AtomicTheory1

(1)

The Development of

Atomic Theory

(2)

The

Atom

• The term atom is derived from the Greek word (atomos)

meaning indivisible

• Democritius (470-370 BC )

suggested that all matter was made up of indivisible particles called

atoms

(3)

Law of Constant

Composition

A compound always contains atoms of two or More elements combined in definite proportions by mass

Example:

Water H

₂₂

O always contains 8

grams of oxygen to 1 gram of

hydrogen

(4)

Law of Multiple

Proportions

Atoms of two or more elements may

combine in different ratios to produce

more than one compound.

Examples:

NO NO

₂₂

N

₂₂

O N

₂₂

O

₅₅

(5)

Dalton’s Atomic

Theory

1. All elements are

composed of indivisible

and indestructible

particles called atoms. 2. Atoms of the same

element are exactly alike, They have the

same masses. 3. Atoms of different

elements have different masses.

4. Atoms combine to form compounds in small

(6)

SomeObjections to

Dalton’s Atomic

Theory

1. Atoms are not indivisible. They

are composed of subatomic

particles.

2. Not all atoms of a particular

element have exactly the same

mass.

3. Some nuclear transformations

alter (destroy) atoms

(7)

Crookes Experiment

Crookes found that passing an electrical current

through a gas at very low pressure caused the gas to

glow. Putting a magnet next to the beam caused it to

be deflected.

(8)

The Electron

1. The electron was the first subatomic particle to be identified.

2. In 1897 J.J Thomson used a cathode ray tube to establish the presence of a

charged particle known as the electron 3. Thomson established the charge to

mass ratio

E

/

m = 1.76 x 108 coulombs/gram

(9)

A Cathode Ray Tube

Thomson found that an electrical field would

also deflect an electron beam. He surmised

that the

(10)

Thomsen’s Charge to

Mass Ratio

Thomson proposed that the cathode rays

were in fact charged particles coming

from the traces gases in the Cathode Ray

Tube.

He then determined that the

ratio of

(11)

Thomsen’s Plum

Pudding Model

Thompson proposed that

an atom was made up of electrons scattered

unevenly through out an elastic sphere. These

charges were surrounded by a sea of positive

charge to balance the electron's charge like plums surrounded by pudding.

This early model of the atom was called The Plum

Pudding Model. A more contemporary American

(12)

Millikan’s Experiment

 By varying the charge on the plates, Millikan found _{By varying the charge on the plates, Millikan found}

that he could suspend the oil drops or make them

levitate.

(13)

Millikan’s Experiment

Millikan used his data to measure the charge of an electron and then to

calculate the mass of the electron from Thomson’s charge to mass ratio.

Given the charge =

1.60 x 10-19 coulomb and

the ratio of E/m = 1.76 x 108 coulombs/gram it is

possible to calculate the mass

Mass

= 9.11 x 10-28 gram

(14)

Protons

First observed by E. Goldstein in 1896

J.J. Thomson established the presence

of positive charges.

The mass of the proton is

1.673 x 10

-24

grams

(15)

Rutherford’s Experiment

1910

Ernest Rutherford

Rutherford oversaw Geiger and Marsden carrying out his famous experiment.

They fired high speed alpha particles (Helium nuclei) at a

piece of gold foil which was only a few atoms thick.

They found that although most of them passed through. About 1 in 10,000 hit and were deflected

(16)

Rutherford’s Experiment

(17)

Rutherford’s

Experiment

(18)

Rutherford’s

Experiment

By studying this

pattern, Rutherford

concluded that

atoms have a very

dense nucleus, but

there are mostly

empty space.

(19)

Subatomic Particles

The diameter of a single atom ranges From 0.1 to 0.5 nm. (1 nm = 10-9 m).

Within the atom are smaller particles: Electrons

Protons Neutrons

(20)

Neutrons

Discovered by James Chadwick in 1932

Slightly heavier than a proton

Mass of a neutron = 1.675 x 10

-24

grams

(21)

The Bohr Model

Niels Bohr

proposed the

Planetary Model in 1913.

Electrons move in

definite orbits around the

nucleus like planets

moving around the

nucleus. Bohr proposed

that each electron moves

in a specific energy level.

(22)

Aspects of the Bohr

Model

Bohr put together Balmer’s and

Plank’s discoveries to form a new atomic model

In Bohr’s model:

1. Electrons can orbit only at certain allowed distances from the

nucleus.

