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A History of Atomic Theory & Basic Atomic Structure Chapter 3: The Atom Big Idea: Physical, chemical and nuclear changes are explained using the location.

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Presentation on theme: "A History of Atomic Theory & Basic Atomic Structure Chapter 3: The Atom Big Idea: Physical, chemical and nuclear changes are explained using the location."— Presentation transcript:

1 A History of Atomic Theory & Basic Atomic Structure Chapter 3: The Atom Big Idea: Physical, chemical and nuclear changes are explained using the location and properties of subatomic particles.

2 Section 3.1 The Atom: From Philosophical Idea to Scientific Theory

3  All matter is made up of tiny, indivisible particles called atomos  Today, we define atom as the smallest particle of an element that retains the chemical identity of that element  Aristotle asked: Democritus – 450 BC What holds the tiny particles together? Democritus: ???

4  Aristotle rejected Democritus reasoning and proposed that matter was a continuum composed of mass and form Marble (mass)  Statue (form)  Later the simplest forms of matter were proposed to be: Earth, Water, Fire, Wind Aristotle – 384 BC

5  The transformation of a substance(s) into one or more new substances is known as a chemical reaction.  Law of Definite Composition: a chemical compound contains the same elements in exactly the same proportions by mass (regardless sample size or source) Foundations of Atomic Theory Sugar: 42.1 % Carbon 51.4 % Oxygen 6.5 % Hydrogen Whether you have a teaspoon or a truckload!

6  Law of Conservation of Mass: mass is neither created nor destroyed during ordinary chemical reactions or physical changes Foundations of Atomic Theory HgO  Hg + O 433.2 g 401.2g + 32g

7  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 H 2 O H 2 O 2 Water Peroxide 2g H 16g O 1:2 Ratio 32g O

8 2

9 Dalton’s Atomic Theory  All matter is composed of extremely small particles called atoms.  Atoms of a given element are identical in size, mass, and other properties; atoms of different elements differ in size, mass, and other properties.  Atoms cannot be subdivided, created, or destroyed.  satisfies Law of Conservation of Mass Dalton - 1808

10  Atoms of different elements combine in simple whole-number ratios to form chemical compounds.  satisfies Law of Multiple Proportions  In chemical reactions, atoms are combined, separated, or rearranged.

11  Not all aspects of Dalton’s atomic theory have proven to be correct.  Atoms can be split into even smaller particles.  A given element can have atoms with different masses (called isotopes)  Some important concepts remain unchanged  All matter is composed of atoms  Atoms are rearranged in chemical reactions  Atoms of any one element are never identical to atoms of another element Modern Atomic Theory

12 Section 3.2 The Structure of the Atom

13  Atom is the smallest particle of an element that retains the chemical properties of that element.  1897 - Joseph John Thomson’s cathode-ray tube (CRT) The Discovery of the Electron

14 Negatively Charged Electrode Positively Charged Electrode Joseph John Thomson’s cathode-ray tube (CRT)

15 Cathode Ray Tube  Scientists studied the flow of electric current in a glass vacuum tube with electrodes at each end.  When connected to electric current the remaining gas glowed forming a BEAM OF LIGHT.  The beam always originated at the NEGATIVE electrode and toward the POSITIVE electrode.  The electrode is named by what type of particle it attracts  Cathode: Negative (-)  Anode: Positive (+)

16  JJ Thomson used magnets to deflect the beam proving that particles had a negative charge.  These negatively charges particles were called electrons.  Major contribution to the atom:  Electrons are in all atoms!

17  Cathode Ray Tube Experiment Cathode Ray Tube Experiment CRT Video

18 Robert A. Millikan - 1909  Continued Thomson’s work –  performed the Oil Drop Experiment  confirmed the negative charge of an electron and measured the mass of an electron  The electron has mass, though 1836 x less than that of soon to be discovered proton.

19  Thomson proposed that the electrons of an atom were spread evenly throughout a positively charged ball of matter.  Known as Plum-pudding model J.J. Thomson’s Plum Pudding Model

20 Plum Pudding Video

21 The Discovery of the Atomic Nucleus  Earnest Rutherford’s Gold Foil Experiment - 1909

22 Gold Foil Experiment  Set up Gold Foil with a detection sheet around it.  Set up radioactive source- emitted alpha particles.  ALPHA PARTICLES shot at gold foil.  MOST particles went through the gold foil  But SOME particles BOUNCED back

23 Gold Foil Conclusions 1.The atom is made up of mostly EMPTY SPACE 2.The center of the atom contains a POSITIVE CHARGE 3.Rutherford called this positive bundle of matter the NUCLEUS  Rutherford’s major contribution to the atom was the discovery of the nucleus. The volume of this is very small compared with the total volume of an atom.

24 Rutherford’s Model of the Atom atomic radius ~ 100 pm = 1 x 10 -10 m nuclear radius ~ 5 x 10 -3 pm = 5 x 10 -15 m

25  When two protons are extremely close to each other, there is a strong attraction between them.  A similar attraction exists when neutrons are very close to each other or protons and neutrons.  The short-range proton-neutron, proton-proton, and neutron-neutron forces that hold the nuclear particles together are referred to as nuclear forces. Composition of the Nucleus

26 Structure of the Atom  The nucleus is a very small region located at the center of an atom.  The nucleus is made up of at least one positively charged particle called a proton and usually one or more neutral particles called neutrons.  Surrounding the nucleus is a region occupied by negatively charged particles called electrons.  P, N, E are often referred to as subatomic particles.

27 ParticleSymbolChargeMass Number Actual Mass (kg) Electrone-e- 09.109 x 10 -31 Protonp+p+ +111.673 x 10 -27 Neutronnono 011.675 x 10 -27

28 Section 3.3 Ions and Isotopes


30 Atomic number (Z) = number of protons in nucleus Mass number (A) = number of protons + number of neutrons = atomic number (Z) + number of neutrons Isotopes are atoms of the same element (X) with different numbers of neutrons in their nuclei X A Z H 1 1 H (D) 2 1 H (T) 3 1 U 235 92 U 238 92 Mass Number Atomic Number Element Symbol

31 An ion is an atom, or group of atoms, that has a net positive or negative charge. cation – ion with a positive charge If a neutral atom loses one or more electrons it becomes a cation. anion – ion with a negative charge If a neutral atom gains one or more electrons it becomes an anion. Na 11 protons 11 electrons Na + 11 protons 10 electrons Cl 17 protons 17 electrons Cl - 17 protons 18 electrons

32 Atomic Mass Unit  One atom is the standard – Carbon  Mass of other elements are based off of the standard  Carbon: 6 p and 6 n = 12 amu  1/12 mass of Carbon atom  Periodic table lists weighted average atomic masses of elements (like a GPA calculation) Relative Atomic Mass

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