the development of atomic theory larry scheffler lincoln high school portland or 1

The Development of The Development of Atomic TheoryAtomic Theory

Larry SchefflerLincoln High School Portland OR

1

The The AtomAtom

• The term atom is derived from the Greek word (atomos) meaning invisible

• Democritius (470-370 BC ) suggested that all matter was made up of invisible particles called atoms

2

Law of Constant Law of Constant CompositionComposition

A compound always contains atoms of two or More elements combined in definite proportions by mass

Example:

Water H22O always contains 8 grams of oxygen to 1 gram of hydrogen

3

Law of Multiple Law of Multiple ProportionsProportions

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

Examples:

NO NONO NO22 N N22O NO N22OO55

4

Dalton’s Atomic Dalton’s Atomic TheoryTheory1. All elements are

composed of indivisible and indestructible particles called atoms.

2. Atoms of the same element are exactly alike, They have the same masses. 3. Atoms of different

elements have different masses.

4. Atoms combine to form compounds in small whole number ratios..5

Objections to Dalton’sObjections to Dalton’sAtomic TheoryAtomic Theory

• Atoms are not indivisible. They are composed of subatomic particles.• Not all atoms of a particular element have exactly the same mass. • Some nuclear transformations alter (destroy) atoms

6

Crookes ExperimentCrookes Experiment

Crookes found that passing an electrical current Crookes found that passing an electrical current through a gas at very low pressure caused the gas to through a gas at very low pressure caused the gas to glow. Putting a magnet next to the beam caused it to glow. Putting a magnet next to the beam caused it to be deflected.be deflected. 7

The ElectronThe Electron1. The electron was the first subatomic

particle to be identified.2. In 1897 J.J Thomson used a cathode ray

tube to establish the presence of a charged particle known as the electron

3. Thomson established the charge to mass ratio

E/m = 1.76 x 108 coulombs/gram

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A Cathode Ray TubeA Cathode Ray Tube

Thomson found that an electrical field would Thomson found that an electrical field would also deflect an electron beam. He surmised also deflect an electron beam. He surmised that the ratio of charge to mass is constant.that the ratio of charge to mass is constant.

Thomson’s Charge to Thomson’s Charge to Mass RatioMass Ratio

E/m = 1.76 x 108 coulombs/gram

Thomsen’s Plum Thomsen’s Plum Pudding ModelPudding Model

Thompson proposed that an atom was made up of electrons scattered unevenly through out an elastic sphere. These charges were surrounded by a sea of positive charge to balance the electron's charge like plums surrounded by pudding.

This early model of the atom was called The Plum Pudding Model. A more contemporary American label might be the “chocolate chip cookie” model 11

Millikan’s ExperimentMillikan’s Experiment

By varying the charge on the plates, Millikan found By varying the charge on the plates, Millikan found that he could suspend the oil drops or make them that he could suspend the oil drops or make them levitate.levitate. 12

Millikan’s ExperimentMillikan’s ExperimentMillikan used his data to measure the charge of an electron and then to calculate the mass of the electron from Thomson’s charge to mass ratio.

Given the charge = 1.60 x 10-19 coulomb and the ratio of E/m = 1.76 x 108 coulombs/gram it is possible to calculate the mass

Mass

= 9.11 x 10-28 gram

13

ProtonsProtons

First observed by E. Goldstein in 1896J.J. Thomson established the presence of positive charges.

The mass of the proton is1.673 x 10-24 grams

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Rutherford’s ExperimentRutherford’s Experiment

1910

Ernest Rutherford

Rutherford oversaw Geiger and Marsden carrying out his famous experiment.

They fired high speed alpha particles (Helium nuclei) at a piece of gold foil which was only a few atoms thick.

They found that although most of them passed through. About 1 in 10,000 hit and were deflected

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Rutherford’s ExperimentRutherford’s Experiment

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Rutherford’s Rutherford’s ExperimentExperiment

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Rutherford’s Rutherford’s ExperimentExperiment

By studying this By studying this pattern, Rutherford pattern, Rutherford concluded that concluded that atoms have a very atoms have a very dense nucleus, but dense nucleus, but there are mostly there are mostly empty space.empty space.

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Subatomic ParticlesSubatomic Particles

The diameter of a single atom rangesFrom 0.1 to 0.5 nm. (1 nm = 10-9 m).

