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Section 3.1—Substances Are Made of Atoms

Objectives:

1. State the three laws that support the existence of atoms.

2. List the five principles of John Dalton’s atomic theory.

Vocabulary:

law of definite proportions

law of conservation of mass

law of multiple proportions

Atomic Theory

Atomic theory, which states that atoms are the building blocks of matter, has been around

as long ago as 400 BCE. However, scientific experiments supporting the existence of

atoms did not appear until the late 17th century.

Experimental results show that the following laws support the current atomic theory:

Law of Definite Proportions—each specific chemical compound always contains

exactly the same elements in exactly the same proportions by number, or by mass

(or by weight), or by volume.

Law of Conservation of Mass—mass cannot be created or destroyed in normal

chemical reactions (the mass of the reactants equals the mass of the products).

Law of Multiple Proportions—atoms combine in whole number ratios.

Dalton’s Atomic Theory

In 1808, John Dalton, an English high school teacher, used the Greek concept of the atom

and the law of definite proportions, the law of conservation of mass, and the law of

multiple proportions to develop an atomic theory. According to Dalton, elements are

composed of one type of atom and compounds are composed of two or more types of

atoms. Dalton’s atomic theory can be summarized by the following statements:

1. All matter consists of atoms that cannot be divided, created, or destroyed.

2. Atoms of the same element are identical in their physical and chemical properties.

3. Atoms of different elements are chemically and physically different.

4. Atoms of different elements combine in simple, whole number ratios to form

compounds.

5. In chemical reactions, the atoms in compounds will separate, rearrange, and

combine but atoms are never created, destroyed, or changed.

The huge advance in human thinking about these statements that these propositions can be

tested. Note that propositions 1 and 2 are no longer true.

YouTube

http://www.youtube.com/watch?v=L2KmCTst2o0 (2 min)

http://www.youtube.com/watch?v=mXMFSVowbSc (4 min)

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Section 3.2—Structure of Atoms

Objectives:

1. Describe the evidence for the existence of electrons, protons, and neutrons, and

describe the properties of these subatomic particles.

2. Discuss atoms of different elements in terms of their numbers of electrons, protons,

and neutrons, and define the terms atomic number and mass number.

3. Define isotope, and determine the number of particles in the nucleus of an isotope.

Vocabulary:

electron

nucleus

proton

neutron

atomic number

mass number

isotope

charge

Historical Development of Atomic Theory

Experiments were conducted in the mid-19th to early 20th century that led to the discovery

that atoms are themselves made of electrons, protons, and neutrons.

J.J. Thomson discovered the electron using a cathode ray tube

Electrons have negative charge because

they are emitted from the negatively

charged cathode. Thomson also placed a

paddle wheel in the tube and it turned,

which indicated that electrons are particles

with mass.

Key hypothesis: atoms have neutral charge,

so there must exist positively charged

subatomic stuff that balanced the negativity

of the electrons. Thomson proposed his

“Plum Pudding Model” of the atom.

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Ernest Rutherford’s Gold Foil Experiment

According to the plum pudding model, all alpha

particles should go straight through. But some were

deflected at backward angles! “It was like firing a 15

inch shell at paper tissue and it bounced back to hit

you.” Using the angle of deflection, scientists

determined:

A heavy, positively charged nucleus much

larger than the alpha particle.

The radius of the atom is 10,000 times larger

than that of the alpha particle. The atom is

mostly empty space! It consists of a tiny, very

dense nucleus surrounded by “distant”

electrons.

The nucleus contains protons (p+) that are

2000 times heavier than the electron (e-). The

nucleus also contains neutrons (no) that are

neutrally charged and mass the same as a

proton. (Irene Joliot-Curie, daughter of Marie

Curie, was a co-discoverer of neutrons when

alpha particles hit barium).

Even though protons have positive charge,

they form a stable nucleus! Neutrons hold

protons together in the nucleus. All atoms

with 2 or more protons in the nucleus also

have neutrons in the nucleus.

Atomic Number and Mass Number

Name Symbol Charge Common

Charge

Notation

Mass Common

Mass

Notation

Electron e- -1.602 x10-19 coulombs -1 9.109 x10-31 kg 0

Proton p+ +1.602 x10-19 coulombs +1 1.673 x10-27 kg 1

Neutron no 0 coulombs 0 1.675 x10-27 kg 1

Fe26

55.847

Mass Number = total

number of protons and

neutrons in the nucleus.

Knowing a mass number

does not help identify an

element since the

number of neutrons can

change for the same

element. (Also called the

atomic mass.)

Atomic Number =

number of protons in

the nucleus. The

atomic number is

unique to each

element; it defines

the type of element.

