The Structure of the Atom Chapter 4

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The Structure of the Atom Chapter 4

BIG IDEA

Atoms are the fundamental

building blocks of matter.

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I. Early Ideas About Matter

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Section I Overview

The Gist

 The ancient Greeks tried to explain

matter, but the

scientific study of the atom began with John Dalton in the early

1800s.

Essential Questions

• What are the similarities and differences of the atomic models of Democritus,

Aristotle, and Dalton?

• How was Dalton’s theory used to

explain the

conservation of mass?

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A. The Roots of Atomic Theory

1. Democritus (460 - 370 B.C.)

• 1st to propose that matter is

made of indivisible particles that move through empty space

• átomos = uncuttable

• Different atoms have different sizes and properties

Democritus, the

“laughing philosopher”

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2. Aristotle (384 - 322 B.C.)

• Disagreed with Democritus

• Rejected the idea of atoms

• Empty space cannot exist

• Views went unchallenged for 2,000 years

Earth, fire, air & water.

“Elements” supported by Aristotle.

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3. (John) Dalton's Atomic Theory (1803)

 Matter is made of indivisible particles called atoms

 Atoms cannot be created or destroyed

 All atoms of the same element are

identical in size, mass, and properties

 Atoms of different elements are different

 During chemical reactions atoms combine, separate, and rearrange

 Atoms of different elements combine in simple, whole number ratios to

form new substances called compounds.

English Quaker John Dalton (1766-

1844) revived the idea of the atom based on his own

and others’

experiments

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B. Conservation of Mass

• Dalton’s Atomic Theory explained conservation of mass

– “Atoms cannot be created or destroyed in a chemical reaction”

• The number of atoms is the same before and after a chem. rxn

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II. Defining the Atom II. Defining the Atom

This image of a single suspended sheet of graphene taken with TEAM 0.5, shows individual carbon

atoms (yellow) on the honeycomb lattice.

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Section II Overview

The Gist

• An atom is made up of a nucleus

containing protons and neutrons, with electrons moving around the nucleus.

• What is an atom?

• How can the subatomic particles be distinguished in terms of relative charge and mass?

• Where are the locations of the subatomic particles within the structure of the atom?

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A. The Atom

• Smallest part of an element that retains the properties of that element

• One of the only ways to observe atoms is with a scanning tunneling microscope

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B. The Electron (e^-)

1. Discovery of the Electron: J.J. Thomson’s Cathode Ray Tube Experiments (1897)

• Measured charge-to-mass ratio of “cathode-ray particle” (1.759 x 10⁸ C/g)

• Work suggested that atoms were not “indivisible”

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2. “Plum Pudding” Model

• In 1904 J.J.T. proposed that atoms are spherical structures composed of a uniformly distributed positive (+) charge in which e^-s reside

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3. Oil-Drop Experiment

• Robert A. Millikan (1909) determined charge of e^-

• Used Thomson’s c/m equation to calculate the

mass of an e^-

– About 1,840 times less than the

mass of H atom

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• Charge of -1 e (-1.602 x 10^-19 C)

• Almost zero mass

• Responsible for chem & phys properties

• Located outside the nucleus in electron cloud

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4. Electron Cloud Model

• Describes a probability of where e^-s might be located

• The position of e- cannot be known

• Nucleus (center) is roughly 100,000 times smaller than electron cloud

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C. The Nucleus

1. Gold Foil Experiment

• Ernest Rutherford (1911) experimented with alpha (α) particles, positively-charged particles given off by certain radioactive materials

• Directed beam of α particles at thin gold foil surrounded by a screen coated w/ zinc sulfide

– light produced when α particle hit screen

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• Almost all α particles passed through

undeflected

• Very few (about 1 in

20,000) deflected at an angle

• A few bounced back

Conclusions:

