Presentation on theme: "Unit II : Atoms, Molecules, Ions and Nuclear Chemistry."— Presentation transcript:
Unit II : Atoms, Molecules, Ions and Nuclear Chemistry
Early History of Chemistry Before 16th Century Alchemy: Attempts (scientific or otherwise) to change cheap metals into gold 17th Century Robert Boyle: First chemist to perform quantitative experiments 18th Century George Stahl: Phlogiston flows out of a burning material. Joseph Priestley: Discovers oxygen gas, dephlogisticated air.
Law of Conservation of Mass -Discovered by Antoine Lavoisier -Mass is neither created nor destroyed -Combustion involves oxygen, not phlogiston
Law of Definite Proportions - Sometimes called the law of constant composition - John Proust (1799) - A given compound always contains exactly the same proportion of elements by mass. - Carbon tetrachloride is always 1 atom carbon per 4 atoms chlorine.
Law of Multiple Proportions - When two elements form a series of compounds, the ratios of the masses of the second element that combine with 1 gram of the first element can always be reduced to small whole numbers. - The ratio of the masses of oxygen in H 2 O and H 2 O 2 will be a small whole number (2).
Daltons Atomic Theory 1Each element is made up of tiny particles called atoms. 2The atoms of a given element are identical; the atoms of different elements are different in some fundamental way or ways. 3Chemical compounds are formed when atoms combine with each other. A given compound always has the same relative numbers and types of atoms. 4Chemical reactions involve reorganization of the atoms - changes in the way they are bound together. The atoms themselves are not changed in a chemical reaction.
Avagadros Hypothesis (1811) At the same temperature and pressure, equal volumes of different gases contain the same number of particles. 5 liters of oxygen 5 liters of nitrogen SAME NUMBER OF PARTICLES
Chemical Symbols The chemical symbols used today were developed by Jons Jakob Berzelius. They consist of one or two letters. The first letter is always capitalized. The second, if there is one, is never capitalized. The second letter is often a letter prominent in the pronunciation. CoCuHe H
Early Experiments to Characterize the Atom Democritis Dalton Thomson Rutherford Bohr Schrodinger
Cathode Ray Streams of negatively charged particles were found to emanate from cathode tubes. J. J. Thompson is credited with their discovery (1897).
Deflection of Cathode Ray by an Applied Electric Field
Rutherford Gold Foil Experiment Ernest Rutherford shot particles at a thin sheet of gold foil and observed the pattern of scatter of the particles.
Rutherford Gold Foil Experiment (a)Expected results if Thomsons Model was correct (b)Actual results
Summary of Atom Democritis - first idea of the atom Dalton - ~1807 Atomic Theory (atom = ball) Thomson - Experiment with Cathode Ray Discovered the electron Plum pudding model Rutherford - Gold Foil Experiment Atom is mostly empty space with a dense, positively charged nucleus Bohr - Solar System Schrodinger - Quantum Mechanical Model
Modern View of Atomic Structure The atom contains 3 types of subatomic particles: Electrons Negatively charged, found outside the nucleus… very small mass… would take ~2000 electrons to equal the mass of 1 of the other subatomic particles Protons Positively charged, found in the nucleus Neutrons Found in the nucleus… like a proton but with no charge
Subatomic Particle Summary Mass (g)ChargeAtomic Mass Scale Relative Mass Relative Charge Proton (p + ) x x c amu 1+1 Neutron (n 0 ) x amu 10 Electron (e - ) x x c x amu 1/1836
Useful Units Atomic Diameter: m (1 Å) = 100 pm = 1x10 -8 cm Nuclear Diameter: 10-13cm 1 amu (atomic mass unit) = x kg Based off of C-12… relative mass of O w/ 8 protons and 8 neutrons is picometer (pm) = 1 x m NOTE: Heaviest atom has a mass of only 4.8x g and a diameter of only 5x m Biggest atom is 240 amu and is 50 Å across. Typical C-C bond length 154 pm (1.54 Å) Molecular models are 1 Å /inch or about 0.4 Å /cm
Atomic Number K Element Symbol Z = atomic number = number of protons A = mass number = number of protons + number of neutrons A = Z + N Mass number = # p + + # n 0
Isotopes Isotopes are atoms of the same element with different masses. Isotopes have different numbers of neutrons. Carbon has two natural isotopes: 12 C and 13 C How does this change things?
Isotopes On the periodic table the given atomic mass is the AVERAGE mass We use percentage to help with this math. For Carbon, there is 98.89% 12 C and 1.11% of 13 C Now what? 12 (mass of 12 C) x = (mass of 13 C) x = =
Isotopes On the periodic table the given atomic mass is the AVERAGE mass We use percentage to help with this math. For Carbon, there is 98.89% 12 C and 1.11% of 13 C
Mass Spectrometer Used to experimentally determine mass and percent abundances of isotopes
Periodicity When one looks at the chemical properties of elements, one notices a repeating pattern of reactivities.
The Periodic Table Dmiti Mendeleev organized 1 st periodic table by lining elements up in horizontal rows in order of increasing atomic weight Left spaces when an element was not known but should exist and have properties similar to the element about it in his table. Periodicity – periodic repetition of the properties of elements H.G.J. Moseley organized elements according to increasing atomic number Law of chemical periodicity The properties of the elements are periodic functions of atomic numbers.
Features of the Periodic Table Vertical Columns: Groups or Families A and B groups. A main group, B transition elements Horizontal Rows: called periods Periodic Law: when the elements are arranged in order of increasing atomic number, there is a periodic repetition of their physical and chemical properties
Metals ~80% of all elements are metals High electrical conductivity and have a high luster when cleaned Ductile (can be made into wires) and Malleable Can form alloys Are solid at room temperature except for one… which one is it?
Nonmetals Generally nonlustrous… which means not shiny Generally poor conductors of electricity With the exceptions of carbon (graphite), none conduct electricity Allotropes A particular element can often exist in several different forms Carbon as graphite or diamond… oxygen as O 2 or O 3 (ozone)
Overview of Groups Group 1A Alkali Metals Metals, solid at room temp, and VERY REACTIVE Only found in nature combined in compounds Group 2A Alkaline Earth Metals Metals, solid at room temp, and also only found in nature in compounds Except for Be, all react with water to produce alkaline solutions
Overview of Groups Group 7A Halogens All exist in diatomic molecules Combine violently with alkali metals to form salts Group 8A Noble Gases (Inert Gases) Least reactive elements All are gases
Overview of Groups Group B Metals (Groups 3-12): Transition Metals & Inner Transition Metals Lanthanides- shiny metals similar in reactivity to alkaline earth metals Actinides- unique in nuclear structures – nuclei are unstable and therefore radioactive Inner Transitions Metals – rare-earth metals
Chemical Formulas The subscript to the right of the symbol of an element tells the number of atoms of that element in one molecule of the compound. Molecular compounds are composed of molecules and almost always contain only nonmetals.
Chemical Formulas: Diatomic Molecules These 7 elements occur naturally as diatomic molecules
Types of Formulas Molecular Formula – CH 4 Structural Formula Ball and Stick Model Space-Filling Model
Types of Formulas Molecular Formula Exact number of each type of atom C 6 H 12 O 6 Empirical Formula Lowest whole number ratio of each type of atom CH 2 O