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Chapter 5 – Early Atomic Theory & Structure

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1 Chapter 5 – Early Atomic Theory & Structure
5.1 Early Thoughts 5.6 Subatomic Parts of the Atom 5.7 The Nuclear Atom 5.2 Dalton's Model of the Atom 5.8 General Arrangement of Subatomic Particles 5.3 Composition of Compounds 5.9 Atomic Numbers of the Elements 5.4 The Nature of Electric Charge 5.5 Discovery of Ions 5.10 Isotopes of the Elements 5.11 Atomic Mass

2 Early Thoughts The earliest models of the atom were developed by the ancient Greek philosophers. Empedocles stated that matter was made of 4 elements: earth, air, fire, and water. Democritus (about B.C.) thought that all forms of matter were divisible into tiny indivisible particles. He called them “atoms” from the Greek “atomos” indivisible.

3 Early Thoughts Aristotle ( B.C.) rejected the theory of Democritus and advanced the Empedoclean theory. Aristotle’s influence dominated the thinking of scientists and philosophers until the beginning of the 17th century.

4 Dalton’s Model of the Atom
2000 years after Aristotle, John Dalton an English schoolmaster, proposed his model of the atom–which was based on experimentation.

5 Dalton’s Atomic Theory
Elements are composed of minute indivisible particles called atoms. Atoms under special circumstances can be decomposed. Modern research has demonstrated that atoms are composed of subatomic particles. Atoms of the same element are alike in mass and size. Atoms of different elements have different masses and sizes. Chemical compounds are formed by the union of two or atoms of different elements.

6 Dalton’s Atomic Theory
Atoms combine to form compounds in simple numerical ratios, such as one to one , two to two, two to three, and so on. Atoms of two elements may combine in different ratios to form more than one compound.

7 Dalton’s atoms were individual particles.
Atoms of each element are alike in mass and size. 5.1

8 Dalton’s atoms were individual particles.
Atoms of different elements are not alike in mass and size. 5.1

9 Daltons atoms combine in specific ratios to form compounds.

10 The Law of Definite Composition
Composition of Compounds The Law of Definite Composition A compound always contains two or more elements combined in a definite proportion by mass.

11 Water always contains the same two elements: hydrogen and oxygen.
Composition of Water Water always contains the same two elements: hydrogen and oxygen. The percent by mass of hydrogen in water is 11.2%. The percent by mass of oxygen in water is 88.8%. Water always has these percentages. If the percentages were different the compound would not be water.

12 The Law of Multiple Proportions
Composition of Compounds The Law of Multiple Proportions Atoms of two or more elements may combine in different ratios to produce more than one compound.

13 Composition of Hydrogen Peroxide
Hydrogen peroxide always contains the same two elements: hydrogen and oxygen. The percent by mass of hydrogen in hydrogen peroxide is 5.9%. The percent by mass of oxygen in hydrogen peroxide is 94.1%. Hydrogen peroxide always has these percentages. If the percentages were different the compound would not be hydrogen peroxide.

14 Combining Masses of Hydrogen and Oxygen
Mass Hydrogen(g) Mass Oxygen(g) Water 1.0 8.0 Hydrogen Peroxide 16.0 Hydrogen peroxide has twice as much oxygen (by mass) as does water.

15 Combining Ratios of Hydrogen and Oxygen
Hydrogen peroxide has twice as many oxygens per hydrogen atom as does water. The formula for water is H2O. The formula for hydrogen peroxide is H2O2.

16 Composition of Compounds

17 The Nature of Electric Charge
Properties of Electric Charge Charge may be of two types: positive and negative. Unlike charges attract (positive attracts negative), and like charges repel (negative repels negative and positive repels positive). Charge may be transferred from one object to another, by contact or induction. The less the distance between two charges, the greater the force of attraction between unlike charges (or repulsion between identical charges). q1 and q2 are charges, r is the distance between charges and k is a constant.

18 Discovery of Ions Michael Faraday discovered that certain substances when dissolved in water conducted an electric current. He found that atoms of some elements moved to the cathode (negative electrode) and some moved to the anode (positive electrode). He concluded they were electrically charged and called them ions (Greek wanderer).

