CH 2 Atomic Theory

7/27/2019 CH 2 Atomic Theory

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2The Atomic Nature

of Matter

Introduction: The Evolution of Atomic Theory

2.1 - Atomic Theory

2.2 - Atomic Architecture: Electrons and Nuclei

2.3 – Atomic Diversity: The Elements

2.4 – Charged atoms: Ions

2.5 – Energy of Atoms and Molecules



The Greek Concept of Atomos:The Indivisible Atom

• Around 440 BC, Leucippus originated the atom concept.• His pupil, Democritus (c460-371 BC) extended it

• There are five major points to their atomic idea.• All matter is composed of atoms, which are too small to be seen.

These atoms CANNOT be further split into smaller portions.• There is a void, which is empty space between atoms.

• Atoms are completely solid.• Atoms are homogeneous, with no internal structure.• Atoms can differ in size, shape, and weight



Aristotle (384-322 BC)

Spoke openly against the concept— the atom concept diminished

Only a few scholars gave it much thought.The Catholic Church accepted Aristotle's position

— equated atomistic ideas with Godlessness

It was not until 1660 that Pierre Gassendi succeededin separating the twonot until 1803 that John Dalton put the atom ona solid scientific basis.



John Dalton 1803-1807Modern Atomic Theory

1) All matter is composed of tiny particles calledatoms.

2) All atoms of a given element have identicalchemical properties that are characteristic of that element.

3) Atoms form chemical compounds bycombining in whole-number ratios.

4) Atoms can change how they are combined,but they are neither created nor destroyed inchemical reactions.



Fig 2-4

Atoms are neither created nor destroyed during physical or chemical processes.

2.1 Atomic TheoryConservation of Atoms & Mass



Fig 2-5

Mass is neither created nor destroyed during physical or chemical processes.

Conservation of Atoms & Mass

2.1 Atomic Theory

Courtesy Patrick Watson



Diffusion : the gradual mixing of atoms & molecules due to their continual motion.

2.1 Atomic TheoryAtoms & Molecules are Continually in Motion

Fig 2-7



2.1 Atomic TheoryAtoms & Molecules are Continually in Motion

Fig 2-9 Courtesy Patrick Watson


http://slidepdf.com/reader/full/ch-2-atomic-theory 9/56Fig 2-6

Atoms Combine to Make Molecules

2.1 Atomic Theory



2.1 Atomic TheoryDynamic Molecular Equilibrium

Fig 2-10

the condition in which a forward and reverse process occur atequal rates, so the system undergoes no net change.

When this example is at dynamic equilibrium, the number of molecules vaporizing equals the number of molecules

condensing at any given moment.



Laws of Mass Conservation &Definite Composition

Law of Mass conservation: The total mass of substancesdoes not change during a chemical reaction.

Law of Definite ( or constant ) composition: No matter

what its source, a particular chemicalcompound is composed of the same elementsin the same parts (fractions) by mass.



Mass Percent Composition of Na 2SO 4

Na 2SO 4 = 2 atoms of Sodium + 1 atom of Sulfur + 4 atoms of Oxygen

Elemental masses

2 x Na = 2 x 22.99 = 45.981 x S = 1 x 32.07 = 32.07

4 x O = 4 x 16.00 = 64.00142.05

Percent of each Element

% Na = Mass Na / Total mass x 100%% Na = (45.98 / 142.05) x 100% =32.37%

% S = Mass S / Total mass x 100%% S = (32.07 / 142.05) x 100% = 22.58%

% O = Mass O / Total mass x 100%

% O = (64.00 / 142.05) x 100% = 45.05%Check % Na + % S + % O = 100%

32.37% + 22.58% + 45.05% = 100.00%



Calculating the Mass of an Elementin a Compound Ammonium Nitrate

Ammonium Nitrate = NH 4 NO 3

How much Nitrogen is in 455 kg of Ammonium Nitrate?The Formula Mass of Cpd is:

4 x H = 4 x 1.008 = 4.032 g2 x N = 2 X 14.01 = 28.02 g3 x O = 3 x 16.00 = 48.00 g

80.052 gTherefore gm Nitrogen/ gm Cpd

28.02 g Nitrogen

80.052 g Cpd= 0.35002249 g N / g Cpd x 100 = 35.00%

455 kg x 1000g / kg = 455,000 g NH 4 NO 3

455,000 g Cpd x 35.00 g N / 100g Cpd = 1.59 x 10 5 g Nitrogen

28.02 kg Nitrogen

80.052 kg NH 4 NO 4

= 159 kg Nitrogen455 kg NH4 NO

3Xor:



