1 Chapter 5 – Early Atomic Theory & Structure 5.1 Early ThoughtsEarly Thoughts 5.2 Dalton's Model of the AtomDalton's Model of the Atom 5.3 Composition.

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Chapter 5 – Early Atomic Theory & Structure

5.1 Early Thoughts

5.2 Dalton's Model of the Atom

5.3 Composition of Compounds

5.4 The Nature of Electric Charge

5.5 Discovery of Ions

5.6 Subatomic Parts of the Atom

5.7 The Nuclear Atom

5.8 General Arrangement of Subatomic Particles

5.9 Atomic Numbers of the Elements

5.10 Isotopes of the Elements

5.11 Atomic Mass

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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.

• The earliest models of the atom were developed by the ancient Greek philosophers.

Early ThoughtsEarly ThoughtsEarly ThoughtsEarly Thoughts

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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.

Early ThoughtsEarly ThoughtsEarly ThoughtsEarly Thoughts

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

Dalton’s Model of the AtomDalton’s Model of the Atom

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2. Atoms of the same element are alike in mass and size.

3. Atoms of different elements have different masses and sizes.

4. Chemical compounds are formed by the union of two or atoms of different elements.

Dalton’s Atomic Theory

Modern research has demonstrated that atoms are composed of subatomic particles.

Atoms under special circumstances can be decomposed.

1. Elements are composed of minute indivisible particles called atoms.

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Dalton’s Atomic Theory

5. Atoms combine to form compounds in simple numerical ratios, such as one to one , two to two, two to three, and so on.

6. Atoms of two elements may combine in different ratios to form more than one compound.

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5.1

Dalton’s atoms were individual particles.

Atoms of each element are alike in mass and size.

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5.1

Dalton’s atoms were individual particles.

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

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Daltons atoms combine in specific ratios to form compounds.

H 2 =

O 1H 1

= O 1

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The Law of Definite Composition

A compound always contains two or more elements combined in a definite proportion by mass.

Composition of CompoundsComposition of Compounds

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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.

• Water always contains the same two elements: hydrogen and oxygen.

Composition of Water

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The Law of Multiple Proportions

Atoms of two or more elements may combine in different ratios to produce more than one compound.

Composition of CompoundsComposition of Compounds

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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.

• Hydrogen peroxide always contains the same two elements: hydrogen and oxygen.

Composition of Hydrogen Peroxide

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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.

Combining Ratios of Hydrogen and Oxygen

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Composition of CompoundsComposition of Compounds

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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).

Properties of Electric Charge

1 22

kq qF =

r

q1 and q2 are charges, r is the distance between charges and k is a constant.

The Nature of Electric ChargeThe Nature of Electric ChargeThe Nature of Electric ChargeThe Nature of Electric Charge

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• 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 IonsDiscovery of IonsDiscovery of IonsDiscovery of Ions

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• 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

Discovery of IonsDiscovery of IonsDiscovery of IonsDiscovery of Ions

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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-

Discovery of IonsDiscovery of IonsDiscovery of IonsDiscovery of Ions

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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.

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 Parts of the AtomSubatomic Parts of the Atom

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• 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.

Subatomic Particle - Electron

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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 - ElectronSubatomic Particle - ElectronSubatomic Particle - ElectronSubatomic Particle - Electron

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This was the discovery of the This was the discovery of the fundamental unit of charge fundamental unit of charge

– the electron.– the electron.

Subatomic Particle - ElectronSubatomic Particle - Electron

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• 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 - ProtonSubatomic Particle - ProtonSubatomic Particle - ProtonSubatomic Particle - Proton

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• James Chadwick discovered the neutron in 1932.

Its actual mass is slightly greater than the mass of a proton.

Subatomic Particle - NeutronSubatomic Particle - NeutronSubatomic Particle - NeutronSubatomic Particle - Neutron

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Subatomic ParticlesSubatomic Particles

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• 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.

IonsIonsIonsIons

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5.4

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

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5.4

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

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• 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.

The Nuclear AtomThe Nuclear AtomThe Nuclear AtomThe Nuclear Atom

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• 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.