Within the atom are smaller particles:ElectronsProtonsNeutrons

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NeutronsNeutrons

Discovered by James Chadwick in 1932

Slightly heavier than a proton

Mass of a neutron = 1.675 x 10-24 grams

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The Bohr ModelThe Bohr ModelNiels Bohr proposed the Planetary Model in 1913. Electrons move in definite orbits around the nucleus like planets moving around the nucleus. Bohr proposed that each electron moves in a specific energy level.

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Aspects of the Bohr Aspects of the Bohr ModelModel

Bohr put together Balmer’s and Plank’s discoveries to form a new atomic model

In Bohr’s model:

1. Electrons can orbit only at certain allowed distances from the nucleus.

2. Electrons that are further away from the nucleus have higher energy levels (explaining the faults with Rutherford’s model).

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The Electromagnetic SpectrumThe Electromagnetic Spectrum

Wave CharacteristicsWave Characteristics

Energy of a waveE = h

Frequency = = number of peaks per unit of time

Speed of lightc =

Emission SpectraEmission Spectra

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Flame TestsFlame Tests

According to BohrAccording to Bohr

Atoms radiate energy whenever an electron jumps from a higher-energy orbit to a lower-energy orbit. Also, an atom absorbs energy when an electron gets boosted from a low-energy orbit to a high-energy orbit.

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Problems with the Bohr Problems with the Bohr ModelModel

The Bohr model provided a model that gave The Bohr model provided a model that gave precise results for simple atoms like hydrogen.precise results for simple atoms like hydrogen.

Using the Bohr model precise energies could Using the Bohr model precise energies could be calculated for energy level transitions in be calculated for energy level transitions in hydrogen.hydrogen.

Unfortunately these calculations did not work Unfortunately these calculations did not work for atoms with more than 1 electron.for atoms with more than 1 electron.

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Weakness of the Bohr Weakness of the Bohr ModelModel

• According to the Bohr model electrons could be found in orbitals with distinct energies.

• When the data for energies measured using spectral methods where compared to the values predicted by the Rydberg equation, they were accurate only for hydrogen.

• By the 1920s, further experiments showed that Bohr's model of the atom had some difficulties. Bohr's atom seemed too simple to describe the heavier elements.

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Modern View of the Modern View of the AtomAtom

The wave mechanical model for the atom was developed to answer some of the objections that were raised about the Bohr model. It is based on the work of a number of scientists and evolved over a period of time

The quantum theorists such as Maxwell Planck suggested that energy consists of small particles known as photons. These photons can have only discreet energies

Maxwell Planck

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Albert Einstein demonstrated the equivalence of matter and energy. Hence matter and energy in Einstein’s theory were not different entities but different expressions of the same thing

Einstein then proposed the equivalence of Matter and Energy given by his famous equation

E = mc2

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Louis de Broglie suggested that if energy could be thought of as having particle properties, perhaps matter could be thought of as having wave like characteristics

Louis de Broglie

32


Louis de Broglie proposed that an electron is not just a particle but it also has wave characteristics.

E = mc2 = h

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Heisenberg proposed that it was impossible to know the location and the momentum of a high speed particle such as an electron.

The more precisely the position is determined, the less precisely the momentum is known in this instant, and vice versa.

--Heisenberg, Uncertainty paper, 1927

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The atom cannot be defined as a solar system with discreet orbits for the electrons. The best that we could do was define the probability of finding an electron in a particular location.

The more precisely the position is determined, the less precisely the momentum is known in this instant, and vice versa. --Werner Heisenberg,

Uncertainty paper, 1927

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Edwin Schroedinger proposed that the electron is really a wave. It only exists when we identify its location. Therefore the electrons are best thought of probability distributions rather than discreet particles.

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Modern View of the AtomModern View of the Atom

The modern view of the atom suggests The modern view of the atom suggests that the atom is more like a cloud. that the atom is more like a cloud. Atomic orbitals around the nucleus Atomic orbitals around the nucleus define the places where electrons are define the places where electrons are most likely to be found. most likely to be found.

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Wave Mechanical ModelWave Mechanical ModelThe location of the The location of the electron in a hydrogen electron in a hydrogen atom is a probability atom is a probability distribution.distribution.

