2 Section 4.1 Defining the Atom OBJECTIVES:Describe Democritus’s ideas about atoms.Explain Dalton’s atomic theory.Identify what instrument is used to observe individual atoms.
3 Section 4.1 Defining the Atom DemocritusFirst to suggest the existence of atoms (from the Greek word “atomos”)He believed that atoms were indivisible and indestructible.-Greek philosopher-not based on the scientific method – but just philosophy
4 Dalton’s Atomic Theory All elements are composed of tiny indivisible particles called atoms.John Dalton(1766 – 1844)2) Atoms of the same element are identical.--Atoms of any one element are different from those of any other element.
5 Dalton’s Atomic Theory Atoms of different elements combine in whole-number ratios to form compounds.4) In chemical reactions, atoms are combined, separated, or rearranged – but never changed into atoms of another element.
6 Sizing up the Atom 100,000,000 atoms = 1 cm 1,000,000 atoms = width of hairCan be observed with scanning tunneling (electron) microscopes
7 Section 4.2 Structure of the Nuclear Atom OBJECTIVES:Identify three types of subatomic particles.Describe the structure of atoms, according to the Rutherford atomic model.
8 Section 4.2 Structure of the Nuclear Atom Atoms are divisible into three subatomic particles:ElectronsProtonsNeutrons
9 Discovery of the Electron J.J. Thomson used a cathode ray tube to discover the negatively charged electronCathode ray tubes pass electricity through a gas that is contained at a very low pressure.
10 Mass of the ElectronMass of the electron is9.11 x gThe oil drop apparatusRobert Millikan determined the mass of the electron: 1/1840 the mass of a hydrogen atom
11 Conclusions from the Study of the Electron: Atoms have no charge, so there must be positive particles to balance the negative charge of the electronsElectrons have so little mass that other particles must account for most of the mass
12 Conclusions from the Study of the Electron: Eugen Goldstein observed positive protonMass of 1 (or 1840 times that of an electron)James Chadwick confirmed the neutral neutronMass nearly equal to a protonProtons have a positive charge. Neutrons have no charge. Mass is same.
13 Subatomic Particles Particle Charge Mass (g) Location Electron (e-) -1 9.11 x 10-28Electron cloudProton (p+)+11.67 x 10-24NucleusNeutron(no)Don’t have to know numbers, just charge and how masses compare, and location.
14 Thomson’s Atomic Model J. J. ThomsonThomson - plum pudding model.Electrons were like plums embedded in a positively charged pudding.
15 Ernest Rutherford’s Gold Foil Experiment - 1911 Alpha particles (helium nuclei) fired at a thin gold foil.Particles that hit on the detecting screen are recorded
16 Rutherford’s Findings Most of the particles passed right throughA few particles were deflected.Conclusions:The nucleus is small, dense, and, positively charged
17 The Rutherford Atomic Model Based on his experimental evidence:Atom is mostly empty space.All the positive charge, and almost all the mass is in the center at the nucleus.
18 The Rutherford Atomic Model Nucleus is made of protons and neutronsElectrons surround the nucleus.Called the “nuclear model”
19 Section 4.3 Distinguishing Among Atoms OBJECTIVES:Explain what makes elements and isotopes different from each other.Calculate the number of neutrons in an atom.Calculate the atomic mass of an element.Explain why chemists use the periodic table.
20 Atomic NumberAtoms are composed of identical protons, neutrons, and electronsHow then are atoms of one element different from another element?
21 Atomic NumberElements are different because they contain different numbers of PROTONSAtomic number - number of protons in the nucleus (smaller #)# protons = # electronsAtomic Number:# p+ : # e- :353535535353
22 Atomic NumberAtomic number (Z) of an element is the number of protons in the nucleus of each atom of that element.Element# of protonsAtomic # (Z)Carbon (C)Phosphorus (P)Gold (Au)6615157979
23 Mass Number Mass # = p+ + n0 Mass number is the number of protons and neutrons in the nucleus of an isotope:Mass # = p+ + n0Atomic Number: Mass Number:# p+ : # e- : #n0 :# p+ : # e- : #n0 :3579.935354553127535374
24 Mass Number Practice Atom p+ n0 e- Mass # 16 8 8 8 33 41 33 74 15 16 Oxygen16888Arsenic33413374Phosphorus15161531
25 Complete SymbolsContain the symbol of the element, the mass number and the atomic number.AtomicnumberXSuperscript →Subscript →Massnumber*
26 Na Symbols 11 11 12 23 11 11 23 Find each of these: number of protons number of neutronsnumber of electronsAtomic numberMass Number111112Na23111123*
27 SymbolsIf an element has an atomic number of 34 and a mass number of 78, what is the:number of protons =number of neutrons =number of electrons =complete symbol34433434X78*
28 If an element has 91 protons and 140 neutrons what is the SymbolsIf an element has 91 protons and 140 neutrons what is theAtomic number =Mass number =number of electrons =complete symbol9113191*
29 If an element has 78 electrons and 117 neutrons what is the SymbolsIf an element has 78 electrons and 117 neutrons what is theAtomic numberMass numbernumber of protonscomplete symbol*
30 Isotopes Dalton was wrong about all elements of the same type being identicalAtoms of the same element can have different numbers of neutrons.Thus, different mass numbers.These are called isotopes.Atomic #:Mass #:# p+:#n0:Atomic #:Mass #:# p+:#n0:Atomic #:Mass #:# p+:#n0:*
31 IsotopesIsotopes are atoms of the same element with different masses, due to varying numbers of neutrons.
32 We can also put the mass number after the name of the element: Naming IsotopesWe can also put the mass number after the name of the element:carbon Mass:carbon-14 Mass:uranium-235 Mass:1214235*
33 What’s the only thing that changes? # of neutrons Isotopes are atoms of the same element having different masses, due to varying numbers of neutrons.IsotopeProtonsElectronsNeutronsNucleusHydrogen–1(protium)1Hydrogen-2(deuterium)Hydrogen-3(tritium)2What’s the only thing that changes? # of neutrons
34 Atomic Mass How heavy is an atom of oxygen? Depends - there are different masses of oxygen atoms.We want the average atomic mass.Based on abundance (percentage) of each variety of that element in nature.*
35 Measuring Atomic MassMeasure atomic mass with the Atomic Mass Unit (amu)Defined as one-twelfth the mass of a carbon-12 atom.Each isotope has its own atomic mass, thus we determine the average from percent abundance.We don’t use grams for this mass because the numbers would be too small.Carbon-12 chosen because of its isotope purity. – 6 p+, 6 n0 almost all mass and high % C12.*
36 To calculate the average: Multiply the atomic mass of each isotope by it’s abundance (expressed as a decimal), then add the results.Expressed as amu.C-12 = 12 amu.*
37 Composition of the nucleus Atomic MassesAtomic mass is the average of all the naturally occurring isotopes of that element.IsotopeSymbolComposition of the nucleus% in natureCarbon-1212C6 protons6 neutrons98.89%Carbon-1313C7 neutrons1.11%Carbon-1414C8 neutrons<0.01%Carbon =*
39 Atomic Mass Example (11.0)(.802) 10.8 amu (10.0)(.198) B-10 = 19.8% At. Mass ==(10.0)(.198)(11.0)(.802)10.8 amu
40 The Periodic Table: A Preview Periodic table - arrangement of elements in which the elements are separated into groups based on a set of repeating properties.Allows easy comparison of the properties of different elements
41 The Periodic Table: A Preview Period - horizontal row (there are 7 of them)Group - vertical columnAlso called a familyElements in a group have similar chemical and physical propertiesIdentified with number and “A” or “B”