Chemistry Timeline #1 B.C. 400 B.C. Demokritos and Leucipposuse the term "atomos” 2000 years of Alchemy • 1500's • Georg Bauer: systematic metallurgy • Paracelsus: medicinal application of minerals 1600's Robert Boyle:The Skeptical Chemist.Quantitative experimentation, identification of elements • 1700s' • Georg Stahl: Phlogiston Theory • Joseph Priestly: Discovery of oxygen • Antoine Lavoisier: The role of oxygen in combustion, law of conservation of • mass, first modern chemistry textbook
Chemistry Timeline #2 • 1800's • Joseph Proust: The law of definite proportion (composition) • John Dalton: The Atomic Theory, The law of multiple proportions • Joseph Gay-Lussac: Combining volumes of gases, existence of diatomic molecules • Amadeo Avogadro: Molar volumes of gases • Jons Jakob Berzelius: Relative atomic masses,modern symbols for the elements • Dmitri Mendeleyev: The periodic table • J.J. Thomson: discovery of the electron • Henri Becquerel: Discovery of radioactivity • 1900's • Robert Millikan: Charge and mass of the electron • Ernest Rutherford: Existence of the nucleus, and its relative size • Meitner & Fermi: Sustained nuclear fission • Ernest Lawrence: The cyclotron and trans-uranium elements
Laws • Conservation of Mass • Law of Definite Proportion – • compounds have a constant composition. • They react in specific ratios by mass. • Multiple Proportions- • When two elements form more than one compound, the ratios of the masses of the second element that combine with one gram of the first can be reduced to small whole numbers.
Proof • Mercury has two oxides. • One is 96.2 % mercury by mass, the other is 92.6 % mercury by mass. • Show that these compounds follow the law of multiple proportion. • Speculate on the formula of the two oxides.
Dalton’s Atomic Theory (1808) • All matter is composed of extremely small particles called atoms • Atoms of a given element are identical in size, mass, and other properties; atoms of different elements differ in size, mass, and other properties John Dalton • 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
Modern Atomic Theory Several changes have been made to Dalton’s theory. Dalton said: Atoms of a given element are identical in size, mass, and other properties; atoms of different elements differ in size, mass, and other properties Modern theory states: Atoms of an element have a characteristic average mass which is unique to that element.
Modern Atomic Theory #2 Dalton said: Atoms cannot be subdivided, created, or destroyed Modern theory states: Atoms cannot be subdivided, created, or destroyed in ordinary chemical reactions. However, these changes CAN occur in nuclear reactions!
The Atomic Scale • Most of the mass of the atom is in the nucleus (protons and neutrons) • Electrons are found outside of the nucleus (the electron cloud) • Most of the volume of the atom is empty space “q” is a particle called a “quark”
About Quarks… Protons and neutrons are NOT fundamental particles. Protons are made of two “up” quarks and one “down” quark. Neutrons are made of one “up” quark and two “down” quarks. Quarks are held together by “gluons”
Isotopes Isotopes are atoms of the same element having different masses due to varying numbers of neutrons.
Atomic Masses Atomic mass is the average of all the naturally isotopes of that element. Carbon = 12.011
Molecules Two or more atoms of the same or different elements, covalently bonded together. Molecules are discrete structures, and their formulas represent each atom present in the molecule. Benzene, C6H6
Covalent Network Substances Covalent network substances have covalently bonded atoms, but do not have discrete formulas. Why Not?? Graphite Diamond
Ions • Cation: A positive ion • Mg2+, NH4+ • Anion: A negative ion • Cl-, SO42- • Ionic Bonding: Force of attraction between oppositely charged ions. • Ionic compounds form crystals, so their formulas are written empirically (lowest whole number ratio of ions).
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Discovery of the Electron In 1897, J.J. Thomson used a cathode ray tube to deduce the presence of a negatively charged particle. Cathode ray tubes pass electricity through a gas that is contained at a very low pressure.
