Atomic Theory Click to begin.

Atomic Theory

Click to begin

Atomic TheoryNew to this

presentation? The Laws

Dalton’s

Atomic Theory

Early Models of the Atom

All About Isotopes

The Modern Atomic Theory

Click on a topic to begin your lesson

Main Menu

exit

Concept map of the atom

The Basic ButtonsThis presentation is intended to be interactive. It is best to click on these buttons only. Please do not just click on the mouse to move you along. You may wind up somewhere you do not want to be! Click a button below to proceed. Have fun learning!

Main Menu

Previous Slide

Next Slide

Menu or first slide of the current section

Returns you to this screen

Exits the presentation

Represents specific movesLast slide viewed

3

In the earliest days of chemistry the chief ‘chemical’ occupations were held by the alchemists. During the middle ages, alchemists tried to transform various metals into gold. They also produced metals from their ores, made glasses and enamels, and dyed fabrics. However, there was no understanding of what happens in these processes.

Dalton’s Atomic Theory


It wasn’t until 1803, however, when an English school teacher, John Dalton, was able to propose the very first atomic theory. This theory was the result of much experimentation into the nature of matter. It paved the way towards a deeper understanding of what chemicals are and what happens when they react.

Dalton’s Atomic Theory1. All elements are composed of atoms, which are indivisible and indestructible particles with no internal structure.

2. All atoms of the same element are exactly alike; in particular, they all have the same mass.

3. Atoms of different elements are different; in particular, they have different masses.

4. Compounds are formed by the joining of atoms of two or more elements. They are joined in a definite whole-number ratio, such as 1 to 1, 2 to 1, 3 to 2, etc.

5. Chemical reactions involve the rearrangement of atoms to make new compounds.


How was Dalton able to develop this atomic theory?

Proceed to the next section on The Laws…..

The Laws

John Dalton based his atomic theory upon the following laws.

Law of Conservation of Mass (1782)

Law of Definite Proportions (1797)

Law of Multiple Proportions (1803)

Law of Definite ProportionsJoseph Proust (1754 - 1826) found that the proportion by mass of the elements in pure samples of a given compound is always the same, regardless of the sample’s origin or size.

There are 50 grams of a chemical in this test tube. Analysis shows it to contain 26.36 g of chlorine and 23.64 g of copper.

What is the ratio of the mass of chlorine to the mass of copper in this compound?

Do this now before you click the answer!

AnswerExample 1

26.36 g Cl23.64 g Cu

= 1.12 g Cl1.00 g Cu

For every gram of copper in the compound, there are 1.12 g of chlorine.

Continued

Law of Definite ProportionsJoseph Proust (1754 - 1826) found that the proportion by mass of the elements in pure samples of a given compound is always the same, regardless of the sample’s origin or size.

There are 20 grams of a chemical in this test tube. Analysis shows it to contain 10.55 g of chlorine and 9.45 g of copper.

What is the ratio of the mass of chlorine to the mass of copper in this compound?

Do this now before you click the answer!

AnswerExample 2

10.55 g Cl 9.45 g Cu

= 1.12 g Cl1.00 g Cu

For every gram of copper in the compound, there are 1.12 g of chlorine.

Summary

Law of Definite Proportions

Both of these samples contain the same substance. Even though there are different quantities in each tube, the ratios or proportions of the elements to one another by mass are the same. Each contains:

1.12 g chlorine1.00 g copper

This is known asThe Law of Definite Proportions.

Summary

End of section.

Law of Multiple ProportionsJohn Dalton (1766 - 1844) proposed that the same elements that make up one compound could also make up another compound based upon the work of French chemist Berthollet in 1790. This law was proved by the Swedish chemist, Berzelius, after Dalton published his atomic theory!

ExampleThese two vials contain different compounds. They are, however, composed of the same elements.

Yellow = K2CrO4

Orange = K2Cr2O7

Law of Multiple ProportionsHow was John Dalton able to make such a statement? Let’s investigate! Click on the red arrow below to proceed.

Suppose the balloons shown below contain two different oxides of carbon, both of which are known to exist. Let’s find the mass ratio for each compound.

More

Law of Multiple Proportions

Analysis of this compound shows there are 16.0 g of oxygen and 12.0 g of carbon. Calculate the ratio of the mass of oxygen to the mass of carbon.

Answer 16 g O12 g C

= 1.33

Continue


Analysis of this compound shows there are 32.0 g of oxygen and 12.0 g of carbon. Calculate the ratio of the mass of oxygen to the mass of carbon.

Answer 32 g O12 g C

= 2.66

Continue


What do you notice about the two mass ratios below? Click the mouse button to see!

Mass ratio is 2.66

Continue

Mass ratio is 1.33

2.661.33

The mass ratio of one is double the other! What can this mean?

