2 Section 4.1 Defining the Atom Greek philosopher Democritus(460 B.C. – 370 B.C.) suggestedexistence of atoms (Greek word“atomos”)Believed atoms were indivisible and indestructibleHis ideas agreed with later scientific theory, but didn’t explain chemical behavior - was not based on scientific methods – only philosophy
3 Dalton’s Atomic Theory (experiment based!) All elements composed of tiny indivisible particles called atomsAtoms of same element identical. Atoms of any one element are different from all other elements.John Dalton(1766 – 1844)Atoms of different elements combine in simple whole-number ratios to form chemical compoundsIn chemical reactions, atoms are combined, separated, or rearranged – but never changed into atoms of another element.
4 Sizing up the AtomElements subdivided into smaller particles – called atoms, and they still have properties of that element1.0 x 108 copper atoms in a single file, they would be approximately 1 cm longindividual atoms are observable with instruments such as scanning tunneling (electron) microscopes
5 Section 4.2 Structure of the Nuclear Atom One change to Dalton’s atomic theory - atoms are divisible into subatomic particles:Electrons, protons, and neutrons
6 Discovery of ElectronJ.J. Thomson used cathode ray tube to deduce presence of negatively charged particle…….the electron
7 Modern Cathode Ray Tubes TelevisionComputer MonitorCRT’s pass electricity through gas contained - very low pressure.
8 Mass of the ElectronMass of the electron is9.11 x gThe oil drop apparatus1916 – Robert Millikan determines mass of electron: 1/1840 the mass of hydrogen atom; has one unit of negative charge
9 Conclusions from the Study of the Electron: Cathode rays have identical properties regardless of element used to produce them. All elements must contain identically charged electrons.Atoms are neutral, so there must be positive particles in atom to balance negative charge of electronsElectrons have so little mass that atoms must contain other particles that account for most of mass
10 Conclusions from the Study of the Electron: Eugen Goldstein in 1886 observed what is now called the “proton” - particles with a positive charge, and a relative mass of 1 (or 1840 times that of an electron)1932 – James Chadwick confirmed the existence of “neutron” – particle with no charge, but mass nearly equal to proton
11 Subatomic Particles Particle Charge Mass (g) Location Electron (e-) -1 9.11 x 10-28Electron cloudProton (p+)+11.67 x 10-24NucleusNeutron(no)
12 Thomson’s Atomic Model J. J. ThomsonBelieved electrons were like plums embedded in + charged “pudding,” called “plum pudding” model.
13 Ernest Rutherford’s Gold Foil Experiment - 1911 Alpha particles - helium nuclei w/ + charge - The alpha particles were fired at thin sheet of gold foilParticles that hit on the detecting screen (film) were recorded
14 Rutherford’s problem: In the following pictures, there is a target hidden by a cloud. To figure out the shape of the target, we shot some beams into the cloud and recorded where the beams came out. Can you figure out the shape of the target?Target #2Target #1
16 Rutherford’s Findings Most particles passed throughFew deflectedVERY FEW greatly deflected“Like howitzer shells bouncing off of tissue paper!”Conclusions:Small nucleusDense nucleus+ charge nucleusThe Atom SongAtoms song - Mark Rosengarten
17 The Rutherford Atomic Model His experimental evidence:atom mostly empty spaceAll positive charge, almost all mass in small center. “Nucleus”protons and neutrons make nucleus!electrons distributed around nucleus…occupy most volumeHis model called “nuclear model”Rutherford’s Atom 3:08
18 # protons in atom = # electrons Section 4.3 Atomic NumberAll atoms composed of identical protons, neutrons, and electronsHow then are atoms of one element different from another element?Elements different b/c they contain different # of PROTONS“atomic number” of element is number of protons in nucleus# protons in atom = # electrons
19 Atomic NumberAtomic number (Z) of element is # of protons in nucleus of each atom of that element.Element# of protonsAtomic # (Z)Carbon6Phosphorus15Gold79
20 Mass NumberMass number is # of protons and neutrons in nucleus of an isotope:Mass # = p+ + n0Nuclidep+n0e-Mass #Oxygen -10-3342- 3115188818Arsenic753375Phosphorus161531
21 Complete SymbolsContain symbol of element, mass number & atomic number.MassnumberXSuperscript →AtomicnumberSubscript →
22 Br Symbols 80 35 Find each of these: number of protons number of neutronsnumber of electronsAtomic numberMass Number80Br35
23 SymbolsIf an element has an atomic number of 34 and a mass number of 78, what is the:number of protonsnumber of neutronsnumber of electronscomplete symbol
24 SymbolsIf an element has 91 protons and 140 neutrons what is theAtomic numberMass numbernumber of electronscomplete symbol
25 SymbolsIf an element has 78 electrons and 117 neutrons what is theAtomic numberMass numbernumber of protonscomplete symbol
26 IsotopesDalton was wrong about elements of same type being identical…Atoms of same element can have different numbers of neutrons.different mass numbersisotopes
27 Isotopes Frederick Soddy (1877-1956) proposed idea of isotopes in 1912 Isotopes - atoms of same element with different masses, b/c varying #s of neutronsWon 1921 Nobel Prize in Chemistryhas a small crater named for him on the far side of the Moon.
28 We can also put mass number after name of the element: Naming IsotopesWe can also put mass number after name of the element:carbon-12carbon-14uranium-235
30 IsotopesElements occur in nature as mixtures of isotopes.
31 Atomic Mass How heavy is an oxygen atom? Depends, b/c different kinds of oxygen atoms exist.We’re more concerned with average atomic mass.Based on abundance (%) of each variety of that element in nature.Don’t use grams - numbers tooooo small.
32 Measuring Atomic Mass Atomic Mass Unit (amu) one-twelfth mass of a carbon-12 atom.Carbon-12 chosen b/c of its isotope purity.Each isotope has own atomic masswe determine average from % abundance.
33 To calculate the average: Multiply atomic mass of each isotope by abundance (decimal), then add results.If not told otherwise, mass of isotope expressed in atomic mass units (amu)
34 Composition of the nucleus Atomic MassesAtomic mass is average of all naturally occurring isotopes of that element.Atomic mass (amu)1213.0014.00IsotopeSymbolComposition of the nucleus% in natureCarbon-1212C6 protons6 neutrons98.89%Carbon-1313C7 neutrons1.11%Carbon-1414C8 neutrons<0.01%What is the average atomic mass of Carbon?12.01
35 - Page 117QuestionKnowns and UnknownSolutionAnswer