Thomson’s Atomic Model Thomson believed that the electrons were like plums embedded in a positively charged “pudding,” thus it was called the “plum pudding” model.
Rutherford’s Gold Foil Experiment • Alpha particles are helium nuclei • Particles were fired at a thin sheet of gold foil • Particle hits on the detecting screen (film) are recorded
The Puzzle of the Atom • Protons and electrons are attracted to each other because of opposite charges • Electrically charged particles moving in a curved path give off energy • Despite these facts, atoms don’t collapse
Electromagnetic radiation propagates through space as a wave moving at the speed of light. c = C = speed of light, a constant (3.00 x 108 m/s) = frequency, in units of hertz (hz, sec-1) = wavelength, in meters
Long Wavelength = Low Frequency = Low ENERGY Wavelength Table Short Wavelength = High Frequency = High ENERGY
Toupee? The Wave-like Electron The electron propagates through space on an energy wave. To understand the atom, one must understand the behavior of electromagnetic waves. Louis deBroglie
Spectroscopic analysis of the visible spectrum… …produces all of the colors in a continuous spectrum
Spectroscopic analysis of the hydrogen spectrum… …produces a “bright line” spectrum
Electron transitionsinvolve jumps of definite amounts ofenergy. This produces bands of light with definite wavelengths.
Schrodinger Wave Equation Equation for probability of a single electron being found along a single axis (x-axis) Erwin Schrodinger
Heisenberg Uncertainty Principle “One cannot simultaneously determine both the position and momentum of an electron.” You can find out where the electron is, but not where it is going. OR… You can find out where the electron is going, but not where it is! Werner Heisenberg
Quantum Numbers Each electron in an atom has a unique set of 4 quantum numbers which describe it. • Principal quantum number • Angular momentum quantum number • Magnetic quantum number • Spin quantum number (n) (l) (m) (s)
Principal Quantum Number Generally symbolized by n, it denotes the shell (energy level) in which the electron is located. Number of electrons that can fit in a shell: 2n2
Angular Momentum Quantum Number The angular momentum quantum number, generally symbolized by l, denotes the orbital (subshell) in which the electron is located. l =3 f
Magnetic Quantum Number The magnetic quantum number, generally symbolized by m, denotes the orientation of the electron’s orbital with respect to the three axes in space.
Assigning the Numbers • The three quantum numbers (n, l, and m) are integers. • The principal quantum number (n) cannot be zero. • n must be 1, 2, 3, etc. • The angular momentum quantum number (l ) can be any integer between 0 and n - 1. • For n = 3, lcan be either 0, 1, or 2. • The magnetic quantum number (ml) can be any integer between -l and +l. • For l = 2, m can be either -2, -1, 0, +1, +2.
Pauli Exclusion Principle No two electrons in an atom can have the same four quantum numbers. Wolfgang Pauli
Spin Quantum Number Spin quantum number denotes the behavior (direction of spin) of an electron within a magnetic field. Possibilities for electron spin:
An orbital is a region within an atom where thereis a probability of finding an electron. This is a probability diagram for the s orbital in the first energy level… Orbital shapes are defined as the surface that contains 90% of the total electron probability.
Sizes of s orbitals Orbitals of the same shape (s, for instance) grow larger as n increases… Nodes are regions of low probability within an orbital.
Orbitals in outer energy levels DO penetrate into lower energy levels. Penetration #1 This is a probability Distribution for a 3s orbital. What parts of the diagram correspond to “nodes” – regions of zero probability?
The s orbital has a spherical shape centered around the origin of the three axes in space. s orbital shape
P orbital shape There are three peanut-shaped p orbitals in each energy level above n = 1, each assigned to its own axis (x, y and z) in space.
d orbital shapes Things get a bit more complicated with the five d orbitals that are found in the d sublevels beginning with n = 3. To remember the shapes, think of: “double peanut” …and a “peanut with a donut”!
Shape of f orbitals Things get even more complicated with the seven f orbitals that are found in the f sublevels beginning with n = 4. To remember the shapes, think of: Flower