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ATOMIC THEORY In 460 B.C., a Greek philosopher, Democritus, develop the idea of atoms. He asked this question: If you break a piece of matter in half, and then break it in half again, how many breaks will you have to make before you can break it no further? Democritus thought that it ended at some point, a smallest possible bit of matter. He called these basic matter particles, atoms. Democritus (460-370 BC) • Proposed that matter is made of tiny particles and empty space • atoms = smallest part of matter • different types of atom for each type of matter People considered Aristotle's opinions very important and if Aristotle thought the atomic idea had no merit, then most other people thought the same also. For more than 2000 years nobody did anything to continue the explorations that the Greeks had started into the nature of matter. Lavoisier (1743-1794) • Made measurements of a chemical change in a sealed container • law of conservation of matter Proust (1799) • Observed that water is 11% H and 89% O. • Law of definite proportions definitions • The principle that the elements that comprise a compound are always in a certain proportion by mass. In the 1800's an English chemist, John Dalton performed experiments with various chemicals that showed that matter, indeed, seem to consist of elementary lumpy particles (atoms). http://antoine.fsu.umd.edu/chem/senese/101/ato ms/dalton.shtml Postulates 1. All matter consists of tiny particles. 2.Atoms are indestructible and unchangeable. 3.Elements are characterized by the mass of their atoms. 4.When elements react, their atoms combine in simple, whole-number ratios. J. J. Thomson (1897) • Discovered the electron and proton • Vacuum Tube Experiment cathode - anode + power source • The beam starts at the cathode. • Cathode ray tube. • CRT cathode - anode + power source • The beam is bent away from a negatively charged plate. • Like charges repel. • The beam must be negative particles. cathode - anode + power source • Positive particles were deflected from the anode. THOMSON’S MODEL In 1897, the English physicist J.J. Thomson discovered the electron and proposed a model for the structure of the atom. Thomson knew that electrons had a negative charge and thought that matter must have a positive charge. His model looked like raisins stuck on the surface of a lump of pudding. The Gold Foil Experiment “You, lab tech, just stay here and count the flashes of light. I’m going out to have a drink with the other professors” - Rutherford to Geiger “I say! It was like firing a gun at a tissue and having the bullet bounce back at you! “You’re wrong. Do you hear me? You’re wrong! I say!” -Rutherford to Thomson Experiment Conclusion • The atom is mostly empty space. • The atom contains a small, dense positive core. – Nucleus – If the nucleus was the size of a golf ball the nearest electron would be one mile away. – model RUTHERFORD’S MODEL Rutherford knew that atoms consist of a compact positively charged nucleus, around which circulate negative electrons at a relatively large distance. The nucleus occupies less than one thousand million millionth of the atomic volume, but contains almost all of the atom's mass. If an atom had the size of the earth, the nucleus would have the size of a football stadium. • The small, dense, positively charged central core of an atom. smallest same mass Comparing Subatomic Particles Particle Symbol Charge Mass (u) Mass (g) Proton P+ +1 1 1.67x10-24 Neutron n0 0 1 1.67x10-24 Electron e- -1 .00055 9.11x10-28 Atomic Number 17 Cl Chlorine 35.45 • Equals the number of protons • defines the element – all chlorine atoms have 17 protons Mass number 17 Cl Chlorine 35.45 • Protons + neutrons • Isotope have different masses – because they have a different # of neutrons • Chart mass is an average of the isotopes Mass number 17 Cl Chlorine 35.45 • What is the unit for mass? • Atomic mass unit • symbol = u • based on carbon-12 • 1p+= 1n0 = 1u Why