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Isotopes

The nucleus of the same element can have a different number of neutrons. Atoms of the

same element (equal number of protons in the nucleus) that have a different number of

neutrons are called isotopes. For example, there are two isotopes of helium: helium-3 and

helium-4.

All isotopes of an element have the same atomic number

Atomic mass numbers of isotopes of the same element are not the same because the

numbers of neutrons are different.

The mass number listed in the periodic table is an average of the mass numbers of

the naturally occurring isotopes of the element.

Name Symbol # neutrons # protons Mass # Abundance

Lead-204 Pb204

82 122 82 204 1.4%

Lead-206 Pb206

82 124 82 206 24.1%

Lead-207 Pb207

82 125 82 207 22.1%

Lead-208 Pb208

82 126 82 208 52.4%

The mass number used in the periodic table is determined as follows:

Mass = Σ (Abundance of Isotope)(Mass # of Isotope)

For example, the average atomic mass of the naturally occurring isotopes of lead is:

Mass # = (1.4%)(204) + (24.1%)(206) + (22.1%)(207) + (52.4%)(208) = 207.2

How do electrons stay close to the nucleus?

Electromagnetic force causes the negatively charged electrons to be attracted to the

positively charged nucleus. The equation for this attraction is called Coulomb’s Law:

𝐸𝑙𝑒𝑐𝑡𝑟𝑜𝑚𝑎𝑔𝑛𝑒𝑡𝑖𝑐 𝐹𝑜𝑟𝑐𝑒 (𝐹)

= 𝑐ℎ𝑎𝑟𝑔𝑒 𝑜𝑛 𝑜𝑛𝑒 𝑝𝑎𝑟𝑡𝑖𝑐𝑙𝑒 (𝑞1) 𝑥 𝑐ℎ𝑎𝑟𝑔𝑒 𝑜𝑛 𝑜𝑡ℎ𝑒𝑟 𝑝𝑎𝑟𝑡𝑖𝑐𝑙𝑒 (𝑞2)

(𝑑𝑖𝑠𝑡𝑎𝑛𝑐𝑒 𝑏𝑒𝑡𝑤𝑒𝑒𝑛 𝑐ℎ𝑎𝑟𝑔𝑒𝑑 𝑝𝑎𝑟𝑡𝑖𝑐𝑙𝑒𝑠)2=

𝑞1𝑥 𝑞2

𝑑2

How can proton form a stable nucleus?

Protons have a positive charge and exist together in the nucleus. The nature of magnetic

force predicts that they should repel! However, at very small distances the strong nuclear

force overcomes Coulombic repulsion and holds protons together. In addition, neutrons

stabilize the nucleus—neutrons are like “glue” holding the nucleus together.

YouTube

History of Atoms (Bozeman, 9:10)

5

http://www.youtube.com/watch?v=njGz69B_pUg

Atom are mostly empty space-Bill Nye (6:37)

http://www.youtube.com/watch?v=cnXV7Ph3WPk

Development of the Atomic Theory (9:52)

http://www.youtube.com/watch?v=eXdWlnBlncM

Cartoon of the History of the Atom Theory

http://www.youtube.com/watch?v=QbWKF9uDF7w

Chemistry Atomic Theory Timeline (2:07)

http://www.youtube.com/watch?v=WSY5H1k3TVU

Have you ever seen an atom? (Nature Video, 2:32)

http://www.youtube.com/watch?v=yqLlgIaz1L0

Early Atomic Theory: Dalton, Thomson, Rutherford, and Millikan

http://www.youtube.com/watch?v=AwJieYzYSF4

Section 3.3: Electron Configuration

Objectives:

1. Compare the Rutherford, Bohr, and quantum models of the atom.

2. Explain how the wavelengths of light emitted by an atom provide evidence for the

modern model of the atom and information about electron energy levels.

3. List the four quantum numbers, and describe their significance.

4. Write the electron configuration of an atom by using the Pauli Exclusion Principle

and the Aufbau Principle.

Vocabulary:

orbital

electromagnetic spectrum

ground state

excited state

quantum number

Pauli exclusion principle

electron configuration

aufbau principle

Hund’s rule

Atomic Models

Building a model helps scientists imagine what may be happening at the microscopic level.

Models have limitations and have to be modified or discarded as new information comes

available.

Dolton’s Model

2 + 2

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Thomson’s Model

Key hypothesis: atoms have neutral charge, so there must exist positively charged

subatomic particles that balance the negativity of the electrons. Thomson proposed his

“Plum Pudding Model” of the atom.