The Rutherford ExperimentThe Rutherford ExperimentThe Rutherford ExperimentThe Rutherford Experiment

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Rutherford’s alpha particle scattering experiment.5.5

The Rutherford ExperimentThe Rutherford Experiment

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• 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 ExperimentThe Rutherford ExperimentThe Rutherford ExperimentThe Rutherford Experiment

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• 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.

The Rutherford ExperimentThe Rutherford ExperimentThe Rutherford ExperimentThe Rutherford Experiment

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

The Rutherford ExperimentThe Rutherford ExperimentThe Rutherford ExperimentThe Rutherford Experiment

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5.5Deflection and scattering of alpha particles by positive gold nuclei.

Deflection

Scattering

The Rutherford ExperimentThe Rutherford Experiment

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• 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 General Arrangement of Subatomic ParticlesSubatomic ParticlesGeneral Arrangement of General Arrangement of Subatomic ParticlesSubatomic Particles

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• 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

General Arrangement of General Arrangement of Subatomic ParticlesSubatomic ParticlesGeneral Arrangement of General Arrangement of Subatomic ParticlesSubatomic Particles

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5.6

General Arrangement of General Arrangement of Subatomic ParticlesSubatomic ParticlesGeneral Arrangement of General Arrangement of Subatomic ParticlesSubatomic Particles

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• 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.

Atomic Numbers of the ElementsAtomic Numbers of the ElementsAtomic Numbers of the ElementsAtomic Numbers of the Elements

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

Every hydrogen atom contains 1 proton in its nucleus.

Atomic Numbers of the ElementsAtomic Numbers of the Elements

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1 proton in the nucleus

atomic number

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

1H

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

Every carbon atom contains 6 protons in its nucleus.

Atomic Numbers of the ElementsAtomic Numbers of the Elements

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6 protons in the nucleus 6C

atomic number

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

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92 protons in the nucleus 92U

atomic number

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

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• 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.

Atomic Numbers of the ElementsAtomic Numbers of the ElementsAtomic Numbers of the ElementsAtomic Numbers of the Elements

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Isotopes of the Same Element Have

Equal numbers of protons

Different numbers of neutrons

Atomic Numbers of the ElementsAtomic Numbers of the ElementsAtomic Numbers of the ElementsAtomic Numbers of the Elements

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Isotopic Notation

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Isotopic Notation

6 protons

C6126 protons + 6 neutrons

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Isotopic Notation

C66 protons

146 protons + 8 neutrons

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Isotopic Notation

88 protons

168 protons + 8 neutrons

O

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Isotopic Notation

88 protons

178 protons + 9 neutrons

O

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Isotopic Notation

88 protons

188 protons + 10 neutrons

O

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Hydrogen has three isotopes

1 proton

0 neutrons

1 proton

1 neutron

1 proton

2 neutrons

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Examples of Isotopes

Element Protons Electrons Neutrons SymbolHydrogen 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

H11

H21

U235192

H31

U236192

Cl3717

Cl3517

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• 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.

Atomic MassAtomic MassAtomic MassAtomic Mass

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A Modern Mass Spectrometer

A mass spectrogram is recorded.

From the intensity and positions of the lines on the mass spectrogram, the different isotopes and their relative amounts can be determined.

Positive ions formed from sample. Electrical field

at slits accelerates positive ions.

Deflection of positive ions occurs at magnetic field.

5.8

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5.9

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

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

Atomic Mass

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.

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126C

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.

Atomic MassAtomic MassAtomic MassAtomic Mass

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 12

1

1 amu = 1.6606 x 10-24 g

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H Average atomic mass 1.00797 amu.

Atomic MassAtomic MassAtomic MassAtomic Mass

K Average atomic mass 39.098 amu.

U Average atomic mass 248.029 amu.

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• 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 12.0000…amu).

Average Relative Atomic Mass

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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 amu63.55 amu

IsotopeIsotopic mass

(amu)Abundance

(%)

Average atomic mass

(amu)

62.9298 69.09

64.9278 30.91

6329 Cu6529 Cu

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ConceptsDalton’s Atomic Mode;Law of Definite CompositionLaw of Multiple ProportionsThree principle subatomic particlesThomson Model of the AtomRutherford alpha-scattering experimentAtomic Number, Mass number, number of neutrons, number of Protons, Number of ElectronsThree Isotopes of HydrogenAverage Atomic Mass of an Element