= 2.00

Law of Multiple ProportionsNotice that the mass of carbon in each sample is the same. This implies there are the same number of carbon atoms in each sample. Since the mass of the oxygen atoms doubles from one sample to the next, could this mean that the number of oxygen atoms also doubles? If we assume the elements in the first compound combine in a 1:1 ratio then the formula of the second compound can be predicted.

Continue

Mass of carbon Mass of oxygen Formula

1 2 g

12 g 32 g

16 g CO

CO2

Note that the ratio of the masses of oxygen in the two compounds is 2 : 1!


In summary, the Law of Multiple Proportions states that:

• The same elements that make up one compound can also make up another compound. Example: CO and CO2. • The mass ratio of the first compound compared to the mass ratio of the second compound will be different by a factor of a whole number.

• The masses of the element that combine with a fixed mass of the other element are themselves a whole number ratio.

Problem

A sample of sulfur dioxide with a mass of 10.00 g contains 5.00 g of sulfur. A sample of sulfur trioxide with a mass of 8.33 g contains 3.33 g of sulfur.

A. What is the mass of oxygen in the sample of sulfur dioxide?

B. What is the mass of oxygen in the sample of sulfur trioxide?

C. For a fixed mass of oxygen in each sample, what is the small whole-number ratio of the mass of sulfur in sulfur dioxide to the mass of sulfur in sulfur trioxide?

Answer

Answer

Answer





Answer

10.00 g total - 5.00 g Sulfur = 5.00 g OxygenAnswer

Answer





Answer

8.33 g total - 3.33 g Sulfur = 5.00 g Oxygen

Answer

Answer





Answer5.00 g S / 3.33 g S = 1.5 / 1 or 3 / 2Remember, 3 : 2 is the ratio because it represents a small whole number ratio. 1.5 : 1 is not a whole number ratio!

Answer

Answer

End of section.

Chemical reactions produce a variety of interesting ‘products’. Perhaps you have seen a clip o f a bomb exploding. Maybe you remember combining vinegar with baking soda. Wow! That produces lots of gas! Light can also be the product of a chemical reaction as shown below.

Antoine Lavoisier (1743 - 1794) wanted to know how the masses of the reactants and products of a chemical reaction compared. He carefully determined the masses of reactants and products in many different chemical reactions. Let’s investigate this ourselves!

The test tube contains a solution of potassium iodide.

The erlenmeyer flask contains a solution of lead (II) nitrate.

When these chemicals are mixed, a yellow precipitate (solid) forms in the liquid!

Let’s Investigate!

That’s pretty cool! Do you think the total mass of the products is greater than the total mass of the original reactants? We’re going to repeat this experiment but this time in a more quantitative fashion to find out.

Here is picture showing the initial mass of our ‘reactants’. Click on the picture to see the reading in a larger view. Please make a note of the mass.

When you are ready, click on the red arrow to watch the chemicals being mixed. On some computers the movie may not work. Just click the white button on the movie screen to continue.

Hmmm…Write down what you think happened to the mass of the system. Try and give a reason for your answer. Remember, two liquids did make a solid! Click a button below to continue.

Increase

Decrease

Stay the same

Well how about that!

The mass stayed the same! Did you guess correctly? Good for you if you did. Can you explain why the mass stayed the same? If you can’t explain why the mass is constant or if you guessed wrong, don’t worry. That’s why you are taking chemistry. This course is designed to help you understand what really goes on during a chemical reaction. In time, you will be able to come back to this demonstration and explain it fully!

We discovered that the total mass of the reactants is equal to the total mass of the products in a chemical reaction. Lavoisier discovered this too. This simple statement is now known as the: (click)

Law ofConservation

of Mass

Problem

Try the following problem to see if you understand this law.

Suppose you have 15 grams of substance A and 13 grams of substance B. How much substance C will be formed if you also produce 9 grams of substance D as shown below?

Reactants Products

A + B C + D15 g 13g ? g 9g

The Law of Conservation of Mass(Click until buttons appear)

The reactants side of the chemical equation shows a total of 28 grams of chemical being used. The products side of the equation shows 9 grams of D produced. Since the law states that the total mass of the reactants must equal the total mass of the products, the amount of C that must be produced is: (click)

Reactants Products

A + B C + D15 g 13g ? g 9g

28 g reactants- 9 g of D 19 g of C

Solution

End of section.

Significant discoveries occurred since Dalton’s time that seriously changed the way people envisioned the atom. A summary of the early models of the atom will appear when you click the mouse button.

Dalton’ Billiard Ball

Model

Thomson’s Plum Pudding

ModelRutherford’s

Planetary Model

Click on a model to learn more or…..

Early Models of Atoms

Dalton’s Billiard Ball model was the accepted model for about 100 years. It was a direct result from his atomic theory.