aren’t the electrons attracted to the nucleus? BOHR’S MODEL In 1912 a Danish physicist, Niels Bohr came up with a theory that said the electrons do not spiral into the nucleus and with some rules for what does happen. RULE 1: Electrons can orbit only at certain allowed distances from the nucleus. RULE 2: Atoms radiate energy when an electron jumps from a higher-energy orbit to a lower-energy orbit. Also, an atom absorbs energy when an electron gets boosted from a low-energy orbit to a highenergy orbit. BOHR-SOMMERFIELD’S MODEL According to the Bohr-Sommerfeld model, not only do electrons travel in certain orbits but the orbits have different shapes and the orbits could tilt in the presence of a magnetic field. Orbits can appear circular or elliptical, and they can even swing back and forth through the nucleus in a straight line. QUANTUM MECHANICAL MODEL The visual concept of the atom now appeared as an electron "cloud" which surrounds a nucleus. The cloud consists of a probability distribution map which determines the most probable location of an electron. ATOM divisions NUCLEUS ELECTRON CLOUD particles particles ELECTRONS PROTONS charge charge NEGATIVE Relative mass NEUTRONS NEUTRAL POSITIVE is 1/1900 If atom is neutral ATOMIC NUMBER (Z) Relative mass 1 is ATOMIC MASS (A) N. Bohr (1885-1962) • Proposed that the electrons orbited the nucleus . • The further away the more energy was needed. • Electrons only occupy orbits of certain energy. •Carry energy not matter •Wavelength is the distance between corresponding points on the wave. •Wavelengths per second = frequency •1 wavelength/second = 1 Hertz (Hz) The Modern Model • Visible light can be split into colored light. • Each color has its own – frequency – wavelength – energy • Red light has – low frequency – long wavelength – less energy • Blue light has – high frequency – short wavelength – more energy Gamma Rays….....Pass through most substances. X rays…....Pass through soft tissue, but not bone. UV light…….....causes sunburn, stopped by ozone Light……………………………………………………..visible Infrared………………………………………..radiant heat Radiowaves…..Carry the sounds of radio stations. The Electromagnetic Spectrum Electrons & Light • When energy is put into an atom, electrons absorb the energy and become “excited”. • Electrons can absorb only fixed amounts of energy. • When excited, electrons move into a higher energy level (state). Electrons & Light • The electrons return to their original ground state. • The electrons must give off energy in the form of light. • This light has a specific frequency (and color.) • This frequency shows up as a line when seen through a prism. The Modern Model Bright Line Spectra When atoms absorb energy, from fire or electricity, the electrons give off light. •When viewed through a prism H a bright line spectra appears. Hg •This is unique for every Ne element. hydrogen The big gaps between spectral lines represent electron transitions from one energy level to another. helium The groups of fine lines indicate that electrons are jumping from energy levels that are close in energy. 3 ENERGY LEVEL 2 1 ULTRAVIOLET 4 INFRARED 7 6 5 Quantum Theory Model (Electron Cloud Model) • Heisenberg Uncertainty Principle – You can never know exactly where an electron is if you know how fast it is moving. – If you know its exact location, you can’t know how fast it’s moving. Electron Cloud Model • Electron paths are not neat orbits. • The region around a nucleus where an electron may be found is the electron cloud or orbital. • An electron cloud is associated with a certain energy level. In 1927 Heisenberg formulated an idea which agreed with test results, that no experiment can measure the position and momentum of a quantum particle simultaneously. Scientists call this the "Heisenberg uncertainty principle." This implies that as one measures the certainty of the position of a particle, the uncertainty in the momentum gets correspondingly larger. Or, with an accurate momentum measurement, the knowledge