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0. Atomic History and Structure: . What comes to mind when you think of the term “atom”?. How do we know what we know about atoms? List any people you can think of. Thales of Miletus (600BC). Noticed what we call static electricity with amber Things would be attracted to it when rubbed

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Atomic history and structure

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Atomic History and Structure:


What comes to mind when you think of the term atom

What comes to mind when you think of the term “atom”?

How do we know what we know about atoms? List any people you can think of.


Thales of miletus 600bc

Thales of Miletus (600BC)

  • Noticed what we call static electricity with amber

    • Things would be attracted to it when rubbed

    • It was a “magical property”

      • The term electron comes from the Greek word for amber: “elektron”


Kanada 600 501bc

Kanada (~600-501BC)

  • Indian attributed with first proposing the idea of atoms (called “parmanu” or “anu”)

  • 5 elements

    • Earth

    • Fire

    • Water

    • Air

    • Ether

  • Atoms were indestructable and eternal


Empedocles 450bc

Empedocles (450BC)

  • 4 elements:

    • Earth

    • Wind

    • Fire

    • Water

      • Everything was different combinations of these

  • This idea didn’t really change until1661!


Leucippus 490 bc

  • Proposed the idea of atoms

    • That two things exist

      • Atoms

      • Empty space.

Leucippus (~490 BC)


Democritus 420bc

  • Student of Leucippus

  • Matter is made up of “eternal, indivisible, indestructible and infinitely small substances which cling together in different combinations to form the objects perceptible to us”

  • “Atomos”

Democritus (420BC)

From :http://www.historyworld.net/wrldhis/PlainTextHistories.asp?historyid=ac20#ixzz1UvX6le4i

100 Greek Drachma, 1967


Atomic history and structure

  • Originally opposed the idea of atoms, then

  • Added hot/cold or moist/dry to the four elements:

    • earth (cold and dry)

    • air (hot and moist)

    • fire (hot and dry)

    • water (cold and moist)

  • The differences in matter where a result of different balances of these atoms

    • Changing the balance could change matter

      • ex: what we know as copper changed to gold

  • Aristotle 384 BC – 322 BC


Benjamin franklin 1752

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Benjamin Franklin (1752)

  • Franklin believed object had 1 of 2 charges (+/-)

  • Opposites attract, like charges repel (Coulomb’s Law, which the Greeks knew a little about)

  • Kite experiment (among others):

    • Electric charges run from + to –

    • Lightening is electricity

  • Words he gave us:

    • battery, conductor, condenser, charge, discharge, uncharged, negative, minus, plus, electric shock, and electrician.


J l proust 1794

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J.L. Proust (1794*)

  • Law of constant composition:

  • A given compound always contains the same elements in the same proportion

    • In other words…a given compound always has the same composition, regardless of where it comes from.

      • Ex: H2O is always 89% oxygen and 11% H by mass

        *not published or recognized until 1811


Dalton s atomic theory 1800

Dalton’s Atomic Theory ~1800

  • John Dalton (1766-1844) proposed an atomic theory

  • While this theory was not completely correct, it revolutionized how chemists looked at matter and brought about chemistry as we know it today instead of alchemy


Dalton s atomic symbols

Dalton’s Atomic Symbols


Dalton s atomic theory

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Dalton’s Atomic Theory

  • Elements are made of very small indivisible particles called atoms.

  • All atoms of a given element are identical (all hydrogen atoms are identical).

  • The atoms of an element are different than the atoms of another element (hydrogen is different than helium).

  • Atoms of one element can combine with the atoms of another element to make compounds. A given compound should have the same relative numbers and types of atoms.

  • Atoms are indivisible in chemical processes…they are not created or destroyed just reorganized.


Problems with dalton s atomic theory

Problems with Dalton’s Atomic Theory?

1. matter is composed of indivisible particles

Atoms Can Be Divided, but only in a nuclear reaction

2. all atoms of a particular element are identical

Does Not Account for Isotopes (atoms of the same element but a different mass due to a different number of neutrons)!

3. different elements have different atoms

YES!

4. atoms combine in certain whole-number ratios

YES! Called the Law of Definite Proportions

5. In a chemical reaction, atoms are merely rearranged to form new compounds; they are not created, destroyed, or changed into atoms of any other elements.

Yes, except for nuclear reactions that can change atoms of one element to a different element


Michael faraday 1832

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Michael Faraday (1832)

  • atoms contain particles with an electric charge

  • structure of atoms related to electricity

    • The electron was the fundamental particle of electricity


Jj berzelius 1779 1848

JJ Berzelius (1779-1848)

  • Came up with how we write chemical formulas

    • Symbols for elements

    • Subscripts to indicate numbers of each element (he used superscripts, though!)