You may review Dalton’s Atomic Theory if you wish by clicking on the picture above.

Billiard Ball Model 1803

Sir J.J. Thomson (1856-1940) proposed that the structure of the atom could be compared to a plum or raisin pudding. The pudding would consist of a positive charged matrix. Embedded in it would be electrons that would neutralize the positive charge. Thus, the atom would be neutral overall. This was the first model that incorporated the experimental evidence that showed atoms must be composed of electrical charges. Click red dots below for more info before you leave this section!

Hey! Read more about me in your book!

Plum Pudding Model 1898

Click here to see a gas discharge

tube.

J.J. Thomson was experimenting with electricity using a gas discharge tube like that shown below. No air was present in the tube. Thomson discovered that an electric current flowed from the cathode to the anode by means of a ray of some kind. These rays later became known as electrons. Your instructor may give you more information about Thomson’s experiments.

Cathode Ray Tube

cathode anode

Glass tubeCathode rays

Discovery of the ElectronIn 1897, J.J. Thomson used a cathode ray tube to deduce the presence of a negatively charged particle.

Cathode ray tubes pass electricity through a gas that is contained at a very low pressure.

Conclusions from the Study of the Electron

Cathode rays have identical properties regardless of the element used to produce them. All elements must contain identically charged electrons.

Atoms are neutral, so there must be positive particles in the atom to balance the negative charge of the electrons

Electrons have so little mass that atoms must contain other particles that account for most of the mass

William Crookes was the first to discover these negative rays in 1879. Thomson was able to conduct experiments that proved these rays were actually negatively charged particles. Thomson is generally given credit for the discovery of the electron.

FYI (For Your Information)

Electrons are very small particles with very little mass and a fixed amount of negative charge. Thomson was able to calculate the ratio of the charge on an electron to its mass (e/m). In 1911, Robert Millikan actually measured the charge on the electron. With this value and Thomson’s e/m ratio, he was able to determine the mass of the electron.

The Electron

e/m= 1.759 x 108 coulombs/gram

m = e 1.759 x 108 coulombs/gram

m = 1.602 x 10-19 coulomb 1.759 x 108 coulombs/gram

m = 9.07 x 19-28 g

Ernest Rutherford’s (1871 - 1937) famous gold foil experiment soon made the plum pudding model obsolete. He and his researchers found that the atom must be composed of a very dense positively charged area they called a nucleus. The results of their experiment further indicated the atom is largely empty space. Rutherford theorized that the electrons must be located outside the nucleus at a relatively large distance. Perhaps the electrons circled about the nucleus like planets around the sun.

Click on me to see a diagram of the gold foil experiment.

Click on me to learn about this model and its

flaws.

Planetary Model 1909

Detector

Undeflected alpha particles

Deflected alpha particles

Gold atoms

Gold Foil Experiment

Rutherford’s Gold Foil ExperimentRutherford’s Gold Foil Experiment

Alpha () particles are helium nuclei Particles were fired at a thin sheet of gold foil Particle hits on the detecting screen (film) are recorded

Rutherford’s Findings

The nucleus is small The nucleus is dense The nucleus is positively charged

Most of the particles passed right through A few particles were deflected VERY FEW were greatly deflected

Conclusions:Conclusions:

NucleusElectron ( - charge)

Protons (+ charge)

Great distance

2. Electrons should attract to the nucleus and collapse the atom. Opposite charges attract!

1. Protons only accounted for half the mass of the atom.

The Model and It’s Flaws

Flaws

Click until buttons appear.

The proton was discovered shortly after the electron using specially designed gas discharge tubes. The proton has a positive charge equal in magnitude to the electron’s negative charge. However, it is about 1800 times heavier than the electron. Every element has a specific number of protons. Change the number of protons and you have a new element. Imagine the possibilities! Rutherford had theorized that there must be another particle that would have the same mass as the proton, no charge, and would also be located in the nucleus. In 1932, James Chadwick proved Rutherford’s theory correct by discovering the neutron.

The Proton

Summary of the proton, electron, and

neutron

(Postulated by William Wein in 1898. In 1920, Rutherford announced its existence.)

Protons, Electrons, Neutrons

Particle Mass Charge

Electron (0)9.11 x 10-28 g 1-Proton (1)1.67 x 10-24 g 1+Neutron (1)1.67 x 10-24 g 0

Find more information in your text!

Chadwick’s discovery of the neutron in 1932 helped take care of one of the flaws of the planetary model. Now, the entire mass of the atom could be accounted for. While the number of protons in an element are constant, it was discovered that atoms of the same element can have different numbers of neutrons. These atoms became known as isotopes of one another. Isotopes of hydrogen will appear below upon mouse clicks.

