MODERN CHEMISTRY CH 3 Atoms: The Building Blocks of Matter

3-1 The Atom: From Philosophical Idea to Scientific Theory Democritus (400B.C.) – Greek philosopher who first used the term atom to describe a basic indivisible particle of matter. Chemical Reaction – transformation of substances into one or more new substances Law of Conservation of Mass – mass is neither created or destroyed during ordinary chemical reactions or physical changes Law of Definite Proportions – chemical compounds contain the same elements in exactly the same proportions my mass (water is always 2-H and 1-O) EX: Every sample of pure salt (NaCl) contains 39.34% Na and 60.66% Cl by mass.

3-1 The Atom: From Philosophical Idea to Scientific Theory Law of Multiple Proportions – If two or more different compounds are composed of the same two elements, then the ratio of the masses of the second element combined with a certain mass of the first element is always a ratio of small whole numbers. - carbon monoxide (CO) contains 1.33 g O and 1.0 g C - carbon dioxide (CO 2 ) contains 2.66 g O and 1.0g C - the ratio of masses of oxygen in the two compounds is 2.66 =

3-1 The Atom: From Philosophical Idea to Scientific Theory John Dalton’s Atomic Theory p. 66 1) All matter is composed of extremely small atoms. 2) Atoms of a given element are identical in size mass, and other properties. 3) Atoms cannot be subdivided, created, or destroyed. 4) Atoms of different elements combine in simple whole number ratios to form compounds. 5 ) In chemical reactions, atoms are combined, separated, or rearranged, but they are NOT created or destroyed. Modern Atomic Theory – -Atoms can be subdivided into smaller particles. -All atoms of the same element are NOT identical.

3-2 The Structure of the Atom Atom – smallest particle of an element that retains the chemical properties of that element Nucleus – small region near the center of the atom (positively charged) Subatomic Particles – positively charged protons, neutral neutrons, negatively charged electrons Cathode Ray Tubes – glass tubes containing gases at low pressure exposed to electrical currents J.J. Thomson – (1897, English Physicist) measured the charge to mass ratio of the cathode rays. The rays were composed of negatively charged particles later named electrons. He developed the Plum Pudding Model of the atom. Robert Millikan – (1909, American Physicist) he used the Oil Drop Experiment to determine the exact charge and mass of the electron. Electron has a mass of 1/1837 th of a single hydrogen atom (basically 1/1837 th the mass of a proton) [p. 71]

3-2 The Structure of the Atom From these two experiments we now know - atoms are electrically neutral – they have as many positive charges as they do negative charges - electron has so much less mass than atoms the atoms must contain other particles to account for their mass. Ernest Rutherford – (1911, New Zealand) In the Gold Foil Experiment he used a radioactive source to fire positively charged alpha particles toward a sheet of gold foil. Protons must be contained in a dense nucleus. The atom is mostly empty space. (p. 72) Niels Bohr – (1913, Denmark) his work showed that electrons follow specific paths or orbits around the nucleus of the atom in an electron cloud. (p )

2-3 The Structure of the Atom Subatomic Particles ParticleSymbolChargeMass #Relative Mass (a.m.u.)Actual Mass (kg) Electrone -, 0 -1 e x Protonp +, 1 1 H x Neutronn 0, 1 0 n x10 -27

3-2 The Structure of the Atom Nuclear Forces – short range proton-neutron, proton-proton, and neutron- neutron forces holding the nuclear particles together (overcomes the repulsive forces of like charges) Electron Cloud – region containing negative charge (electrons) Atomic Radii – range from pm (1 pm = 1x )

3-3 Counting Atoms Atomic Number (Z) – identifies the element and gives the number of protons in the nucleus of each atom of that element (whole number on the periodic table) Neutral Atoms - # protons = #electrons Isotopes – atoms of the same element have the same #p +, but different #n 0 and, therefore different masses (p ) EX:Hydrogen isotopes: Protium (Hydrogen -1) 1 proton, 1 electron Deuterium (Hydrogen -2)1 proton, 1 electron, 1 neutron Tritium (Hydrogen – 3)1 proton, 1 electron, 2 neutrons Mass Number – total number of protons and neutrons (#p + + #n 0 = mass#) Nuclear Symbol – shows composition of isotope’s nucleus (superscript is the mass number and the subscript is the atomic number) EX: U (uranium – 235) and 2 1 H (hydrogen – 2)

3-3 Counting Atoms Nuclide – general term for any isotope of any element EX #1:How many protons, electrons, and neutrons are there is an atom of chlorine – 37? atomic number = 17 #p + = 17 #e - = 17 mass number = 37 #n 0 = mass number – atomic number = = 20 Pause for LOTS of practice with this! (Atomic Structure WS I and II)

3-3 Counting Atoms Relative Atomic Mass – the mass of an atom compared to the carbon-12 isotope Atomic Mass Unit – exactly 1/12 the mass of a carbon-12 atom Average Atomic Mass – weighted average of the atomic masses of the naturally occurring isotopes of an element EX:What is the average atomic mass of naturally occurring copper which consists of 69.17% copper-63 (mass of amu) and % copper-65 (mass of amu)? K:69.17% Cu % Cu-65 unk:average atomic mass (0.6917)( amu) + (0.3083)( ) = amu

3-3 Counting Atoms Composition Stoichiometry - this is basically doing dimensional analysis with atoms and mass Mole (mol) – SI unit for the amount of substance -1 mole contains as many particles as there are in exactly 12 g of carbon-12 -Avogadro’s Number = 6.022x10 23 the number of particles in exactly one mole of a pure substance -Molar Mass – the mass of one mole of a pure substance (g/mol); numerically equal to the atomic mass of the element in atomic mass units (periodic table) EX:1 mol Xe = g Xe 1 mol Na = g Na

3-3 Counting Atoms 1 mole = 6.022x10 23 atoms (or particles) 1 mole = weight in grams from the P.T. We will now work SEVERAL compositional stoichiometry practice problems from your book p

3-3 Counting Atoms Compositional Stoichiometry Examples