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Chapter 4 “Atomic Structure”. Section 4.1 Defining the Atom OBJECTIVES: OBJECTIVES: Describe Democritus’s ideas about atoms. Describe Democritus’s ideas.

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Presentation on theme: "Chapter 4 “Atomic Structure”. Section 4.1 Defining the Atom OBJECTIVES: OBJECTIVES: Describe Democritus’s ideas about atoms. Describe Democritus’s ideas."— Presentation transcript:

1 Chapter 4 “Atomic Structure”

2 Section 4.1 Defining the Atom OBJECTIVES: OBJECTIVES: Describe Democritus’s ideas about atoms. Describe Democritus’s ideas about atoms. Explain Dalton’s atomic theory. Explain Dalton’s atomic theory. Identify what instrument is used to observe individual atoms. Identify what instrument is used to observe individual atoms.

3 Section 4.1 Defining the Atom Democritus Democritus First to suggest the existence of atoms (from the Greek word “atomos”) First to suggest the existence of atoms (from the Greek word “atomos”) He believed that atoms were indivisible and indestructible. He believed that atoms were indivisible and indestructible.

4 Dalton’s Atomic Theory 1)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 3)Atoms of different elements combine in whole-number ratios to form compounds. Dalton’s Atomic Theory 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 hair  Can be observed with scanning tunneling (electron) microscopes

7 Section 4.2 Structure of the Nuclear Atom OBJECTIVES: OBJECTIVES: Identify three types of subatomic particles. Identify three types of subatomic particles. Describe the structure of atoms, according to the Rutherford atomic model. 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: Atoms are divisible into three subatomic particles: Electrons Electrons Protons Protons Neutrons Neutrons

9 Discovery of the Electron J.J. Thomson used a cathode ray tube to discover the negatively charged electron

10 Mass of the Electron Robert Millikan determined the mass of the electron: 1/1840 the mass of a hydrogen atom The oil drop apparatus Mass of the electron is 9.11 x g

11 Conclusions from the Study of the Electron: a)Atoms have no charge, so there must be positive particles to balance the negative charge of the electrons b)Electrons 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 proton  Mass of 1 (or 1840 times that of an electron)  James Chadwick confirmed the neutral neutron  Mass nearly equal to a proton

13 Subatomic Particles ParticleCharge Mass (g) LocationElectron (e - ) (e - ) 9.11 x x Electron cloud Proton (p + ) x x Nucleus Neutron (n o ) (n o ) x x Nucleus

14 Thomson’s Atomic Model Thomson - plum pudding model. Electrons were like plums embedded in a positively charged pudding. J. J. Thomson

15 Ernest Rutherford’s Gold Foil Experiment  Alpha particles (helium nuclei) fired at a thin gold foil.  Particles that hit on the detecting screen are recorded

16 Rutherford’s Findings a) The nucleus is small, dense, and, positively charged  Most of the particles passed right through  A few particles were deflected. Conclusions:

17 The Rutherford Atomic Model Based on his experimental evidence: Based on his experimental evidence: Atom is mostly empty space. Atom is mostly empty space. All the positive charge, and almost all the mass is in the center at the nucleus. All the positive charge, and almost all the mass is in the center at the nucleus.

18 Nucleus is made of protons and neutronsNucleus is made of protons and neutrons Electrons surround the nucleus.Electrons surround the nucleus. Called the “nuclear model”Called the “nuclear model” The Rutherford Atomic Model

19 Section 4.3 Distinguishing Among Atoms OBJECTIVES: OBJECTIVES: Explain what makes elements and isotopes different from each other. Explain what makes elements and isotopes different from each other. Calculate the number of neutrons in an atom. Calculate the number of neutrons in an atom. Calculate the atomic mass of an element. Calculate the atomic mass of an element. Explain why chemists use the periodic table. Explain why chemists use the periodic table.

20 Atomic Number Atoms are composed of identical protons, neutrons, and electrons Atoms are composed of identical protons, neutrons, and electrons How then are atoms of one element different from another element? How then are atoms of one element different from another element?