19 Discovery of Ions NaCl → Na+ + Cl- Δ
Svante Arrhenius reasoned that an ion is an atom (or a group of atoms) carrying a positive or negative electric charge. Arrhenius accounted for the electrical conduction of molten sodium chloride (NaCl) by proposing that melted NaCl dissociated into the charged ions Na+ and Cl-. NaCl → Na+ + Cl- Δ

20 Discovery of Ions NaCl → Na+ + Cl-
In the melt the positive Na+ ions moved to the cathode (negative electrode). Thus positive ions are called cations. In the melt the negative Cl- ions moved to the anode (positive electrode). Thus negative ions are called anions.

21 The diameter of an atom is 0.1 to 0.5 nm.
Subatomic Parts of the Atom This is 1 to 5 ten billionths of a meter. The diameter of an atom is 0.1 to 0.5 nm. If the diameter of this dot is 1 mm, then 10 million hydrogen atoms would form a line across the dot. Even smaller particles than atoms exist. These are called subatomic particles.

22 Subatomic Particle - Electron
In 1875 Sir William Crookes invented the Crookes tube. Crookes tubes experiments led the way to an understanding of the subatomic structure of the atom. Crookes tube emissions are called cathode rays.

23 Subatomic Particle - Electron
In 1897 Sir Joseph Thompson demonstrated that cathode rays: travel in straight lines. are negative in charge. are deflected by electric and magnetic fields. produce sharp shadows are capable of moving a small paddle wheel.

24 Subatomic Particle - Electron
This was the discovery of the fundamental unit of charge – the electron.

25 Subatomic Particle - Proton
Eugen Goldstein, a German physicist, first observed protons in 1886: Thompson determined the proton’s characteristics. Thompson showed that atoms contained both positive and negative charges. This disproved the Dalton model of the atom which held that atoms were indivisible.

26 Subatomic Particle - Neutron
James Chadwick discovered the neutron in 1932. Its actual mass is slightly greater than the mass of a proton.

27 Subatomic Particles

28 Ions Positive ions were explained by assuming that a neutral atom loses electrons. Negative ions were explained by assuming that extra electrons can be added to atoms.

29 When one or more electrons are lost from an atom, a cation is formed.
5.4

30 When one or more electrons are added to a neutral atom, an anion is formed.
5.4

31 The Nuclear Atom Radioactivity was discovered by Becquerel in 1896.
Radioactive elements spontaneously emit alpha particles, beta particles and gamma rays from their nuclei. By 1907 Rutherford found that alpha particles emitted by certain radioactive elements were helium nuclei.

32 The Rutherford Experiment
Rutherford in 1911 performed experiments that shot a stream of alpha particles at a gold foil. Most of the alpha particles passed through the foil with little or no deflection. He found that a few were deflected at large angles and some alpha particles even bounced back.

33 The Rutherford Experiment
Rutherford’s alpha particle scattering experiment. 5.5

34 The Rutherford Experiment
An electron with a mass of 1/1837 amu could not have deflected an alpha particle with a mass of 4 amu. Rutherford knew that like charges repel. Rutherford concluded that each gold atom contained a positively charged mass that occupied a tiny volume. He called this mass the nucleus.

35 The Rutherford Experiment
If a positive alpha particle approached close enough to the positive mass it was deflected. Most of the alpha particles passed through the gold foil. This led Rutherford to conclude that a gold atom was mostly empty space.

36 The Rutherford Experiment
Because alpha particles have relatively high masses, the extent of the deflections led Rutherford to conclude that the nucleus was very heavy and dense.

37 The Rutherford Experiment
Scattering Deflection Deflection and scattering of alpha particles by positive gold nuclei. 5.5

38 General Arrangement of Subatomic Particles
Rutherford’s experiment showed that an atom had a dense, positively charged nucleus. Chadwick’s work in 1932 demonstrated the atom contains neutrons. Rutherford also noted that light, negatively charged electrons were present in an atom and offset the positive nuclear charge.

39 General Arrangement of Subatomic Particles
Rutherford put forward a model of the atom in which a dense, positively charged nucleus is located at the atom’s center. The negative electrons surround the nucleus. The nucleus contains protons and neutrons

40 General Arrangement of Subatomic Particles
5.6

41 Atomic Numbers of the Elements
The atomic number of an element is equal to the number of protons in the nucleus of that element. The atomic number of an atom determines which element the atom is.