    • Considered one of the fathers of modern chemistry

      • Along with

        • John Dalton

        • Antoine Lavoisier

        • Robert Boyle


Up until the 1900 s

Up until the 1900’s….

  • Atomic structure was thought about, but not well known. It took a few more people to really put things together, and build off of each other’s knowledge to come up with what we know today.


Atomic history and structure

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  • Lord William Thomson Kelvin (1903)

    • Proposed the Plum Pudding Model, but didn’t name it

      • Electrons embedded in a positive, spherical cloud


Jj thomson 1904

JJ Thomson (1904)

  • Discovered electrons (1897)

    • cathode ray tube

    • Called electrons corpuscles

      • Name electron came from George JohnstoneStoney, who proposed the concept in 1874 and 1881, and the word came in 1891

  • Named the “Plum Pudding” model of the atom (1904)


Atomic history and structure

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Cathode Ray Tube


Atomic history and structure

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Cathode ray tube


Hantaro nagaoka 1904

HantaroNagaoka (1904)

  • Proposed the planetary (Saturnian) model of the atom

    • Positive, massive nucleus

    • Electrons bound to the nucleus via gravity in charged rings

  • Both were confirmed by Rutherford

  • He abandoned the model in 1908 due to errors that were not confirmed by new studies (charged rings)


Rutherford s gold foil experiment

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Rutherford’s Gold Foil Experiment

  • alpha (α) particles: positively charged particles directed at thin metal foil

  • most particles made it through → empty space

  • others were deflected back → since alpha particles are positive, they had to bounce off of something positive

    So…there is a dense positive charge (nucleus) that the electrons move around.

Gold Foil Animation


Rutherford s experiment led to the nuclear view of the atom 1909 published 1911

Rutherford’s experiment led to the nuclear view of the atom (1909/ published 1911)

(side note- it was actually Geiger- Marsden Experiment. Scientists Hans G. and undergraduate Ernest M. worked for Rutherford.)

“It was quite the most incredible event that has ever happened to me in my life. It was almost as incredible as if you fired a 15-inch shell at a piece of tissue paper and it came back and hit you. On consideration, I realized that this scattering backward must be the result of a single collision, and when I made calculations I saw that it was impossible to get anything of that order of magnitude unless you took a system in which the greater part of the mass of the atom was concentrated in a minute nucleus. It was then that I had the idea of an atom with a minute massive center, carrying a charge.[2]”

—Ernest Rutherford


Gold foil and the models of the atom

Gold Foil and the Models of the Atom


Jj thomson 1912

JJ Thomson (1912)

  • Determined isotopes of atoms exist (1912)

    • Used anode rays

    • Found Ne deflected in two different paths using what we now call mass spectroscopy


Niels bohr 1885 1962

Niels Bohr (1885-1962)

  • Bohr Model or the Solar System Model

    • Niels Bohr in 1913 introduced his model of the hydrogen atom.

    • Electrons circle the nucleus in orbits, which are also called energy levels.

    • An electron can “jump” from a lower energy level to a higher one upon absorbing energy, creating an excited state.

    • The concept of energy levels accounts for the emission of distinct wavelengths of electromagnetic radiation during flame tests.


Bohr s orbit model 1913

Bohr’s Orbit Model (1913)

Electrons occupy orbitals around the nucleus according to their energy..


Glenn seaborg 1912 1999

Glenn Seaborg(1912-1999 )

  • Discovered 8 new elements.

  • Only living person for whom an element was named.


Atomic history and structure

Which brings us to the modern day view of the atom….


Atomic structure

ATOMIC STRUCTURE

The atom is mostly

empty space

  • protons and neutrons in the nucleus.

  • the number of electrons is equal to the number of protons.

  • electrons in space around the nucleus.

  • extremely small.

    • One teaspoon of water has 3 times as many atoms as the Atlantic Ocean has teaspoons of water.


Atomic composition

ATOMIC COMPOSITION

  • Protons (p+)

    • positive (+) electrical charge

    • mass = 1.672623 x 10-24 g

    • relative mass = 1.007 atomic mass units (amu)

      • but we can round to 1

  • Electrons (e-)

    • negative (-) electrical charge

    • relative mass = 0.0005 amu

      • but we can round to 0

  • Neutrons (no)

    • no electrical charge

    • mass = 1.009 amu

      • but we can round to 1


Atomic history and structure

The following four slides are for additional information only; you will not be tested on the fundamental particles. However, they could appear as extra credit on a test or quiz.