• center of an atom is very small, dense &

positively (+) charged

• electrons surround the positive (+)

center

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C. The Nucleus

2. Rutherford’s (Planetary) Model of the Atom

• Atoms are mostly empty space

• Nucleus - an extremely small, positively (+) charged, dense center of an atom

• e^-s located outside of nucleus

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C. The Nucleus

3. The Proton (p⁺) & the Neutron (n⁰)

• Proton - located in the nucleus (center)

• Discovered by Eugene Goldstein in 1886

• Positive (+1) charge is equal and opposite of an e^-

• relative mass = 1 (about 1,836 times more than e^-)

• # of p⁺ determines the element

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C. The Nucleus

3. The Proton (p⁺) & the Neutron (n⁰)

• neutron - located in the nucleus

• mass nearly equal to p⁺

• no charge

• relative mass = 1

• n⁰ stabilize the nucleus

• discovered by James Chadwick (1932)

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D. (Rutherford)-Bohr Model

1. Neils Bohr (1913) expanded on Rutherford’s model

• small, positively-charged nucleus surrounded by negative electrons

• e^-s travel in specific, circular orbits (shells) around the nucleus

• Later refined (quantum mechanics)

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E. Electrical Forces

• opposite charges attract (+ , -)

• like charges repel (- , -) or (+ , +)

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F. Summary of atoms

• subatomic particles (p⁺, n⁰ & e^-) are still being researched

• p⁺ & n⁰ are made of smaller particles called quarks

• e^- are responsible for the behavior of atoms

• Almost all of the mass of an atom is in the tiny nucleus

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E. Summary of atoms

Particle Symbol Location

Charge Mass

e coulombs Relative Actual (g) Electron e^- Cloud -1 -1.602 x 10^-19 1/1840 9.11 x 10^-28

Proton p⁺ Nucleus +1 +1.602 x 10^-19 1 1.673 x 10^-24

Neutron n⁰ Nucleus 0 0 1 1.675 x 10^-24

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III. How Atoms Differ

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Section III Overview

The Gist

• The number of protons and the mass number define the type of atom.

• How is the atomic number used to

determine the identity of an atom?

• Why are atomic masses not whole numbers?

• Given the mass number and atomic number, how are the number of

electrons, protons, and neutrons in an atom calculated?

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A. Atomic Number

• # of p⁺ in the nucleus

• # of p⁺ determines the element's ID

• the periodic table is arranged by atomic #

• since atoms are (usually) electrically neutral that means that there is an equal # of p⁺s and e^-s

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B. Isotopes

• atoms with the same # of p⁺ but a different # of n⁰

• K-39 has 19 p⁺ and 20 n⁰

• K-40 has 19 p⁺ and 21 n⁰

• K-41 has 19 p⁺ and 22 n⁰

• isotopes are chemically alike because they have the same # of p⁺s

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B. Isotopes

1. Mass of Isotope

• isotopes containing more n⁰ have greater mass

• isotopes have the SAME chem behavior

• e^-s determine the behavior

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B. Isotopes

2. Isotope Notation

• mass number = protons + neutrons

• the sum of all nucleons

• since e^-s have such a small mass they are ignored

• shorthand: C-14

• shorthand:

• calculating the number of neutrons:

• (mass #) - (atomic #) = number of n⁰s

Mass Number --> 14 Atomic Number --> 6

C

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B. Isotopes

3. Natural Abundance of Isotopes

• in nature most elements are found as a mixture of isotopes

• Ex: potassium (atomic # 19)

• 93.26% have 20 n⁰ (K-39)

• 6.73% have 21 n⁰ (K-40)

• 0.01% have 22 n⁰ (K-41)

K-39 K-40 K-41

0 10 20 30 40 50 60 70 80 90 100

Isotopes of Potassium

% Abundance

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C. Mass of Atoms

1. Atomic Mass Unit (amu)

• since the mass of atoms is extremely small we will NOT use grams

• C-12 is used as a reference isotope

• atomic mass unit (amu) - 1/12 the mass of a C-12 atom

• the mass of a single p⁺ or n⁰ is about 1 amu

• the mass of an e- is about 1/1840 amu

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2. Atomic Mass (Weight)

