Atomic Theory &

Atomic Structure

Tro, Chapter 4 & 9

Sections 4.1 – 4.4, 4.8, 4.9; 9.2 – 9.9

Document BIG IDEAS about:

• Atomic structure– Electrons (mass, size, position)– Protons and neutrons (mass, position)– Isotopes

• Changes in (MODERN) thought – Dalton– Thomson, Rutherford and Bohr

• Quantum theory (CONTEMPORARY)

Early Atomic Theories

Democritis (400 BCE)

• First to propose idea of atom• Atom = “a” + “tomos” = cannot be

cut• Based solely on logic; not

supported by experiments

Alchemy(12-1500 CE)

• Modern word ‘chemistry’ came from Arabic ‘alkimiya’

• recognized importance of experimentation

• Responsible for developing lab equipment & procedures still used today

NOTE: Alchemy is a field, NOT a person…

Galileo(~1600 CE)

• Birth of modern science - combining logic, experimenting, publishing results

Lavosier & Priestly(1700’s)

• Quantitative analysis of chemicals

Law of Conservation of Mass:

Matter can neither be created nor destroyed

Proust(1700’s)

• Developed Law of Definite Proportions

Law of Definite Proportions:Different samples of the same compound always contain its

constituent elements in the same proportions by mass

Law of Definite Proportions

• Copper carbonate always contains – 5.3 parts copper– 4 parts oxygen– 1 part carbon

by mass

Dalton(1800’s)

• School teacher that proposed the first modern-day idea of atoms

Law of Multiple Proportions:If 2 elements combine to form more than one compound, the masses of one element that

combine with a fixed mass of the other element are in small whole # ratios

Law of Multiple Proportions

Dalton’s Atomic Theory - 1808• All matter is composed of atoms which

cannot be subdivided• Atoms of same element are identical

(size, mass, reactivity)• Atoms combine to form compounds in

simple, whole # ratios• Chemical reactions involve the

separation, combination, or rearrangement of atoms; it does not result in their creation or destruction

Modern Atomic Theories

General Principle #1 Electric Charges

Objects with an equal amount of positive and negative charge are said

to be electrically neutral

+ – positive negative

General Principle #2Forces between Charges

• Objects with like charge repel

• Objects with opposite charge attract

+ + ++

– + – +

Forces between Charges• Electrostatic force becomes

greater with more charge• Electrostatic force becomes

smaller the greater the distance between the charges

Thomson’s Atomic Model (1904)

Cathode Ray Experiments• Any metal worked

for anode• Negative electric

field repelled beam• Object placed in

path of glow blocked beam

J.J. Thomson’s Contribution

• Discovered the electron (1897)• Plum Pudding model• Determined the charge-to-mass

ratio of an electron using data from cathode ray tube experiments

Evidence & Conclusions• cathode rays consisted of subatomic

particles from atoms of anode• cathode rays are negatively charged

• must also be positive charge

• Millikan (oil drop experiment, 1909) calculated electron’s mass to be 9.11 x 10-31 kg

Modern View of Atomic Structure

Particle

SymbolRelativ

e Charge

Mass (kg)

proton p+ +1 1.6726 x 10-27

neutron

n0 0 1.67510 x 10-27

electron

e- -1 9.1096 x 10-31

+

0

nucleons

Modern View of Atomic Structure

Particle

Relative

ChargeMass (kg)

Relative mass (amu)

p+ +1 1.6726 x 10-27 ~1

n0 0 1.67510 x 10-

27 ~1

e- -1 9.1096 x 10-31 ~0

+

0

Rutherford’s Problems• How is nucleus held together?• Why don’t electrons collapse into

nucleus?• H atom has 1 proton & He atom has

2 protons, mass ratio should be 2:1; instead the ratio is 4:1

…there must be another particle

The Gold Foil Experiment: Hypothesis

• The α-particles will pass straight through the atoms

What is an () alpha particle?

It is a positively charged Helium nucleus

Rutherford’s Gold Foil Experiment

The Gold Foil Experiment: Outcome

What’s happening?