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Progression of Atomic ModelsProgression of Atomic Models

Our view of the atom has changed over timeOur view of the atom has changed over time 39

ATOMIC STRUCTUREATOMIC STRUCTURE

Particle

proton

neutron

electron

Charge

+ charge

- charge

No charge

1

1

0

Mass

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ATOMIC NUMBER AND MASS NUMBERATOMIC NUMBER AND MASS NUMBER

the number of protons in an atom

the number of protons and neutrons in an atomHH

ee22

44

Atomic Number

Mass Number

Number of electrons = Number of protons

in a neutral atom 41

Atomic MassAtomic MassThe atomic mass of an atom is a relative number that is used to compare the mass of atoms.

An atomic mass unit is defined as 1/12 of the mass of an atom of carbon 12.

The atomic masses of all other atoms are a ratio to carbon 12

42

IsotopesIsotopesMany elements have atoms that have multiple forms

Different forms of the same element having different numbers of neutrons are called isotopes.

For example: Carbon exists as both Carbon 12 and Carbon 14

Carbon 12 Carbon 14

6 electrons 6 electrons

6 protons 6 protons

6 neutrons 8 neutrons 43

Isotopes and Atomic Isotopes and Atomic MassMass

Many elements have atoms that have multiple isotopes.

Isotopes vary in abundance. Some are quite common while others are very rare.

The atomic mass that appears in the periodic table is a weighted average taking into account the relative abundance of each isotope.

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Isotope:Isotope: one of two or more atoms having the one of two or more atoms having the same number of protons but different same number of protons but different numbers of neutronsnumbers of neutrons

or Na-23 or Na-24

Measuring Atomic MassMeasuring Atomic Mass--the Mass Spectrometer--the Mass Spectrometer

The mass The mass spectrometespectrometer can be r can be used to used to determine determine the atomic the atomic mass of mass of isotopes.isotopes.

Mass Spectrum of NeonMass Spectrum of Neon The mass spectrum neon shows three isotopes with the The mass spectrum neon shows three isotopes with the

isotope at atomic mass = 20 accounting for more than isotope at atomic mass = 20 accounting for more than 90% of neon.90% of neon.

Mass Spectrum of GermaniumMass Spectrum of Germanium The mass spectrum of germanium shows 5 peaks at The mass spectrum of germanium shows 5 peaks at

relative atomic masses of 70, 72,73,74, and 75relative atomic masses of 70, 72,73,74, and 75

Calculating the average Calculating the average relative atomic massrelative atomic mass

The average atomic mass that is shown in the The average atomic mass that is shown in the periodic table is really the weighted average of periodic table is really the weighted average of the atomic masses of each of the elements the atomic masses of each of the elements isotopes. Germanium has 5 isotopes whose isotopes. Germanium has 5 isotopes whose relative atomic masses are shown in the tablerelative atomic masses are shown in the table

Mass Number % Abundance 70 20.55 72 27.37 73 7.67 74 36.74 75 7.67

Calculating the Average Calculating the Average Relative Atomic MassRelative Atomic Mass

To calculate the average atomic mass multiply the To calculate the average atomic mass multiply the atomic mass of each isotope by its abundance atomic mass of each isotope by its abundance (expressed as a decimal fraction)(expressed as a decimal fraction)

Mass Number % Abundance 70 20.55 72 27.37 73 7.67 74 36.74 75 7.67

Average atomic mass = (0.2055)(70) + (0.2737)(72) + (0.0767)(73) + (0.3674)(74)+ (0.0767)(75)

= 72.36Note: atomic masses are ratios so they do not have real units although they are sometimes called atomic mass units or amu

ProblemProblem The mass spectrum of an element, The mass spectrum of an element, AA, contained , contained

4 lines at mass/charge ratios of 54, 56, 57 and 4 lines at mass/charge ratios of 54, 56, 57 and 58 with relative intensities of 5.84, 91.68, 2.17 58 with relative intensities of 5.84, 91.68, 2.17 and 0.31 respectively. and 0.31 respectively. CalculateCalculate the relative the relative atomic mass of element A.atomic mass of element A.

Average atomic mass = (0.0584)(54) + (0.9168)(56) + (0.0217)(57) + (0.0031)(58)

= 56.02

The NucleusThe Nucleus

The The nucleus nucleus is very small — of the order is very small — of the order of 10of 10-14-14 meter whereas the atom is of the meter whereas the atom is of the order of 10order of 10-9-9 meters. By analogy, the meters. By analogy, the nucleus occupies as much of the total nucleus occupies as much of the total volume of the atom as a fly in a cathedralvolume of the atom as a fly in a cathedral

Protons and NeutronsProtons and Neutrons Protons and neutrons have nearly equal masses, and Protons and neutrons have nearly equal masses, and

their combined number, their combined number, the the mass numbermass number, is , is approximately equal to approximately equal to the the atomic massatomic mass of an atom. of an atom.