about the particle's position gets correspondingly less. Uncertainty Principle The position of a particle as well as its attributes cannot be determined with absolute precision. In association with this fact, a particle may disappear from one place and reappear elsewhere acausally. The velocity (causal indeterminism) is coupled with uncertainty of position (acausal indeterminism). These indeterminate factors are added to the deterministic velocity to obtain the motion of the electron as the changing position of its "locus of activity". • The visual concept of the atom now appeared as an electron "cloud" which surrounds a nucleus. The cloud consists of a probability distribution map which determines the most probable location of an electron. For example, if one could take a snap-shot of the location of the electron at different times and then superimpose all of the shots into one photo. • Note: Just as no map can equal a territory, no concept of an atom can possibly equal its nature. These models of the atom simply served as a way of thinking about them, albeit they contained limitations (all models do). What should a Model look like? Scientific models may not always look like the actual object. A model is an attempt to use familiar ideas to describe unfamiliar things in a visual way. This is a painting of a young woman by Pablo Picasso. Does it actually look like a young woman? Is this really an Atom? Many of the models that you have seen may look like the one below. It shows the parts and structure of the atom. Even though we do not know what an atom looks like, scientific models must be based on evidence. The model above represents the most modern version of the atom. (Artist drawing) Atomic Models Electron Cloud Model • A certain amount of electrons can exist in each energy level –1st level = 2 –2nd level = 8 –3rd level = 18 • Electrons in the outer most levels are valence electrons. Sublevels & Orbitals • Each energy level can be split into “sublevels” • Each sublevel can be divided into “orbitals”. Quantum Theory Model Orbitals One “s” orbital Three “p” orbitals Five “d” orbital Tells the amount of energy Tells the shape of orbital Is a specific volume around the nucleus where two electrons exist. 1st energy level has 1 sublevel = 1s 2nd energy level has 2 sublevels = 2s, 2p 3rd energy level has 3 sublevels = 3s, 3p and 3d 4th energy level has 4 sublevels = 4s,4p, 4d and 4f s sublevel holds maximum p sublevel holds maximum d sublevel holds maximum f sublevel holds maximum 2 e6 e10 e14 e- Placing electrons in orbitals: •orbital can hold a maximum of 2 e•orbital has the same name as the sublevel •orbitals in the “s” sublevel are “s” orbitals y 2s z x Placing electrons in orbitals: •orbital can hold a maximum of 2 e•orbital has the same name as the sublevel •orbitals in the “p” sublevel are “p” orbitals 2py 2px 2pz 1, 2, 3, 4, 5, 6 or 7 s, p, d, or f (for the 4th energy level) x, y, or z (for the p sublevel) SUBLEVEL ORBITALS ELECTRONS s 1 2 p 3 6 d 5 10 f 7 14 The most stable arrangement of electrons in sublevels and orbitals. Electron Configurations For Aluminunum 2 2 6 2 1s 2s 2p 3s 3p sublevel 1 number of electrons in the sublevel Electron Configurations For Manganese 2 2 6 2 6 2 1s 2s 2p 3s 3p 4s 3d sublevel number of electrons in the sublevel 5 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 6d 4f 5f Sublevel Filling Order 7s 7p 7d 6s 6p 6d 6f 5s 5p 5d 5f 4s 4p 4d 4f 3s 3p 3d 2s 2p 1s Try these electron configurations! 2 2s2 2p6 1s Ne = 10 2 2s2 2p6 3s1 Na = 1s 11 2 2s2 2p6 3s2 Mg = 1s 12 2 2s2 2p6 3s2 3p6 Ar = 1s 18 19K = 1s2 2s2 2p6 3s2 3p6 4s1 Déjà vu! 2 2s2 2p6 Look! It’s 1s Ne = 10 neon! 2 2s2 2p6 3s1 Na = 1s 11 2 2s2 2p6 3s2 Mg = 1s 12 2 2s2 2p6 3s2 3p6 Ar?! Ar = 1s 18 19K = 1s2 2s2 2p6 3s2 3p6 4s1 Let’s make this shorter. 2 2s2 2p6 1s Ne = 10 2 2s2 2p6 3s1 Na = 1s [Ne] 11 2 [Ne] 2 2p6 3s2 Mg = 1s 2s 12 2 2s2 2p6 3s2 3p6 Ar = 1s 18 2p6 3s2 3p6 4s1 19K = 1s2 2s2 [Ar] Noble Gas Notation 2 2s2 2p6 1s Ne = 10 1 Na = [Ne] 3s 11 2 Mg = [Ne] 3s 12 2 2s2 2p6 3s2 3p6 Ar = 1s 18 19K = [Ar] 4s1 1s2 2p6 3p6 4p6 5p6 6p6 Last Filled Sublevel for Noble Gases Try these electron configurations! 