Rutherford’s Model

Rutherford’s gold foil experiment led replacement of the Plum Pudding model with the

nuclear model of the atom. Rutherford suggested that electrons revolved around the nucleus

like planets orbited around the sun. But this did not explain why the negative electrons did

not get pulled into the positive nucleus, because opposite charges attract.

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Bohr’s Model

Electrons can only exist in energy levels (also called orbitals) that are specific distances

from the nucleus. The electrons do not exist in between the energy levels. Energy levels

closer the nucleus are lower in energy. The energy levels further from the nucleus are

higher in energy.

Normally an electron hangs out in the lowest possible energy level, called the ground state

level. If an atom absorbs energy, the electron can jump to a higher energy level. As the

atom cools down, the excited electron falls into a lower orbital, loses energy, causing the

release of light energy as a photon, or light packet. The color of the released light is

determined by how far the electron falls and how much energy it gives up.

Energy Transition Calculations Using Bohr’s Model

1 3026 17 31

relative energies

a

b

c

d

ef

g

h

i

The location of an electron in an energy level is called its energy state. This diagram shows

a cross-section of the transitions that an electron can make from a higher (excited state) to a

lower (lower excited state or ground state) energy level. The lower energy levels are closer

to the nucleus.

The type of light given off when an electron drops from a higher to a lower level is

determined by the difference in the relative energies of the energy levels. For example, the

energy released when an electron falls from the level having a relative energy of 31 to the

ground state level (indicated by arrow d) is 31 – 1 = 30. This energy transition might

correspond to the color red.

Quantum Mechanics: The Structure of Atom (6:11 min)

http://www.youtube.com/watch?v=-YYBCNQnYNM

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Electromagnetic Spectrum

Line Spectrum

When a particular atom absorbs energy, its electrons are elevated from their ground state

into higher or excited states. As the atom cools, the excited electrons fall into lower energy

states, releasing light in the process. Each element has a unique set of energy levels, and

gives off a unique set of colors called a line spectrum. The line spectrum can be used to

identify the type of element, similar to using a fingerprint to identify a person. The

following diagram shows the emission line spectrum for hydrogen, mercury, and neon.

Also shown is the absorption line spectrum for hydrogen.

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Modern Model—Quantum View of Atomic Structure

The quantum model is the present-day view of the atom. The quantum model places

electrons in one of several types of orbitals, or regions of high probability for finding a

particular electron. It is not possible to know exactly where the electron is or the direction it

is going. The following diagram attempts to show a three dimensional view of the four

types of orbitals, the s, p, d and f. Note that the s orbital has one sublevel the p orbital has

three sublevels, the d orbital has five sublevels, and the f orbital has seven sublevels. Each

sublevel can hold a maximum of two electrons.

See: http://www.kentchemistry.com/links/AtomicStructure/PauliHundsRule.htm

Electron Placement

Electrons are placed into a particular orbital based on a set of quantum numbers that are

unique to each electron in the atom.

The principal quantum number (n) indicates the main energy level occupied by the

electron and divides the periodic table into rows. n can have values of 1, 2, 3, 4, 5, 6,

and 7. Higher n values are further from the nucleus and have higher levels of energy.

The principal quantum number is further divided into the l, m, and spin sublevels.

Angular momentum quantum number (l) indicates the shape of the orbital and can

have values of 1, 2, 3, …

Magnetic quantum number (m) is a subset of l and indicates the orientation and can

have values of …, -2, -1, 0, 1, 2, …

Spin quantum number indicates the spin (+½, -½ or ↑ and ↓) of the electrons

magnetic field and can hold a maximum of two electrons

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Electron Placement Rules

Aufbau Principal: electrons fill orbitals starting at the lowest available (possible)

energy states before filling higher states

Pauli Exclusion Principal: An orbital can hold 0, 1, or 2 electrons only, and if there

are two electrons in the orbital, they must have opposite (paired) spins.

incorrect; electrons must spin in opposite directions

correct; the unpaired electrons have the same spin

Hund’s Rule: If multiple orbitals of the same energy are available, unoccupied

orbitals will be filled before occupied orbitals are reused (by electrons having

different spins).

Electron Configuration

The arrangement of electrons in a particular atom can be found by

determining the electron configuration of the electrons. The

electron configuration can be determined filling the orbitals

according to the yellow brick road in the following schematic:

For example, the neutral sulfur atom has 16 electrons and thus has a

configuration of: S: 1s22s22p63s23p4

Each element’s configuration builds on the previous elements

configuration. To save space, use the configuration of the previous

noble gas. Ne has 10 electrons, so: S: [Ne] 3s23p4

Taking this process to the extreme, an atom of the element with

atomic number 118 has 118 electrons. By following the yellow brick road, the element

would have the following electron diagram.