Subatomic particles can also be further broken down into fundamental particles

He

Subatomic Particles can also be further broken down into Fundamental Particles

  • Quarks

    • component of protons & neutrons

    • 6 types

      • Up, down

      • Spin, charm

      • Top, bottom

  • 3 quarks = 1 proton or 1 neutron


Subatomic particles and quarks

Subatomic Particles and Quarks


What about electrons

What about electrons?

  • Electrons are electrons

    • They are not made from quarks

      • Which is why they weigh so much less than p+ or no

    • Classified as a lepton


Subatomic particles

Subatomic Particles

More information at http://www.lns.cornell.edu/~nbm/NBM_INTRO_TO_HEP1.htm


Atomic number z

Atomic number

Atom symbol

AVERAGE Atomic Mass

Atomic Number, Z

All atoms of the same element have the same number of protons in the nucleus, Z

13

Al

26.981


Atomic history and structure

Atoms are neutral because the numbers of protons and electrons are equal - the opposite charges cancel.

  • 11 electrons

  • 11negative charges

  • 11 protons

  • 11 positive charges


Atomic history and structure

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Ions

  • A charged atom because of a gain or loss of electrons.

    • If an atom is neutral, the # of p+= # of e-

    • If it has lost1 e-, the atom has a 1+ charge

    • If it has gained 1 e-, the atom has a 1- charge


Atomic history and structure

IONS

  • Taking awayelectrons from an atom gives a CATION with a positive charge

  • Addingelectrons to an atom gives an ANION with a negative charge.

  • Atoms may gain or lose more than 1 e-

  • To tell the difference between an atom and an ion, look to see if there is a charge in the superscript!

  • Examples: Na+ Ca+2 I- O-2

    Na Ca I O compared to


Predicting ion charges

PREDICTING ION CHARGES

In general

  • metals lose electrons ---> cations

  • nonmetals gain electrons ---> anions


Charges on common ions

-3

-2

-1

+1

+2

Charges on Common Ions

-/+4

+3

By losing or gaining e-, atom has same number of e-’s as nearest Group 8A atom.


Mass number a

Mass Number, A

  • C atom with 6 protons and 6 neutrons is the mass standard

    • = 12 atomic mass units

  • Mass Number (A)

    • =(# protons) + (# neutrons)

  • NOTon the periodic table…(that is the AVERAGE atomic mass on the table)

    • Ex: A boron atom can have A = 5 p + 5 n = 10 amu


Atomic math

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Atomic Math

On periodic table- but not all PTs look exactly like this set up, but they have the same information


Think back

Think Back…

  • John Dalton stipulatedthat all atoms of a particular element were identical

    • Their atomic numbers were the same, and also their #’s of neutrons were identical

  • In 1912, J.J. Thomson discovered that this was not accurate

    • In an experiment measuring the mass-to-charge ratios of positive ions in neon gas, he made a remarkable discovery:

      • 91% of the atoms had one mass

      • The remaining atoms were 10% heavier

      • All of the atoms had 10 protons, however some had more neutrons


Isotopes

Isotopes

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  • atoms with the same number of protons (Z) but a different number of neutrons

    • same element, different atomic mass number (A)


Atomic history and structure

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1H (hydrogen): A=1 Z=1

2H (Deuterium): A=2Z=1

3H (Tritium): A=3Z=1


Isotopes their uses

Isotopes & Their Uses

Bone scans with radioactive technetium-99.


Isotopes their uses1

Isotopes & Their Uses

The tritium content of ground water is used to discover the source of the water, for example, in municipal water or the source of the steam from a volcano.


Learning check

Learning Check

Which of the following represent isotopes of the same element? Which element?

234 X234 X235 X238 X

92939292


Learning check1

Learning Check

Which of the following represent isotopes of the same element? Which element? The red ones are isotopes of Uranium

234 X234 X235 X238 X

92939292


Atomic math1

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Atomic Math

  • Atomic number (Z)

    • the number of protons in the nucleus

    • gives the element’s identity

  • (Atomic) Mass Number (A)

    • sum of the protons and neutrons for a given isotope of an element

  • Atomic Mass (also called Atomic Weight)

    • Weighted average mass of the atoms (accounts for all the isotopes) is average atomic mass


Counting protons neutrons and electrons

Counting Protons, Neutrons, and Electrons

  • Protons: Atomic Number (from periodic table)

  • Neutrons: Mass Number minus the number of protons (mass number is protons and neutrons because the mass of electrons is negligible)

  • Electrons:

    • If it’s an atom, the protons and electrons must be the SAME so that it is has a net charge of zero (equal numbers of + and -)

    • If it does NOT have an equal number of electrons, it is not an atom, it is an ION. For each negative charge, add an extra electron. For each positive charge, subtract an electron (Don’t add a proton!!! That changes the element!)