Law of Multiple ProportionsIf elements A and B react to form two compounds,

the different masses of B that combine with a fixedmass of A can be expressed as a ratio of small wholenumbers:

Example: Nitrogen Oxides I & II

Nitrogen Oxide I : 46.68% Nitrogen and 53.32% Oxygen Nitrogen Oxide II : 30.45% Nitrogen and 69.55% Oxygen

in 100 g of each Compound: g O = 53.32 g & 69.55 gg N = 46.68 g & 30.45 g

g O /g N = 1.142 & 2.284Cmpd II 2.284 2

Cmpd I 1.142 1=

Cmpd I Cmpd II



2.2 Atomic Architecture: Electrons & NucleiForces

Fig 2-12

Gravitational force : the force which pulls object toward thecenter of the Earth.

Electrical force : theattraction or repulsionbetween two chargedobjects.



2.2 Atomic Architecture: Electrons & NucleiForces

Fig 2-12

Gravitational force : the force which pulls object toward thecenter of the Earth.

Electrical force : the attraction or repulsion between twocharged objects.

Magnetic force : the force generated by charged objects inmotion.

Courtesy Patrick Watson



Davy, Faraday1807 Humphry Davy:

• Forces holding compounds together areelectrical

1833 Michael Faraday:— Atomic mass and the electricity

needed to free elements fromcompounds are related



More Important developments

• 1864 Heinrich Geissler: developed a pump• produced good vacuums in sealed glass tubes

• 1870’s William Crookes: Cathode rays • are negatively charged• same regardless of cathode metal• are particles with mass

• 1891 George Stoney• electricity exists in units called "electrons‖



Fig 2-13 Gas Discharge Tube

Conclusions:

• Atoms are made up of smaller positive and negative fragments.

•The negatively charged particles are electrons, which areuniform in behavior, regardless of their source.

2.2 Atomic Architecture: Electrons & NucleiElectrons





1897 J.J. Thomson

DISCOVERED THE ELECTRON



1897 J.J. Thomson

• Deflected an electron beam by both magneticand electric attraction/repulsion

• Measured the electron’s mass/charge ratio



Fig 2-14 Cathode Ray Tube

Conclusion:

Charge

Mass

e

m 1.76x10 11 C/kg

2.2 Atomic Architecture: Electrons & NucleiElectrons



Cathode Rays

• Attracted to the positive electrode• Not visible but could make things ―glow‖

• Traveled in a straight line• Could be bent by electric or magneticfields

• A plate in it’s path acquired a negativecharge• Same regardless of material



The quest continues• 1905 Albert Einstein: Photoelectric Effect

• light causes electrons to be emittedfrom metals

• quantized energy transfer causes the

emission—E=mc 2



He was SOOOOO excited!!





Fig 2- 15 Millikan’s Oil Drop Experiment

Conclusions:

•Electrons are particles.

•Electrons have a massof 9.1 x 10 -19 kg

Electrons

2.2 Atomic Architecture: Electrons & Nuclei

Charge = n (-1.6x10 -19 C)

ChargeMass

em

1.76x10 11 C/kg



Nucleus Discovered

1911 Ernest Rutherford:

small dense positive nucleus

nucleus is most of the mass of an atomelectrons are in the space around the nucleus





Ernest Rutherford (1871-1937)

• Won the Nobel Prize in Chemistryin 1908

• “It was quite the most incredibleevent..... It was almost as if a gunner were to fire a shell at a piece of tissueand the shell bounced right back!!!!! ”

2 2 At i A hit t El t & N l i



Fig 2-17 Schematic drawing of an atom

• Most of the mass & all of thepositive charge of the atomare in the nucleus , whichoccupies only 1 part in 10 14 of

the atoms volume.

• Electrons occupy a hugevolume in comparison to thenucleus, but have relativelysmall masses.


The Nucleus

2 2 At i A hit t El t & N l i



Fig 2-17 Schematic drawing of an atom

• Protons account for thepositive charge of the nucleusand have a positive chargewith a magnitude equal to thenegative charge of an

electron. Protons have amass about 2000 times thatof an electron.


The Nucleus

• The mass of a neutrons isabout the same as the massof a proton, but a neutron iselectrically neutral.