42 Atomic Numbers of the Elements
Every atom with an atomic number of 1 is a hydrogen atom. Every hydrogen atom contains 1 proton in its nucleus.

43 atomic number 1H Every atom with an atomic number of 1 is a hydrogen atom. 1 proton in the nucleus

44 Atomic Numbers of the Elements
Every atom with an atomic number of 6 is a carbon atom. Every carbon atom contains 6 protons in its nucleus.

45 atomic number 6C Every atom with an atomic number of 6 is a carbon atom. 6 protons in the nucleus

46 atomic number 92U Every atom with an atomic number of 92 is a uranium atom. 92 protons in the nucleus

47 Atomic Numbers of the Elements
Atoms of the same element can have different masses. They always have the same number of protons, but they can have different numbers of neutrons in their nuclei. The difference in the number of neutrons accounts for the difference in mass. These are isotopes of the same element.

48 Atomic Numbers of the Elements
Isotopes of the Same Element Have Equal numbers of protons Different numbers of neutrons

49 Isotopic Notation

50 Isotopic Notation 6 protons + 6 neutrons C 6 12 6 protons

51 Isotopic Notation 6 protons + 8 neutrons 14 C 6 6 protons

52 Isotopic Notation 8 protons + 8 neutrons 16 O 8 8 protons

53 Isotopic Notation 8 protons + 9 neutrons 17 O 8 8 protons

54 Isotopic Notation 8 protons + 10 neutrons 18 O 8 8 protons

55 Hydrogen has three isotopes
1 proton 0 neutrons 1 proton 1 neutron 1 proton 2 neutrons

56 Examples of Isotopes Element Protons Electrons Neutrons Symbol
Hydrogen Hydrogen Hydrogen Uranium Uranium Chlorine Chlorine

57 Atomic Mass The mass of a single atom is too small to measure on a balance. Using a mass spectrometer, the mass of the hydrogen atom was determined.

58 A Modern Mass Spectrometer
Positive ions formed from sample. Electrical field at slits accelerates positive ions. Deflection of positive ions occurs at magnetic field. From the intensity and positions of the lines on the mass spectrogram, the different isotopes and their relative amounts can be determined. A mass spectrogram is recorded. 5.8

59 A typical reading from a mass spectrometer
A typical reading from a mass spectrometer. The two principal isotopes of copper are shown with the abundance (%) given. 5.9

60 Atomic Mass Using a mass spectrometer, the mass of one hydrogen atom was determined to be x g. To overcome this problem of such a small mass relative atomic masses using “atomic mass units” was devised to express the masses of elements using simple numbers.

61 Atomic Mass The standard to which the masses of all other atoms are compared to was chosen to be the most abundant isotope of carbon – carbon 12. A mass of exactly 12 atomic mass units (amu) was assigned to carbon 12. 1 amu is defined as exactly equal to the mass of a carbon-12 atom 1 amu = x g

62 Atomic Mass Average atomic mass 1.00797 amu.

63 Average Relative Atomic Mass
Most elements occur as mixtures of isotopes. Isotopes of the same element have different masses. The listed atomic mass of an element is the average relative mass of the isotopes of that element compared to the mass of carbon-12 (exactly …amu).

64 Average atomic mass (amu)
To calculate the atomic mass multiply the atomic mass of each isotope by its percent abundance and add the results. Isotope Isotopic mass (amu) Abundance (%) Average atomic mass (amu) 69.09 30.91 ( amu) = 43.48 amu ( amu) = 20.07 amu 63.55 amu

65 Relationship Between Mass Number and Atomic Number
The mass number minus the atomic number equals the number of neutrons in the nucleus. mass number atomic number mass number atomic number number of neutrons - = 62 109 - 47 =

66 Concepts Dalton’s Atomic Mode; Law of Definite Composition
Law of Multiple Proportions Three principle subatomic particles Thomson Model of the Atom Rutherford alpha-scattering experiment Atomic Number, Mass number, number of neutrons, number of Protons, Number of Electrons Three Isotopes of Hydrogen Average Atomic Mass of an Element


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