All About Isotopes

Protium Deuterium Tritium

Proton

Neutron

Not all isotopes have names like the isotopes of hydrogen. Usually an isotope is identified by it’s Mass Number. This represents the total number of protons and neutrons in the nucleus of the atom. For example, protium would also be known as hydrogen-1. What would deuterium and tritium be called? Click the mouse button to see if you are correct.

Naming Isotopes

Hydrogen-2 Hydrogen-3Hydrogen-1

Protium Deuterium Tritium

A simple isotope notation is often used to convey information about an isotope. Sample isotope notations are shown below.

Isotope Notation

1

H 1

Hydrogen-1

6

Li 3

Lithium-6

24

Mg 12

Magnesium-24

Which number in each notation represents the mass number?

Mass Number

24

Mg 12

The # of protons + the # of neutrons

What does the number 12 stand for?

Mass Number

Click here until buttons appear.

(Isotope notation for Mg-24)

Atomic Number (Z)(Henry Moseley in 1915)

24

Mg 12

The # of protons (p+) in the atom or….

Atomic Number

The # of electrons (e-) if the atom is electrically neutral ( i.e., not an ion. Click here to learn more about ions).

Click here until buttons appear.

Problem! 24

Mg 12

How many neutrons are there in this

isotope?

24 (neutrons + protons) - 12 protons = 12 neutrons

Fill in the blanks!Isotopenotation

Atomicnumber

Massnumber

Protons Neutrons Electrons Charge

184 74 0

82 128 +2

8 9 10

74

-2

80

74

17

110

82210

8 17 O -2

8

184 W 74

210 Pb +2

82

Click the mouse button repeatedly to view the order in whichthe boxes are solved. The navigation buttons will appear after the table is filled out.

56

Problem!Write the isotope notation for the

element Chlorine with atomic number 17 on the periodic

table.

Is There a Problem?(Click)

Yes! Chlorine has a number that is a decimal! How can this be? The number on the periodic table is an Atomic

Mass and not a Mass Number!

Mass Number vs. Atomic Mass

Mass number = # protons + # neutrons

Atomic mass = weighted average mass of all the isotopes of the element.

Atomic Mass is shown on the periodic table.

Atomic Mass

Also known as………

Weighted Average Atomic Mass

Average Atomic Mass

Average Atomic Mass

Calculated by…….

[(Mass of Isotope 1) x (% abundance of Isotope 1)]

+ [(Mass of Isotope 2) x (% abundance of Isotope 2)]

+ ………

= The average atomic mass as written on the periodic table

Average Atomic Mass…Problem

Calculate the average atomic mass of magnesium given the following information:

Isotope Relative Abundance Atomic Mass

Mg-24 78.70% 23.985 Mg-25 10.13% 24.986 Mg-26 11.17% 25.983

Average Atomic Mass…Problem

Mg-24 0.7870 x 23.985 = 18.876 Mg-25 0.1013 x 24.986 = 2.5310 Mg-26 + 0.1117 x 25.983 = 2.9023

The average atomic mass of magnesium is 24.309

For this isotope

11 p+

+ 11e-

0 net charge

23

Na 11

23

Na+

11

For this isotope

11 p+

+ 10e-

1+ net charge

Positive Ions (Atoms with a positive electric charge!)

Notice that only the electrons can change. Changing the protonswould change the atom!

Click until button appears

35

Cl 17

37

Cl-

17

Negative Ions (Atoms with a negative electric charge!)

Notice that only the electrons can change. Changing the protonswould change the atom!

For this isotope

17 p+

+ 17 e- 0 net charge

For this isotope

17 p+

+ 18 e- 1- net charge

Click until button appears

Modern Atomic Theory

Concept map of the atom

1. Atoms are not indivisible. They are made up of protons, electrons, and neutrons.

2. Atoms of the same element can, and do have different masses. These atoms are called isotopes. Isotopes have the same number of protons but a different number of neutrons.

3. Atoms of different elements are different. They differ in their physical and chemical properties.

4. Compounds are formed by the joining of atoms of two or more elements. They are joined in a definite whole-number ratio, such as 1 to 1, 2 to 1, 3 to 2, etc.

5. Chemical reactions involve the rearrangement of atoms to make new compounds.

Modern Atomic Theory All matter is composed of atoms Atoms cannot be subdivided, created, or destroyed in ordinary chemical reactions. However, these changes CAN occur in nuclear reactions!Atoms of an element have a characteristic average mass which is unique to that element.Atoms of any one element differ in properties from atoms of another element

Atom

Nucleus Electrons

NeutronsProtons

Positive charge

Equal to atomic number

Neutral charge

Equal to mass number

- protons

Negative charge

Equal to protons if not

an ion

contains

has have a

have a have a

69

Thanks for visiting!I certainly hope you learned something about

the history of Modern Atomic Theory!

EndShow

Atomic Theory Click to begin.

Documents

Transcript of Atomic Theory Click to begin.