orientation of orbital Magnetic quantum number, mlDescribes the direction that the orbital projects in space.Think in terms of axes “x, y, z”
54 Spin of the electronPauli added one additional quantum number that would allow only two electrons to be in an orbital.Spin quantum number, ms.An electron can spin clockwise or counterclockwise
55 Rules Pauli exclusion principle Pauli proposed that no two electrons in an atom can have the same set of four quantum numbersUnless you live together, none of you have the exact same address
56 Other Rules Aufbau Principle Electrons are placed into orbitals, subshells, and shells in order of increasing energy
57 Other Rules Hund’s Rule The most stable arrangement of electrons in a subshell is the one in which electrons have the most number of parallel spins possible.
58 Orbital notation of electrons Graphical representation of an electron configurationOne arrow represents one electronShows spin and which orbital within a sublevelFollow all rules(Aufbau principle, two electrons in each orbital, etc. etc.)
59 Orbital notation Use atomic number as number of electrons in an atom HeBeMgSiNe
60 Magnetism Diamagnetism Paramagnetism Elements have all of their electrons spin pairedAll of an element’s subshells are completedNot affected by magnetic fieldsParamagnetismNot all electrons are spin paired in an elementMost elements are thisAffected by magnetic fields
61 Electron Configuration A list of all the electrons in an atom (or ion)Must go in order (Aufbau principle)2 electrons per orbital, maximumWe need electron configurations so that we can determine the number of electrons in the outermost energy level.These are called valence electrons.1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14… etc.
62 Electron configuration 2p4Number of electrons in the sublevelEnergy LevelSublevel
64 Periodic Table Orbitals grouped in s, p, d, and f orbitals (sharp, proximal, diffuse, and fundamental)s orbitalsd orbitalsp orbitalsf orbitals
65 Why are d and f orbitals always in lower energy levels d and f orbitals require LARGE amounts of energyIt’s better (lower in energy) to skip a sublevel that requires a large amount of energy (d and f orbtials) for one in a higher level but lower energy
66 Example: [Ne] 3s2 3p5 Noble Gas Notation A way of abbreviating long electron configurationsSince we are only concerned about the outermost electrons, we can skip to places we know are completely full (noble gases), and then finish the configurationFind the closest noble gas to the atom (or ion), WITHOUT GOING OVER the number of electrons in the atom (or ion). Write the noble gas in brackets [ ].Step 2: Find where to resume by finding the next energy level.Step 3: Resume the configuration until it’s finished.Example: [Ne] 3s2 3p5
67 Exceptions Remember d and f orbitals require LARGE amounts of energy If we can’t fill these sublevels, then the next best thing is to be HALF full (one electron in each orbital in the sublevel)There are many exceptions, but the most common ones areFor the purposes of this class, we are going to assume that ALL atoms (or ions) that end in d4 or d9 are exceptions to the rule. This may or may not be true, it just depends on the atom.
68 Exceptions d4 is one electron short of being HALF full In order to become more stable (require less energy), one of the closest s electrons will actually go into the d, making it d5 instead of d4.For example: Cr = [Ar] 4s2 3d4Since this ends exactly with a d4 it is an exception to the rule. Thus, Cr = [Ar] 4s1 3d5Remember, half full is good… and when an s loses 1, it too becomes half full!d9 works the same way