The Gold Foil Experiment: Conclusions

Atoms :• must be mostly

space• must have a very

small, dense area of + charge

• Protons have same charge as e-, but almost 2000x more mass!

The Neutron• Discovered by James Chadwick in

1932.• Neutron is electrically neutral &

has slightly greater mass than a proton

Mystery solved.

Atomic theory timeline

Updating Dalton’s Atomic Theory

3 major differences between modern atomic theory & Dalton’s atomic theory:

• Atoms are NOT indivisible – they are made up of protons, neutrons, and electrons

• Atoms of the same element are NOT exactly alike – they can have different masses (isotopes)

• Atoms CAN be changed from one element to another, but not by chemical reactions (nuclear reactions)

Atomic Structure & Isotopes

Atomic Mass Unit (amu)• defined as a more convenient unit

for reporting mass of small numbers of atoms

• 12C is used as the reference• 1 amu is defined as exactly 1/12 of

a 12C atom

Getting Information from the Periodic Table

6C

12.0111

Atomic # = # p+ in nucleus

Elemental symbol

Atomic mass (more on this later)

Isotopic Notation• Atomic number (Z) = # of p+ in the nucleus• Mass number (A) = sum of # p+ & n0 in

nucleus• For a neutral atom, # e- = # p+

H11 He4

2 C126 O16

8 Zn6330

Mass number (A)

Examples

Atomic number (Z)

Isotopes• All atoms in an element have the

same atomic number• However, 2 atoms of the same

element can have different mass numbers – called isotopes

• Isotopes have:– Same # of p+

– Different # of no

Some Common Isotopes

H

H

H

11

21

31

C

C

C

126

136

146

U

U

23592

23892

Relative Abundance

Mass Spectrometry• Technique used to determine

atomic mass

e-

Atom bombarded by stream of high energy electrons

e-

e- collides with atom, “bounces” off, but transfers some energy to it

e-

+Atom dissipates excess energy by expelling an electron

Mass Spectrometry, cont.• Ions are accelerated through a magnetic field• Amount of deflection depends on the ion’s mass• Highest mass deflected least

• Lowest mass deflected most

N

S

++ +

+

++++

Mass Spectrometry, cont.

Mass (amu)

Sample mass spec for chlorine

Relative abundance of each isotope can be determined from relative peak heights

35 37

Relative Abundance & Atomic Mass

• Relative isotopic abundance is then used to calculate atomic mass

• Atomic mass is the weighted average of the mixture of isotopes

Example

average atomic mass = (atomic mass 35Cl)(fraction 35Cl) + (atomic mass 37Cl)(fraction 37Cl)= (34.968 amu)(0.7577) + (36.965 amu)(0.2423)= 35.45 amu

Calculate the atomic mass of Cl given the relative abundances of its isotopes:35Cl – 75.77%37Cl – 24.23%

PracticeCopper, a metal known since ancient times, is used in electrical cables & pennies, among other things. The atomic masses of its 2 stable isotopes, 63Cu (69.09%) and 65Cu (30.91%) are 62.93 amu and 64.9278 amu, respectively. Calculate the average atomic mass of copper – the relative abundances of each ion is given in parentheses. Answer: 63.54 amu

The Bohr Model

Electromagnetic Spectrum

Light

c =

c = speed of light (3.0 x 108 m)= wavelength= frequency

Frequency vs. Wavelength

Light

• Energy as frequency • Energy as wavelength • Light behaves like a particle

(photon) as well as a wave

c =

Emission Spectrums• When electricity is run through a

sample of hydrogen gas, hydrogen atoms gain energy

• H atoms loose that energy by emitting photons

• Resulting spectrum is discontinuouscontinuous

discontinuous

What’s happening?

Bohr Model• Electrons move in

circular orbits around the nucleus

• Only certain energy levels are “permitted ” (this explains the discrete lines for the emission spectrum of hydrogen)

Schroedinger/Heisenburg• Experiments used mathematics

(probability) to predict behavior of electrons– Schroedinger equation

approximated the probability of finding a single electron for H within a region close to the nucleus

– Heisenburg [Uncertainty Principle] reinforces the idea that we just don’t know!

Math in Context: Blackbody Experiments