THE ATOM

THE ATOMAP CHEMISTRY

1. Take out a pencil and sit in testing seatsAtomic Theory of MatterThe theory that atoms are the fundamental building blocks of matter reemerged in the early 19th century, championed by John Dalton.

2Figure 2.1 John Dalton (1766-1844)Daltons PostulatesEach element is composed of extremely small particles called atoms.

3Daltons PostulatesAll atoms of a given element are identical to one another in mass and other properties, but the atoms of one element are different from the atoms of all other elements.

4Daltons PostulatesAtoms of an element are not changed into atoms of a different element by chemical reactions; atoms are neither created nor destroyed in chemical reactions.

5Daltons PostulatesCompounds are formed when atoms of more than one element combine; a given compound always has the same relative number and kind of atoms.

6Law of Constant CompositionJoseph Proust (17541826)Also known as the law of definite proportions.The elemental composition of a pure substance never varies.7Multiple ProportionsThe law states that when chemical elements combine, they do so in a ratio of small whole numbers (Ex. Carbon and oxygen react to form carbon monoxide (CO) or carbon dioxide (CO2) but not CO13

If two elements form more than one compound between them, the ratio of the masses of the second element to a mass of the first elements will also be in small whole numbers8Law of Conservation of MassThe total mass of substances present at the end of a chemical process is the same as the mass of substances present before the process took place.9The Discovery of an ElectronEarly scientists knew about charges and in fact, Benjamin Franklin gave the names positive and negative to the two different charges.

But Michael Faraday (1791-1867) discovered electrons as cathode rays by applying a high voltage to the ends of a cathode ray tube.

The electrons emitted provide an image of their path when they strike a fluorescence zinc sulfide screen.

The Discovery of an ElectronJ.J. Thomson (1856-1940) used a specially designed cathode ray tube to apply both electric and magnetic fields simultaneously to the beam of cathode rays.

By balancing the effect of the electric field against that of the magnetic field he was able to calculate the charge to mass (e/m) ratio for the particles in the beam.

Thomsons experiments also demonstrated that electrons had a negative charge. In addition, he obtained the same mass to charge ratio with twenty different metals.

These results suggested that electrons are present in atoms of all elements.

The Atom, circa 1900:Plum pudding model, put forward by Thompson.Positive sphere of matter with negative electrons imbedded in it.

12Figure 2.9The Millikan Oil Drop ExperimentAmerican Physicist Robert Millikan (1868 - 1953) performed an experiment in which he sprayed oil droplets into a chamber from an atomizer. The oil droplets were allowed to settle slowly towards the bottom of the chamber.

Millikan had two charged plates on the top and bottom of the chamber which allowed him to control the rate at which the oil droplets fell.

By varying the voltage on the plates, he was able to just stop the oil droplets from falling.

As a result of his experiment, Millikan was able to calculate the following charges on the oil dropletsWhat can you conclude about the charges which Millikan obtained?They are all integer multiples of -1.6 x 10-19 C. Since the charge on an electron is the smallest possible charge, it is called an elementary charge.

RADIOACTIVITY The spontaneous emission of radiation by an atom.

First observed by Henri Becquerel.

Also studied by Marie and Pierre Curie.RADIOACTIVITY Three types of radiation were discovered by Ernest Rutherford: particles particles rays

Discovery of the NucleusErnest Rutherford shot particles at a thin sheet of gold foil and observed the pattern of scatter of the particles.

17Figure 2.10The Nuclear AtomSince some particles were deflected at large angles, Thompsons model could not be correct.

18Figure 2.11The Nuclear AtomRutherford postulated a very small, dense nucleus with the electrons around the outside of the atom.Most of the volume of the atom is empty space.

19Figure 2.12The Rutherford Scattering ExperimentIn 1910, Ernest Rutherford (1871 - 1937) and an associate, Hans Geiger (1882 - 1945) and a student, Ernest Marsden (1889 - 1970) performed an experiment in which they bombarded various metal foils with alpha particles. They observed that most of the alpha particles passed through the metal foil with little or no deflection. Some of the alpha particles actually bounced back towards the alpha source.

An Enlarged View of Rutherfords Experiment

CONCLUSIONS As a result of Rutherfords experiment, scientists were able to conclude two important features of the atom:

1. The atom is mostly empty space.2. The atoms positive charge is concentrated in a central core within the atom.Some Properties of the Gold Atom

What can you conclude about the relative size of the gold atom, nucleus, and an electron?

If an atom were the size of a football field (- 100 meters), the nucleus would be about 3 millimeters in diameter (the size of a BB) and an electron would be roughly half the size of the nucleus. That is, the atom is mostly empty space.

Some Properties of the Gold AtomIn other words, when Rutherford fired alpha particles at the gold foil, it would be like trying to hit a BB in the middle of a football field with a BB gun from many miles away, nearly impossible.

Rutherford also observed that only one in twenty thousand alpha particles were significantly deflected as they passed through the gold foil. Explain Rutherfords observation in terms of the relative size of the atom and the nucleus.