21 Atomic Number Elements are different because they contain different numbers of PROTONS Elements are different because they contain different numbers of PROTONS Atomic number - number of protons in the nucleus (smaller #) Atomic number - number of protons in the nucleus (smaller #) # protons = # electrons # protons = # electrons Atomic Number: # p+ : # e- : Atomic Number: # p+ : # e- : 35 53

22 Atomic Number Atomic number (Z) of an element is the number of protons in the nucleus of each atom of that element.Element # of protons Atomic # (Z) Carbon (C) Phosphorus (P) Gold (Au)

23 Mass Number Mass number is the number of protons and neutrons in the nucleus of an isotope: Mass # = p + + n 0 Atomic Number: Mass Number: # p+ : # e- : #n 0 : Atomic Number: Mass Number: # p+ : # e- :#n 0 :

24 Atom p+p+p+p+ n0n0n0n0 e-e-e-e- Mass # Oxygen 8 Arsenic Phosphorus Mass Number Practice

25 Complete Symbols Contain the symbol of the element, the mass number and the atomic number. Contain the symbol of the element, the mass number and the atomic number. X Mass number Atomic number Subscript → Superscript →

26 Symbols n Find each of these: a) number of protons b) number of neutrons c) number of electrons d) Atomic number e) Mass Number Na

27 Symbols n If an element has an atomic number of 34 and a mass number of 78, what is the: a) number of protons = b) number of neutrons = c) number of electrons = d) complete symbol X

28 Symbols n If an element has 91 protons and 140 neutrons what is the a) Atomic number = b) Mass number = c) number of electrons = d) complete symbol

29 Symbols n If an element has 78 electrons and 117 neutrons what is the a) Atomic number b) Mass number c) number of protons d) complete symbol

30 Isotopes Dalton was wrong about all elements of Dalton was wrong about all elements of the same type being identical Atoms of the same element can have different numbers of neutrons. Atoms of the same element can have different numbers of neutrons. Thus, different mass numbers. Thus, different mass numbers. These are called isotopes. These are called isotopes. Atomic #: Mass #: # p+: #n 0 : Atomic #: Mass #: # p+: #n 0 : Atomic #: Mass #: # p+: #n 0 :

31 Isotopes Isotopes are atoms of the same element with different masses, due to varying numbers of neutrons.

32 Naming Isotopes We can also put the mass number after the name of the element: We can also put the mass number after the name of the element: carbon-12 Mass: carbon-12 Mass: carbon-14Mass: carbon-14Mass: uranium-235Mass: uranium-235Mass:

33 Isotopes are atoms of the same element having different masses, due to varying numbers of neutrons.Isotope Proton s ElectronsNeutronsNucleusHydrogen–1 (protium) (protium)110 Hydrogen-2(deuterium)111 Hydrogen-3(tritium)112 What’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 Mass Measure atomic mass with the Atomic Mass Unit (amu) Measure atomic mass with the Atomic Mass Unit (amu) Defined as one-twelfth the mass of a carbon-12 atom. 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. Each isotope has its own atomic mass, thus we determine the average from percent abundance.

36 To calculate the average: Multiply the atomic mass of each isotope by it’s abundance (expressed as a decimal), then add the results. Multiply the atomic mass of each isotope by it’s abundance (expressed as a decimal), then add the results. Expressed as amu. Expressed as amu. C-12 = 12 amu. C-12 = 12 amu.

37 Atomic Masses IsotopeSymbol Composition of the nucleus % in nature Carbon C 6 protons 6 neutrons 98.89% Carbon C 6 protons 7 neutrons 1.11% Carbon C 6 protons 8 neutrons <0.01% Atomic mass is the average of all the naturally occurring isotopes of that element. Carbon =

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39 Atomic Mass Example B-10 = 19.8% B-11 = 80.2% 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 column  Also called a family  Elements in a group have similar chemical and physical properties  Identified with number and “A” or “B”

42 Draw an arrow and label a period and a group. Group Period


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