The combined mass of the electrons is very small in The combined mass of the electrons is very small in comparison to the mass of the nucleus, since protons comparison to the mass of the nucleus, since protons and neutrons weigh roughly 2000 times more than and neutrons weigh roughly 2000 times more than electrons.electrons.

Charge Mass Relative Mass

Electron -1 9.109383 x 10-28 g 1/1837

Proton +1 1.6726217 x 10-24 g 1

Neutron 0 1.6749273 x 10-24 g 1

Atomic Mass UnitsAtomic Mass Units An atomic mass unit (amu) is equal to exactly 1/12 of An atomic mass unit (amu) is equal to exactly 1/12 of

the mass of an atom of Carbon 12.the mass of an atom of Carbon 12. One atomic mass unit is equal to 1.66054 x 10One atomic mass unit is equal to 1.66054 x 10-24-24 grams. grams.

Note that this is slightly less than the mass of a proton Note that this is slightly less than the mass of a proton or a neutron. or a neutron.

An atomic mass unit is sometimes called a An atomic mass unit is sometimes called a Dalton (D). Dalton (D). 1.00 g = 6.02214 x 101.00 g = 6.02214 x 102323 amu. This number is also known amu. This number is also known

as Avogadro’s Number and it defines the size of a as Avogadro’s Number and it defines the size of a quantity we call a mole.quantity we call a mole.

Radioactive NucleiRadioactive Nuclei The presence of neutrons in the nucleus tends The presence of neutrons in the nucleus tends

to buffer the repulsions of multiple protons in the to buffer the repulsions of multiple protons in the nucleus.nucleus.

There appears to be an optimal number of There appears to be an optimal number of neutrons for the number of protons in a given neutrons for the number of protons in a given atom in a stable atom.atom in a stable atom.

In a radioactive element the nucleus may In a radioactive element the nucleus may disintegrate releasing either an alpha particle or disintegrate releasing either an alpha particle or a beta particle as well as some high energy a beta particle as well as some high energy gamma radiation.gamma radiation.

Alpha ParticlesAlpha Particles An alpha particle consists of two protons and two neutrons. This An alpha particle consists of two protons and two neutrons. This

makes it equivalent to a helium nucleus. makes it equivalent to a helium nucleus. When a radioactive element undergoes alpha decay its nucleus is When a radioactive element undergoes alpha decay its nucleus is

decreased in mass by 2 protons and 2 neutrons. decreased in mass by 2 protons and 2 neutrons. Since the number of protons changes, it has a new atomic number Since the number of protons changes, it has a new atomic number

and hence it is a different element. The mass number decreases and hence it is a different element. The mass number decreases by 4. by 4.

238U -->

234Th +

4He

92 90 2

146 neutrons 92 protonsRatio = 1.52

144 neutrons 90 protonsRatio = 1.60

Alpha decay raises the neutron to proton ratio. It occurs in radioactive nuclei where the ratio is too low

Beta ParticlesBeta Particles A beta particle consists of a high-speed electron A beta particle consists of a high-speed electron

released from the nucleus. released from the nucleus. When a radioactive element undergoes beta decay, When a radioactive element undergoes beta decay,

the number of protons increases by one and the the number of protons increases by one and the number of neutrons decreases by one. The mass number of neutrons decreases by one. The mass number remains the same.number remains the same.

14C -->

14N +

0e-

6 7 -1

8 neutrons6 protonsRatio = 1.333

7 neutrons7 protonsRatio = 1.00

Beta decay lowers the neutron to proton ratio. It occurs in radioactive nuclei where the ratio is too high

The Half-LifeThe Half-LifeThe rates at which various radioactive elements undergo decay The rates at which various radioactive elements undergo decay vary considerably. The half-life of a radioactive element if the vary considerably. The half-life of a radioactive element if the time required for half of the nuclei in a sample of radioactive time required for half of the nuclei in a sample of radioactive nuclei to disintegrate.nuclei to disintegrate.

Isotope Type Half-life

Uranium 238 Alpha 4.51 x 109

Carbon 14 Beta 5730 years

Iodine 131 Beta 8 days

Radon-222 Alpha 3.825 days

Cesium -137 Beta 30 years

Polonium 210 Alpha 138 days

the development of atomic theory larry scheffler lincoln high school portland or 1

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