What comparison can you make about the mass of the nucleus compared to the mass of the atom?

Almost all of the mass of an atom is contained in the nucleus.ATOMIC STRUCTUREThe atom consists of two parts:The positively charged nucleus containing most the atoms mass.Negatively charged electrons found in the empty space around the nucleus.

Other Subatomic ParticlesProtons were discovered by Rutherford in 1919.Neutrons were discovered by James Chadwick in 1932.26CONCLUSION QUESTIONSGive It Some Thought Questions on pg. 41, 44

Go Figure Questions on pg. 42, 43THE ATOMAP CHEMISTRYDO NOW:1. If an atom has 15 protons, how many electrons does it have?2. Where do the protons reside in an atom?3. The diameter of a US dime is 17.9 mm, and the diameter of a silver atom is 2.88. How many silver atoms could be arranged side by side across the diameter of a dime?Equivalence statements: 1Ag atom = 2.881 = 1 x10 -10 m

SOME IMPORTANT ATOMIC STRUCTURE TERMSAtomic number - The atomic number equals the number of protons in the nucleus. The atomic number determines the element.

Nucleons - particles in the nucleus (protons plus neutrons)

Nuclear charge - The charge on the nucleus (the nuclear charge equals the number of protons)

Mass number - The mass number (an integer) is the sum of the neutrons plus protons. The mass number also equals the number of nucleons.

Isotopes - Isotopes are atoms of the same element (same atomic number) with a different number of neutrons.

Nuclide - A particular type of atom having a characteristic nucleus.

Ion - An ion is a charged particle.Subatomic ParticlesProtons and electrons are the only particles that have a charge.Protons and neutrons have essentially the same mass.The mass of an electron is so small we ignore it.

30Table 2.1IMPORTANT SUBATOMIC PARTICLESParticle Symbol ChargeMass Atomic Mass units, amu

How many?Protonp+ 11.007276Atomic Number Neutron n 01.008665Mass Number Atomic Number Electrone-10.0005486( the massof a proton)Number of protons Give the number of subatomic particles in each of the following atoms:This particle is called carbon - 14.

Protons: _______ Nuclear charge: ______

Neutrons: _______

Electrons: _______ Number of nucleons: ______

Protons: _______

Neutrons: _______

Electrons: _______ Mass number: ______

What is the name of this particle? ______________

Give the number of subatomic particles in each of the following atoms:A certain particle has 11 protons, 13 neutrons, and 10 electrons.1. What is the name of the element?2. What is the mass number of this particle?3. How many nucleons are found in this particle?4. What is the overall charge on this particle?5. Is this particle an atom or an ion?6. Give the symbol for the nucleus of this particle.Answer each of the following questions based on the following symbol:1. What is the mass number of this particle?

2. How many protons, neutrons, and electrons are found in this particle?

3. What is the nuclear charge on this element?

Answer each of the following questions based on the following symbol:1. What is the mass number of this particle?

2. How many protons, neutrons, and electrons are found in this particle?

3. What is the nuclear charge on this element?

Isotopes of Hydrogen

Give the number of protons, neutrons, and electrons for each of the isotopes of hydrogen listed above.

37Which of the following nuclides are isotopes?

Properties of Heavy Water and Ordinary Water PROPERTY ORDINARY WATER HEAVY WATER

Molecular Mass18 grams/mol 20 grams/mol Density at 25 DC0.997 g/cc 1.105 g/cc Temperature at max. Density4.0 C 11.6 C Melting Point0.00 C 3.802 C Boiling Point100.0 C 101.42 C Heat of Fusion1436 cal/mol 1510 cal/mol Heat of Vaporization10,484 cal/mol 10,743 cal/mol Dielectric Constant 81.5 80.7 Refractive Index1.333 1.328 Surface Tension72.75 dynes/cm 72.8 dynes/cm Viscosity at 25DC13.10 mpoise 16.85 mpoise Units of Atomic MassAn atomic mass unit (amu or u) is defined as exactly 1/12 the mass of a carbon-12 atom. or the mass of a carbon-12 atom is exactly 12.000 . . . amu.

The atomic masses of all the other elements are determined by comparing their masses with carbon-12.

Units of Atomic MassFor example, the mass of an average hydrogen atom is 8.400% the mass of a carbon-12 atom. Thus, the mass of an average hydrogen atom is 0.08400 x 12 amu = 1.008 amu.

Likewise, an average atom of magnesium has a mass which is 2.0254 times the mass of a carbon- 12 atom. Therefore, the mass of an average magnesium atom is 2.0254 x 12 amu = 24.305 amu.

