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The diameter of an atom is 0.1 to 0.5 nm. This is 1 to 5 ten billionths of a meter. ... the way to an understanding of the subatomic structure of the atom. ...

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chapter 5 early atomic theory structure
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 ofSubatomic Particles

5.3 Composition of Compounds

5.9 Atomic Numbers of theElements

5.4 The Nature of Electric Charge

5.5 Discovery of Ions

5.10 Isotopes of the Elements

5.11 Atomic Mass

early thoughts
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 470-370 B.C.) thought that all forms of matter were divisible into tiny indivisible particles. He called them “atoms” from the Greek “atomos” indivisible.
early thoughts3
Early Thoughts
  • Aristotle (384-322 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.
slide4
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.

dalton s atomic theory
Modern research has demonstrated that atoms are composed of subatomic particles.

Atoms under special circumstances can be decomposed.

Dalton’s Atomic Theory
  • Elements are composed of minute indivisible particles called atoms.
  • 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.
dalton s atomic theory6
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.
slide7
Dalton’s atoms were individual particles.

Atoms of each element are alike in mass and size.

5.1

slide8
Dalton’s atoms were individual particles.

Atoms of different elements are not alike in mass and size.

5.1

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.

slide11
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.
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.

slide13
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.
slide14
Combining Ratios of Hydrogen and Oxygen
  • The formula for water is H2O.
  • The formula for hydrogen peroxide is H2O2.
  • Hydrogen peroxide has twice as many oxygens per hydrogen atom as does water.
the nature of electric charge
q1 and q2 are charges, r is the distance between charges and k is a constant.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).
discovery of ions
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).
discovery of ions18
Δ

NaCl → Na+ + Cl-

Discovery of Ions
  • Arrhenius accounted for the electrical conduction of molten sodium chloride (NaCl) by proposing that melted NaCl dissociated into the charged ions Na+ and Cl-.
  • Svante Arrhenius reasoned that an ion is an atom (or a group of atoms) carrying a positive or negative electric charge.
discovery of ions19
Discovery of Ions
  • 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.

NaCl → Na+ + Cl-

the diameter of an atom is 0 1 to 0 5 nm
Subatomic Parts of the AtomThe diameter of an atom is 0.1 to 0.5 nm.

This is 1 to 5 ten billionths of a meter.

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.

subatomic particle electron
Subatomic Particle - Electron
  • Crookes tubes experiments led the way to an understanding of the subatomic structure of the atom.
  • Crookes tube emissions are called cathode rays.
  • In 1875 Sir William Crookes invented the Crookes tube.
subatomic particle electron22
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.
subatomic particle proton
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.
subatomic particle neutron
Subatomic Particle - Neutron
  • Its actual mass is slightly greater than the mass of a proton.
  • James Chadwick discovered the neutron in 1932.
slide27
Ions
  • Negative ions were explained by assuming that extra electrons can be added to atoms.
  • Positive ions were explained by assuming that a neutral atom loses electrons.
the nuclear atom
The Nuclear Atom
  • 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.
  • Radioactivity was discovered by Becquerel in 1896.
the rutherford experiment
The Rutherford Experiment
  • 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.
  • Rutherford in 1911 performed experiments that shot a stream of alpha particles at a gold foil.
slide32
The Rutherford Experiment

Rutherford’s alpha particle scattering experiment.

5.5

the rutherford experiment33
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.
the rutherford experiment34
The Rutherford Experiment
  • Most of the alpha particles passed through the gold foil. This led Rutherford to conclude that a gold atom was mostly empty space.
  • If a positive alpha particle approached close enough to the positive mass it was deflected.
the rutherford experiment35
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.
slide36
The Rutherford Experiment

Deflection

Scattering

Deflection and scattering of alpha particles by positive gold nuclei.

5.5

general arrangement of subatomic particles
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.
general arrangement of subatomic particles38
General Arrangement of Subatomic Particles
  • The negative electrons surround the nucleus.
  • The nucleus contains protons and neutrons
  • Rutherford put forward a model of the atom in which a dense, positively charged nucleus is located at the atom’s center.
atomic numbers of the elements
Atomic Numbers of the Elements
  • The atomic number of an atom determines which element the atom is.
  • The atomic number of an element is equal to the number ofprotons in the nucleus of that element.
slide41
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.

slide42
atomic number

1 proton in the nucleus

1H

Every atom with an atomic number of 1 is a hydrogen atom.

slide44
atomic number

6 protons in the nucleus

6C

Every atom with an atomic number of 6 is a carbon atom.

slide45
atomic number

92 protons in the nucleus

92U

Every atom with an atomic number of 92 is a uranium atom.

atomic numbers of the elements46
Atomic Numbers of the Elements
  • 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.
  • Atoms of the same element can have different masses.
atomic numbers of the elements47
Atomic Numbers of the Elements

Isotopes of the Same Element Have

Equal numbers of protons

Different numbers of neutrons

slide49
12

C

6

Isotopic Notation

6 protons + 6 neutrons

6 protons

slide50
Isotopic Notation

6 protons + 8 neutrons

14

C

6

6 protons

slide51
Isotopic Notation

8 protons + 8 neutrons

16

O

8

8 protons

slide52
Isotopic Notation

8 protons + 9 neutrons

17

O

8

8 protons

slide53
Isotopic Notation

8 protons + 10 neutrons

18

O

8

8 protons

slide54
Hydrogen has three isotopes

1 proton

0 neutrons

1 proton

1 neutron

1 proton

2 neutrons

slide55
Examples of Isotopes

ElementProtonsElectronsNeutronsSymbol

Hydrogen 1 1 0

Hydrogen 1 1 1

Hydrogen 1 1 2

Uranium 92 92 143

Uranium 92 92 146

Chlorine 17 17 18

Chlorine 17 17 20

atomic mass
Atomic Mass
  • Using a mass spectrometer, the mass of the hydrogen atom was determined.
  • The mass of a single atom is too small to measure on a balance.
slide57
Positive ions formed from sample.

Electrical field at slits accelerates positive ions.

Deflection of positive ions occurs at magnetic field.

A Modern Mass Spectrometer

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

slide58
A typical reading from a mass spectrometer. The two principal isotopes of copper are shown with the abundance (%) given.

5.9

atomic mass59
Atomic Mass

Using a mass spectrometer, the mass of one hydrogen atom was determined to be 1.673 x 10-24 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.

atomic mass60
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 = 1.6606 x 10-24 g

atomic mass61
Atomic Mass

Average atomic mass 1.00797 amu.

Average atomic mass 39.098 amu.

Average atomic mass 248.029 amu.

average relative atomic mass
Average Relative Atomic Mass
  • 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 12.0000…amu).
  • Most elements occur as mixtures of isotopes.
slide63
To calculate the atomic mass multiply the atomic mass of each isotope by its percent abundance and add the results.

(62.9298 amu)

0.6909 =

43.48 amu

(64.9278 amu)

0.3091 =

20.07 amu

63.55 amu

slide64
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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