Learning check counting

Learning Check – Counting

State the number of protons, neutrons, and electrons in each of these ions.

39 K+16O -241Ca +2

198 20

#p+ ___________________

#no ___________________

#e- ___________________


Learning check counting1

Learning Check – Counting

State the number of protons, neutrons, and electrons in each of these ions.

39 K+16O -241Ca +2

198 20

#p+19 820

#no 20821

#e-181018


Learning check counting2

Learning Check – Counting

Naturally occurring carbon consists of three isotopes, 12C, 13C, and 14C. State the number of protons, neutrons, and electrons in each of these carbon atoms.

12C 13C14C

6 6 6

#p+ _______ _______ _______

#no _______ _______ _______

#e- _______ _______ _______


Answers

Answers

12C 13C14C

6 6 6

#p+666

#no678

#e-6 66


Learning check2

Learning Check

An atom has 14 protons and 20 neutrons.

A.Its atomic number is

1) 142) 163) 34

B. Its mass number is

1) 142) 163) 34

C. The element is

1) Si2) Ca3) Se

D.Another isotope of this element is

1) 34X 2) 34X 3) 36X

16 14 14


Learning check3

Learning Check

An atom has 14 protons and 20 neutrons.

A.Its atomic number is

1) 142) 163) 34

B. Its mass number is

1) 142) 163) 34

C. The element is

1) Si2) Ca3) Se

D.Another isotope of this element is

1) 34X 2) 34X 3) 36X

16 14 14


Atomic symbols nuclide notation

Atomic Symbols: Nuclide Notation

  • Nuclide: atomic species determined by nuclear contents

  • Show the name of the element, a hyphen, and the mass number in hyphen notation

    sodium-23

  • Show the mass number and atomic number in nuclear symbol from

    mass number

    23 Na

    atomic number11


Atomic history and structure

Nuclide notation: p+, charge, and average atomic mass

Mass number

(protons + neutrons)

Cl

37

Atomic number

(number of protons)

17

number of neutrons

A-Z =20

As atoms have no charge, the number of electrons is the same as the number of protons. This atom has 17 electrons.


Atomic history and structure

Nuclide notation – ions

Mass number

23

Na+

11

Atomic number

number of neutrons=

1+ charge means 1 electron less than the number of protons. This atom has 10 electrons.


Atomic history and structure

Nuclide notation –ions

Mass number

(protons + neutrons)

16

O2–

Atomic number

(number of protons)

8

number of neutrons=

2– charge means 2 electrons more than the number of protons. This atom has 10 electrons.


Learning check4

Learning Check

Write the nuclear symbol form for the following atoms or ions:

A. 8 p+, 8 n, 8 e-___________

B.17p+, 20n, 17e-___________

C. 47p+, 60 n, 46 e-___________


Atomic history and structure

Solution

A. 8 p+, 8 n, 8 e-16O

8

B. 17p+, 20 n, 17e- 37Cl

17

C. 47p+, 60 n, 46 e-107Ag+

47


Atomic history and structure

Learning Check

1. Which of the following pairs are isotopes of the same

element?

2. In which of the following pairs do both atoms have

8 neutrons?

A. 15X 15X

8 7

B. 12X 14X

66

C. 15X 16X

7 8


Atomic history and structure

Solution

B. 12X 14X

66

Both nuclear symbols represent isotopes of carbon with six protons each, but one has 6 neutrons and the other has 8.

C. 15X 16X

7 8

An atom of nitrogen (7) and an atom of oxygen (8) each have 8 neutrons.


Isotopes and average atomic mass

Isotopes and Average Atomic Mass

  • We are used to calculating #’s of p+, no and e- using whole numbers; however on the Periodic Table we often see a decimal number  Why?

  • Atomic Mass (on the Periodic Table)

    • The average of the isotopic masses, weighted according to the naturally occurring abundances of the isotopes of the element

    • In a weighted average we must assign greater importance – give greater weight – to the quantity that occurs more frequently


Isotopes and atomic mass

Isotopes and Atomic Mass

  • The atomic mass for each element on the periodic table reflects the relative abundance of each isotope in nature.