COMPARISON OF ATOMIC MASS SCALES Based on Oxygen = exactly 16

Based on Carbon-12 = exactly 12

Oxygen-16

15.99560 amu

15.99491 amu

Oxygen

16 amu exactly

15.9999 amu

Carbon-12

12.00052 amu

12 amu exactly

Carbon

12.011 amu

12.010 amu

Silver

107.873 amu

107.868 amu

Determining the Mass of Atoms The mass of atoms is determined by use of a mass spectrograph. The sample of matter is ionized in a vacuum chamber. The resulting positive ions are then accelerated by means of a negatively charged screen. Most of the ions pass through the screen, though a slit to focus the ion beam, and then into a magnetic field. By varying the accelerating voltage, the ions can be made to strike the detector. Velocities up to 150,000 miles/sec can be obtained with a voltage of 400-4000 volts.

Atomic MassAtomic and molecular masses can be measured with great accuracy with a mass spectrometer.

44Figure 2.13Atomic MassMost accurate means for determining atomic weights

A gaseous sample is introduced and bombarded by a stream of high energy electrons

Collisions between the electrons and the atoms of molecules of the gas produce positively charged particles that are then accelerated toward a negatively charged grid .45Figure 2.13Atomic MassAfter particles pass through the grid, they encounter two slits that allow only a narrow beam of particles to pass

This beam then passes between the poles of a magnet, which deflects the particles into a curved path

Particles with the same charge, the extent of deflection depends on mass the more massive the particle, the less the deflection 46Figure 2.13Atomic MassParticles are separated by according to their masses

By hanging the strength of the magentic field or the accelerating voltage on the grid, charged particles of various masses can be selected to enter the detector

A graph of the intensity of the detector signal versus particle atomic mass is called a mass spectrum 47Figure 2.13Atomic MassAnalysis of a mass spectrum gives both the masses of the charged particles reaching the detector and their relative abundances, which are obtained from signal intensities

Knowing atomic mass and abundance of each isotope allows us to calculate the atomic weight of an element

48Figure 2.13Atomic MassUsed extensively today to identify chemical compounds and analyze mixtures of substancesAny molecules that loses electrons can fall apart, forming an array of positively charged fragmentsMass spec measures the masses of these fragments, producing a chemical fingerprint of the moelcule and providing clues about how the atoms were connected in the original molecule Application: used to determine the molecular structure of a newly synthesized compound or to identify a pollutant in the environment

49Figure 2.13

Atomic Mass of an Element The atomic mass of any element is the weighed average of its naturally occurring isotopes.

To calculate the atomic mass of the element, a weighed average, multiply the mass of each isotope by its percent abundance divided by 100, and find the sum of these values.

Atomic Mass of an Element

Atomic Mass Practice Find the atomic mass of boron if there are two naturally occurring isotopes, 10B and 11B. 19.6% of the atoms have a mass of 10.01292 amu and 80.4% of the atoms have a mass of 11.00931 amu.

Atomic Mass Practice A sample of element X contains 90. percent 35X atoms, 8.0 percent 37X atoms, and 2.0 percent 38X atoms. The average isotopic mass is closest to

Atomic Mass Practice Element X has two isotopes. If 72.0% of the element has an isotopic mass of 84.9 atomic mass units, and 28.0% of the element has an isotopic mass of 87.0 atomic mass units, the average atomic mass of element X is numerically equal to

Atomic Mass Practice If 75.0% of the isotopes of an element have a mass of 35.0 amu and 25.0% of the isotopes have a mass of 37.0 amu, what is the atomic mass of the element?

Atomic Mass Practice Given the following table containing the isotopic abundance of magnesium, calculate the atomic mass of an average atom of magnesium.

Atomic Mass Practice

Atomic Mass Practice There are two naturally occurring isotopes of copper, copper-63 and copper-65. Calculate the atomic mass for an average atom of copper if 69.17% of the atoms have a mass of 62.9396 amu and 30.83% of the atoms have a mass of 64.9278 amu.

Atomic Mass Practice A naturally occurring sample of an element contains 20.% of the atoms with a mass of 120 amu and 80.% of the atoms with a mass of 110 amu. Calculate the average atomic mass of the element.

Atomic Mass Practice A sample of a naturally element was found to contain 60.0% of its atoms have an isotopic mass of 70.0 amu and 40.0% of its atoms have an isotopic mass of 65.0 amu. Calculate the average atomic mass of the element.

Converting Atomic Mass Units to Grams The SI unit for the amount of substance is the mole (mol), which is the amount of substance that contains as many elementary particles (atoms, molecules, ions, etc.) as there are atoms in exactly 12 grams of the carbon-12 isotope. The term mole is the name of a number of items similar to a pair (2 items), a dozen (12 items), and a gross (144 items).

Converting Atomic Mass Units to Grams The currently accepted value for one mole is

This number is called Avogadros number

Converting Atomic Mass Units to Grams Likewise, the mass of 6.022 x 1023 atoms of the carbon-12 isotope is 12.00 grams.

Similarly, the mass of one Avogadros number of any particle in grams is equal to its mass in atomic mass units (amu) x 6.022 x 1023.

For example, the mass of 6.022 x 1023 atoms of an average oxygen atom is 15.99 grams.

CONCLUSION QUESTIONS/HW2.33, 2.34, 2.35, 2.36, 2.92, 2.93