2.2 Atomic Architecture: Electrons & NucleiThe Nucleus

Name Symbol Charge Mass

Electron e -1.6022 x10 -19 C 9.1091 x10 -31 kgProton p +1.6022 x10 -19 C 1.6726 x10 -27 kgNeutron n 0 1.6749 x10 -27 kg

Table 2-1 Atomic Building Blocks



2.3 Atomic Diversity

Atomic number , Z: nulcear charge, number of protons

An element is identified by the charge of its nucleus

Atomic Definitions I: Symbols



Atomic Definitions I: Symbols,Isotopes,Numbers

XA

Z

X = Atomic symbol of the element, or element symbol

A = The Mass number; A = Z + N

Z = The Atomic Number, the Number of Protons in the Nucleus

N = The Number of Neutrons in the Nucleus

Isotopes = atoms of an element with the same number of protons,

but different numbers of Neutrons in the Nucleus

The Nuclear Symbol of the Atom, or Isotope



Example : How many protons, neutrons and electrons

do each of the following have?O C C


Atom with the same number of protons, butdifferent number of neutrons.

168

126

146

Isotopes

Z A

X Element symbolMass number

Atomic number



Examples:

O

C

C

16

8

126

146

6 protons, 6 neutrons, 6 electrons



X A

Z


Isotopes



46 Ti = 8.2%

47 Ti = 7.4%

48 Ti = 73.8%

49 Ti = 5.4%

50 Ti = 5.2%

2.3 Atomic DiversityIsotopes

Courtesy of Sachtleben Chemie GmbH (Paint)Courtesy M. Freeman/PhotoLink/Phoyo Disc (Knee Joint)




Cl

Fig 2-18 Natural abundance of the isotopes of Cl, Cr, Ge, Sn

Cr Ge Sn

Isotopes



Problem: Calculate the abundance of the two Bromine isotopes:



Problem: Calculate the abundance of the two Bromine isotopes:79Br = 78.918336 g/mol and 81Br = 80.91629 g/mol , given thatthe average mass of Bromine is 79.904 g/mol.

Plan: Let the abundance of 79

Br = X and of 81

Br = Y and X + Y = 1.0Solution: X(78.918336) + Y(80.91629) = 79.904

X + Y = 1.00 therefore X = 1.00 - Y

(1.00 - Y)(78.918336) + Y(80.91629) = 79.904

78.918336 - 78.918336 Y + 80.91629 Y = 79.904

1.997954 Y = 0.985664 or Y = 0.4933

X = 1.00 - Y = 1.00 - 0.4933 = 0.5067

%X = % 79Br = 0.5067 x 100% = 50.67% = 79Br %Y = % 81Br = 0.4933 x 100% = 49.33% = 81Br



Pg 55 2.3.2 – 2.3.3

M d R t f th At i Th



Modern Reassessment of the Atomic Theory

1. All matter is composed of atoms. Although atoms are composedof smaller particles (electrons, protons, and neutrons), the atomis the smallest body that retains the unique identity of the element.

2. Atoms of one element cannot be converted into atoms of another element in a chemical reaction. Elements can only be converted intoother elements in Nuclear reactions in which protons are changed.

3. All atoms of an element have the same number of protons andelectrons, which determines the chemical behavior of the element.

Isotopes of an element differ in the number of neutrons, and thusin mass number, but a sample of the element is treated as thoughits atoms have an average mass.

4. Compounds are formed by the chemical combination of two or moreelements in specific ratios, as originally stated by Dalton.



2.4 Charged Atoms: IonsFormation of Cations

Ions : electrically charged atomic or molecular particlesCations : ions with positive charges



2.4 Charged Atoms: IonsFormation of Anions

Anions : ions with negative charges

Net electrical charge is always conserved





2.4 Charged Atoms: Ions

Fig 2-21

Ionic Compounds

2 Na (s) + Cl 2(g) 2 NaCl (s)

A solid containing cations& anions in a balancedwhole-number ratio.

Courtesy Michael Watson



Fig. 2.19




Fig 2-23

Ionic SolutionsPure water Sugar + water

Salt + water

Courtesy Ken Karp




Fig 2-24

Ionic Solutions

+-

2 5 E f A d M l l



Fig 2-24

2.5 Energy of Atoms and Molecules




Kinetic Energy, E kinetic : the energy of directed motionof an object.

E kinetic = 1 / 2 mu 2 (2-1)

1 kg m 2s -2 = 1 J

Thermal Energy : the energy of random motion,translational, rotational and vibrational. The thermalenergy of an object is equal to the sum of the kineticenergy of its atoms.




Potential Energy: energy that is storedChemical Energy: potential energy stored aschemical bonds

Electrical Energy: potential energy that is the result of electrical forces between charged objects

Radiant Energy : the energy of electromagnetic radiant(light, photons)

E electric kq1q 2

r (2-2)



Energy is neither created nor destroyed in any process only transferred from one body to another,

or changed from one form to another.

Conservation of Energy




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CH 2 Atomic Theory