2. Electrons that are further away from the nucleus have higher energy levels (explaining the

(23)

The Electromagnetic Spectrum

(24)

Emission Spectra

(25)

Flame Tests

(26)

According to Bohr

Atoms radiate energy

whenever an electron jumps from a higher-energy orbit to a lower-energy orbit. Also, an

atom absorbs energy when an electron gets boosted from a low-energy orbit to a

high-energy orbit.

(27)

Problems with the Bohr

Model

 The Bohr model provided a model that gave _{The Bohr model provided a model that gave}

precise results for simple atoms like hydrogen.

 Using the Bohr model precise energies could _{Using the Bohr model precise energies could}

be calculated for energy level transitions in

hydrogen.

 Unfortunately these calculations did not work _{Unfortunately these calculations did not work}

for atoms with more than 1 electron.

(28)

Weakness of the Bohr

Model

• _{According to the Bohr model electrons could}

be found in orbitals with distinct energies.

• When the data for energies measured using spectral methods where compared to the

values predicted by the Rydberg equation, they were accurate only for hydrogen.

• _{By the 1920s, further experiments showed that}

Bohr's model of the atom had some difficulties. Bohr's atom seemed too simple to describe the heavier elements.

(29)

Modern View of the

Atom

The wave mechanical model for the

atom was developed to answer some of the objections that were raised about the Bohr model. It is based on the work of a number of scientists and evolved over a period of time

The quantum theorists such as Maxwell Planck suggested that energy

consists of small particles known as

photons. These photons can have only discreet energies

Maxwell Planck

(30)

Modern View of the

Atom

Albert Einstein

demonstrated the equivalence of matter and energy. Hence matter and energy in Einstein’s theory were not different entities but

different expressions of the same thing

Einstein then proposed the equivalence of Matter and Energy given by his famous equation

E = mc

2

(31)

Modern View of the

Atom

Louis de Broglie suggested that if energy could be

thought of as having particle properties, perhaps matter

could be thought of as having wave like characteristics

Louis de Broglie

(32)

Modern View of the

Atom

Louis de Broglie proposed that an electron is not just a particle but it also has wave characteristics.

mc

2

= h



(33)

Modern View of the

Atom

Heisenberg proposed that it was impossible to know the location and the momentum of a high

speed particle such as an electron.

The more precisely the position is determined, the less precisely the momentum is known in this instant, and vice versa.

--Heisenberg, Uncertainty paper, 1927

(34)

Modern View of the

Atom

The atom cannot be defined as a solar system with discreet orbits for the electrons. The best that we could do was define the

probability of finding an electron in a particular location.

The more precisely the position is determined, the less precisely the momentum is known in this

instant, and vice versa.

--Werner Heisenberg,

Uncertainty paper, 1927

(35)

Modern View of the

Atom

Edwin Schroedinger proposed that the electron is really a

wave. It only exists when we identify its location. Therefore the electrons are best thought of probability distributions rather than discreet particles.

(36)

Modern View of the Atom

The modern view of the atom suggests The modern view of the atom suggests that the atom is more like a cloud.

that the atom is more like a cloud.

Atomic orbitals around the nucleus Atomic orbitals around the nucleus define the places where electrons are

define the places where electrons are

most likely to be found.

(37)

Wave Mechanical Model

The location of the

electron in a hydrogen

atom is a probability

distribution.

(38)

Progression of Atomic Models

Our view of the atom has changed over time

(39)

ATOMIC STRUCTURE

Particle

proton

neutron

electron

Charge

+ charge

- charge

No charge

1

0 Mass

(40)

ATOMIC NUMBER AND MASS NUMBER

the number of protons in an atom

the number of protons and neutrons in an atom

He

2

4

Atomic Number Mass Number

Number of electrons = Number of protons

(41)

Atomic Mass

The atomic mass of an atom is a relative number that is used to compare the

mass of atoms.

An atomic mass unit is defined as 1/12 of the mass of an atom of carbon 12.

The atomic masses of all other atoms are a ratio to carbon 12

(42)

Isotopes

Many elements have atoms that have multiple forms

Different forms of the same element having

different numbers of neutrons are called isotopes.

For example: Carbon exists as both Carbon 12 and Carbon 14

Carbon 12 Carbon 14 6 electrons 6 electrons

6 protons 6 protons

(43)

Isotopes and Atomic

Mass

Many elements have atoms that have multiple isotopes.

Isotopes vary in abundance. Some are quite common while others are very rare. The atomic mass that appears in the

periodic table is a weighted average taking into account the relative abundance of each isotope.

(44)

Isotope:

one of two or more atoms having the one of two or more atoms having the same number of protons but different

same number of protons but different

numbers of neutrons

(45)

Measuring Atomic Mass

--the Mass Spectrometer

The mass

spectromete

r can be