  • The mass on the periodic table is NOT the atomic mass number (A)


Atomic history and structure

AMUs and Atomic Weight

  • Atomic mass unit (amu) is the unit for relative atomic masses of the elements

    • 1 amu =1/12 the mass of C-12 isotope.

    • 1 amu = 1.6605x10-24grams

  • Protons (p+)

    • mass = 1.672623 x 10-24 g

    • relative mass = 1.007 atomic mass units (amu) but we can round to 1*

  • Electrons (e-)

    • relative mass = 0.0005 amubut we can round to 0*

  • Neutrons (no)

    • mass = 1.009 amubut we can round to 1*

    • *most times, like now; when we get to nuclear chemistry, we will not be able to!


Comparative example your grades

Comparative Example – Your Grades

  • To calculate your overall average, we use a weighted average instead of a simple average since different tasks are worth more

  • For example:

    (30/100 x 80)

    + (30/100 x 75)

    + (10/100 x 70)

    + (30/100 x 70)

    = 74.5%


To calculate average atomic mass

To Calculate Average Atomic Mass

  • You add up (fractional abundance X mass) for each isotope to get the weighted average

    • Fractional abundance = natural abundance/100

  • Ex: If something has 3 isotopes:

    (fractional abundance)isotope 1 X (mass)isotope 1

    +(fractional abundance)isotope 2X (mass)isotope 2

    +(fractional abundance)isotope 3X (mass)isotope 3

    = average atomic mass


Example

Example

  • Naturally occurring copper exists with the following abundances:

  • 69.17% is Cu-63 w/ atomic mass 62.93 amu

  • 30.83% is Cu-65 w/ atomic mass 64.93 amu

    (.6917) x (62.93)

    +(.3083) x (64.93)

    = 63.55 amu


Learning check5

Learning Check:

3 Isotopes of Ar occur in nature

  • 0.337% as Ar-36, 35.97 amu

  • 0.063% Ar-38, 37.96 amu

  • 99.6% Ar-40, 39.96 amu

  • Calculate the Average Atomic Mass


Answer to learning check

Answer to Learning Check

(.00337) x (35.97)

+ (.00063) x (37.96)

+ (.996) x (39.96)

= 39.95amu


Atomic history and structure

  • In J.J. Thomson’s experiment, he found that the percent abundances of neon are as follows:

    • Neon – 20 = 90.51%

    • Neon – 21 = 0.27%

    • Neon – 22 = 9.22%

  • Calculate the average atomic mass of neon showing all of your work


If a mass is not specifically given for an isotope

If a mass is not specifically given for an isotope

  • Then make the assumption that the mass is the same as the atomic mass number

    • It isn’t exactly correct, but it will be close


Average atomic mass

11B

10B

AVERAGE ATOMIC MASS

  • Boron is 20% 10B and 80% 11B. That is, 11B is 80 percent abundant on earth.

  • For boron, atomic weight=

    = 0.20 (10 amu) + 0.80 (11 amu) = 10.8 amu


Calculating abundance

Calculating & Abundance

  • Chlorine has two isotopes: chlorine-35 (mass 34.97 amu) and chlorine-37 (mass 36.97 amu).

  • What is the percent abundance of these two isotopes if chlorine's atomic mass is 35.453?


Answer check part 1

Answer Check Part 1

  • if 2 isotopes, then the total is 100%. assume one is x% (x), the other is automatically 100-x%, (1-x)

  • x(34.97) + (1-x)(36.97) = 35.453


Answer check part 2

Answer Check Part 2

  • x(34.97) + (1-x)(36.97)=35.453

  • Solve for x

  • 34.97x+36.97-36.97x=35.453

  • -2x+36.97=35.453

  • -2x=-1.517

  • x=.7585

  • 1-x=.2415


Answer check part 3

Answer Check Part 3

  • Therefore Cl-35 has a % abundance of 75.85% and Cl-37 has a % abundance of 24.15%


Problem 1

Problem 1

  • The two naturally occurring isotopes of nitrogen are nitrogen-14, with an atomic mass of 14.003074 amu, and nitrogen-15, with an atomic mass of 15.000108 amu. What are the percent natural abundances of these isotopes?

  • The atomic mass of nitrogen is 14.00674amu


Answer check

Answer Check

  • The atomic mass of nitrogen is 14.00674amu

  • 14.00674 = p(14.003074) + (1 -p)(15.000108)14.00674 = 14.003074p + 15.000108 - 15.000108p-0.997034p = -0.993368

  • p = 0.9963 = 99.63% (N14)1 - p = 0.0037 = 0.37% (N15)


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