• the weighted average mass of all naturally occurring isotopes of an element

• atomic masses do not tend to be whole #s

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2. Atomic Mass

• Ex: hydrogen

• H has an atomic mass of 1.0079 amu

• this means that H-1 is much more abundant than other isotopes such as H-2 and H-3

• H-1 (1.0078) 99.985%

• H-2 (2.0141) 0.015%

• H-3 (3.0160) negligible%

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3. Calculating Atomic Mass

1) Multiply the isotopes mass by the relative abundance (% as a decimal)

2) The atomic mass is the sum of these contributions

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3. Calculating Atomic Mass

• Ex: Copper has 2 isotopes:

• Cu-63 (62.930 amu) 69.2%

• Cu-65 (64.928 amu) 30.8%

• (62.930 amu)(.692) = 43.5 amu

• (64.928 amu)(.308) = 20.0 amu

• 43.5 amu + 20.0 amu = 63.5 amu

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IV. Unstable Nuclei and Radioactive Decay

Cherenkov Radiation is electromagnetic radiation

emitted when a charged particle (such as an electron)

passes through a dielectric medium at a speed greater

than the phase velocity of light in that medium.

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Section IV Overview

The Gist

• Unstable atoms emit

radiation to gain stability.

• What is the relationship between unstable nuclei and radioactive decay?

• How are alpha, beta, and gamma radiation

characterized in terms of mass and charge?

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A. Radioactivity

• the process by which certain elements spontaneously emit radiation

• caused by an unstable nucleus

• Radiation - the rays and particles emitted by a radioactive material

Stable elements

Radioactive elements with very long-lived isotopes

Radioactive elements that may present low health hazards.

Radioactive elements that are known to pose high safety risks.

Highly radioactive elements.

Extremely radioactive element. s

The trefoil symbol is used to indicate radioactive material.

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A. Radioactivity

2. Radioactive Decay

• an unstable nucleus loses E by emitting radiation

• atoms will decay until they form a stable atom

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B. Types of Radiation

1. Alpha Particles (α or He)

• 2 p⁺ and 2 n⁰

• (+2) charge 4 amu

• He-4 nucleus

• not much penetrating power

• stopped by a piece of paper or 0.5 mm body tissue

• can damage soft tissue

4 2

Litvinenko suddenly fell ill and was hospitalized. He died three weeks later,

becoming the first confirmed victim of lethal polonium-210-induced acute radiation syndrome. Po-210 is an alpha emitter that has a half-life of 138.4 days; it

decays directly to its stable daughter isotope, ²⁰⁶Pb.

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1. Alpha Particles (α or He)

• nuclear equation - shows the atomic #s and mass

#s of the particles involved

• ex: U Th + He (α-particle)

4 2

238 92

234 90

4 2

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2. Beta Particles (β or e)

• fast-moving e^-

• (-1) charge

• 0 amu (1/1840 amu)

• more penetrating power than an α-particle

• stopped by metal Al foil or 4 mm body tissue

• ex: n H + e (β-particle)¹₀ ¹₁ ⁰_-1

0 -1

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3. Gamma Radiation (γ)

• nuclear equation - shows the atomic #s and mass

#s of the particles involved

• high energy electromagnetic (EM) wave

• no charge

• no mass

• very similar to X-rays

• have EXTREME penetrating power

• potentially very dangerous to living organisms

• ex: Th Ra + He + γ²³⁰

90 226

88

4 2

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This diagram demonstrates the constitution of different kinds of ionizing radiation and

their ability to penetrate matter. Alpha particles are stopped by a sheet of paper whilst beta particles halt to an

aluminium plate. Gamma radiation is dampened when it

penetrates matter. Gamma rays can be stopped from 4 meters of lead. Tungsten and

tungsten alloys can stop Gamma radiation with much

less mass then lead

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4. Nuclear Stability

• stability of the nucleus depends on the n⁰ to p⁺ ratio