26Fe = 47Ag = 79Au = 20Ca = 40Zr = [Ar] 4s2 3d6 [Kr] 5s2 4d9 [Xe] 6s2 4f14 5d9 [Ar] 4s2 [Kr] 5s2 4d2 Principle and Angular quantum numbers ( n and l ) • Principle quantum number ( n ) - describes the SIZE of the orbital. Since the distance from of an electron from the nucleus is directly ( l ) - describes the SHAPE of the orbital. proportional to the energy of the electron (as described in the Bohr model), the principle quantum number is also a measure of the orbital. • Angular quantum number ( l ) - describes the SHAPE of the orbital. – The s orbitals are spherical ( l = 0 ). – The p orbitals are polar ( l = 1 ). – The d orbitals are clover-leaf shaped (l = 2 ). l - describes SHAPE Magnetic quantum number ( m ) • Magnetic quantum number ( m ) - describes an ORIENTATION of orbital in space. • For s orbitals (l = 0), there is only one orientation possible, so m must equal 0. • For p orbitals (l = 1), there are three possible orientations, so m can be -1, 0, or 1. • For d orbitals (l = 2), there are five possible orientations, so m can be -2, -1, 0, 1, or 2. m - describes ORIENTATION Spin quantum number (s) • Spin quantum number ( s ) - describes the SPIN or direction (clockwise or counterclockwise) in which an electron spins. • If there are two electrons in any one orbital, they will have opposite spins, that is, one will have a + spin and the other will have a - spin. • The maximum number of electron in any one orbital is two. Rules for Allowable Combinations of Quantum Numbers • The three quantum numbers ( n, l, and m ) that describe an orbital must be integers. • “ n " cannot be zero. " n " = 1, 2, 3, 4... • "l " can be any integer between zero and ( n-1). • e.g. If n = 4, l can be 0, 1, 2, or 3. • "m" can be any integer between -l and +l. • e.g. If l = 2, m can be -2, -1, 0, 1, or 2. • "s" is arbitrarily assigned as + or - , but for any one subshell (n , l , m combination), there can only be one of each. • SUMMARY • 1. De Broglie first pointed out the wave-particle duality of nature. His idea was that all particles exhibit some wave characteristics, and vice versa. • 2. The momentum of an object is the product of its mass and velocity. Wavelength varies inversely as momentum. • 3. Heisenberg's uncertainty principle concerns the process of observing an electron's position or its velocity. It is impossible to know accurately both the position and the momentum of an electron at the same time. • 4. Schrodinger developed a mathematical equation which describes the behavior of the electron as a wave. The solution set of the wave equation can be used to calculate the probability of finding an electron at a particular point. • 5. Because of the electron's high velocity, it effectively occupies all volume defined by the path through which it moves. This volume called the electron cloud. • 6. The principal quantum number (n = 1, 2, 3,..) is the number of energy level and describes the relative electron cloud size. • 7. Each energy level has as many sublevels as the principal quantum number. The second quantum number (1 = s, p, d, f...) describes the of the cloud. • 8. The third quantum number, m, describes the orientation in space of orbital. 9:10, 9 • 9. The fourth quantum number, s, describes the spin direction of an electron. • 10. Each orbital may contain a maximum of one pair of electrons. Electrons in the same orbital have opposite spins. s:12 • 11. Pauli's exclusion principle states that no two electrons in an atom have the same set of quantum numbers. • 12. Electrons occupy first the empty orbital giving the atom the lowest energy. • 13. The diagonal rule can be used to provide the correct electron configuration for most atoms. • 14. The chemist is primarily concerned with the electrons in the outer energy level. Electron dot diagrams are useful in representing the outer level electrons.