1s22s22p63s23p64s23d104p65s24d105p66s24f145d106p67s25f146d107p6

Orbital Diagrams

The orbital diagram shows the energy levels and represents electrons by arrows placed

inside the sub orbitals represented by boxes. The orbital diagram of sulfur is shown below.

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The orbital diagram and the electron configuration process governed by the following

rules:

1. No two electrons can be in the same place (Pauli Exclusion Principle)

2. Electrons assume the lowest possible energy level (aufbau principle)

3. Each individual orbital holds a maximum of two oppositely spinning electrons

(Hund’s Rule)

4. Given a choice, electrons prefer to occupy all available orbitals before pairing up.

For 4 electrons in p orbitals, this means ↓↑ ↑ ↑ rather than ↓↑ ↓↑ __

Section 3.4: Counting Atoms

Objectives:

1. Compare the quantities and units for atomic mass with those for molar mass.

2. Define mole and explain why this unit is used to count atoms.

3. Calculate either mass with molar mass or number with Avogadro’s number given

an amount in moles.

Vocabulary:

atomic mass

mole

molar mass

Avogadro’s number

Atomic Mass

Average mass of one copper atom:

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1 atom Cu = 1.0551 x 10-23 g

Mass of one (pre-1982) penny is 3.13 g Cu. Therefore, the number of atoms in a

copper penny is:

(3.13 g Cu) (1 atom Cu

1.0551 x 10−23g Cu) = 2.97 x 1022

Number of stars in the universe = 1021 = 1000000000000000000000

There are more atoms of Cu in a penny than stars in the universe!

Using grams for the mass of atoms is inconvenient so chemists use the atomic mass unit

(amu).

One atomic mass unit = 1/12th the mass of a carbon-12 atom (a carbon atom with 6 protons

and 6 neutrons).

One amu = 1.6605402 x 10-24 g.

Introduction to the Mole

Samples of elements have great numbers of atoms. To make working with these numbers

easier, chemists use a unit called a mole (mol). A mole is defined as the number of atoms

in exactly 12 grams of carbon-12. This number equals 6.02 x 1023 and is called Avogadro’s

number. Like the dozen, the mole is used to count things.

The molar mass is the mass in grams of one mole (or 6.02 x 1023 atoms) of the element.

Molar mass has the units of grams per mole (g/mol). The molar mass is equal to the atomic

mass of an element.

The molar mass of an element is used to convert between moles and grams.

? 𝑚𝑜𝑙 (𝑎𝑚𝑜𝑢𝑛𝑡)𝑥 ? 𝑔

𝑚𝑜𝑙= ? 𝑔 (𝑚𝑎𝑠𝑠)

? 𝑔 (𝑚𝑎𝑠𝑠) 𝑥 1 𝑚𝑜𝑙

? 𝑔= ? 𝑚𝑜𝑙 (𝑎𝑚𝑜𝑢𝑛𝑡)

Avogadro’s number can be used to convert between moles and number of atoms.

? 𝑚𝑜𝑙 (𝑎𝑚𝑜𝑢𝑛𝑡)𝑥 6.02 𝑥 1023𝑎𝑡𝑜𝑚𝑠

1 𝑚𝑜𝑙 = ? 𝑎𝑡𝑜𝑚𝑠 (𝑛𝑢𝑚𝑏𝑒𝑟 𝑜𝑓 𝑎𝑡𝑜𝑚𝑠)

? 𝑎𝑡𝑜𝑚𝑠 (𝑛𝑢𝑚𝑏𝑒𝑟 𝑜𝑓 𝑎𝑡𝑜𝑚𝑠) 𝑥 1 𝑚𝑜𝑙

6.02 𝑥 1023 𝑎𝑡𝑜𝑚𝑠 = ? 𝑚𝑜𝑙 (𝑎𝑚𝑜𝑢𝑛𝑡)

YouTube

How big is a mole? (Not the animal, the other one.) (4:33 min)

http://www.youtube.com/watch?v=TEl4jeETVmg

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1-1 The Mole & Avogadro’s Number (4:19 min)

http://www.youtube.com/watch?v=g_BelGwRxG8

Mr. Mole Man (0:43 min)

http://www.youtube.com/watch?v=h17bd70-ous

A mole is a unit (2:27 min)

http://www.youtube.com/watch?v=U1frmqkNqW0

Happy Mole Day to You (2:17 min)

http://www.youtube.com/watch?v=ReMe348Im2w

Happy Mole Day (Party in the USA) (2:34 min)

http://www.youtube.com/watch?v=WjjmzM-YV1s

Michael Offutt-A Mole is a Unit-Music Video (3:18 min)

http://www.youtube.com/watch?v=Qg0Lajwew3A