THE ATOM Objectives: Understand the

experimental design and conclusions

used in the development of modern

atomic theory, including Dalton’s

Postulates, Thomson’s discovery of

electron properties, Rutherford’s

nuclear atom, and Bohr’s nuclear

atom.

Democritus

• Made his discovery

around the year 250

B. C.

• This was the first

discovery about the

atom, the next would

come in another 2000

years.

The First Atom

• Democritus took a sea shell and broke it in

half.

• Than he broke it in half again.

• When the pieces got to small he use a

mortar and pestle to crush the shell.

• He finally believed he got to the smallest

piece possible and called it the ATOM;

which in Greek means INDIVISIBLE.

John Dalton (1766-1844)

A New System of Chemical

Philosophy (1808)

Dalton’s Atom Model

1. All matter is made on atoms; and atoms are indivisible.

2. Atoms of the same element are all identical.

3. Compounds are formed by a combination of two or more different atoms and they always have the same proportion of elements. THE LAW OF DEFINITE COMPOSITION

4. A chemical reaction is a rearrangement of atoms and the atoms are neither created nor destroyed. THE LAW OF CONSERVATION OF MATTER

J. J. Thomson (1856-1940)

Joseph John Thomson

• English physicist who in 1897 discovered a particle smaller than the atom ; the electron.

• Particle has a negative charge and is much smaller than the atom so must come from the inside of the atom.

• Electrons are scattered around the atom like raisins in pudding. (THE PLUM PUDDING MODEL)

Thomson and Rutherford

Rutherford’s Gold

Foil Experiment

Rutherford’s Gold

Foil Experiment

Ernest Rutherford (1871-1937)

• New Zealand born physicist; worked in England

• 1911 conducted the “Gold Foil Experiment” the proved the existence of a small positively charged center of the atom.

• Disproved the “Plum Pudding Model”

• THE NUCLEAR MODEL

• Discovered the proton.

• Thought that the electrons orbited the nucleus like planets orbited the sun.

Millikan’s Oil Drop Experiment

• A fine mist of oil droplets is introduced into the chamber.

• The oil is ionized by x-rays.

• The electrons adhere to the oil drops.

• The value for the charge of the electron can be calculated.

Niels Bohr (1885-1962)

• Danish physicist,

produced his model in

1911.

• Saw problems with

Rutherford’s model.

• If electrons “orbit” than

they are changing

direction so they are

accelerating.

• That would require

energy.

The Orbital Model

• Electrons do not

“orbit” but are in

allowable ENERGY

LEVELS.

• When the electrons

stay in these levels,

which are at specific

distances from the

nucleus, they do not

give off energy.

Bright Line Spectrum

• But, if the electron moves from one level to another it gives off or absorbs energy.

• These Bright Line Spectrums are produced when the electrons “fall back” to a lower energy level and give off energy.

• Every element has a unique Bright Line Spectrum.

The Subatomic Particles

THE PROTON

• p+

• positively charged

• located in the nucleus

• relative mass = 1 atomic mass unit

• mass = 1.673 x 10-24 grams

• equal to atomic number

• number of protons “defines” the atom

The Subatomic Particles

THE NEUTRON

• n0

• neutral (no electrical) charge

• located in the nucleus

• relative mass = 1 atomic mass unit

• mass = 1.675 x 10-24 grams

• equal to mass number minus atomic number

• mass number is protons + neutrons

• James Chadwick proposed the existence of the

neutron.

The Subatomic Particles

THE NEUTRON • Isotopes – different atoms of the same element that

have the same number of protons but different numbers of neutrons

• some isotopes are radioactive – they emit energy when the nucleus of the atom breaks down spontaneously

• most radioactive isotopes are not dangerous

• to determine if an isotope is radioactive calculate the proton to neutron ratio

• if ratio is greater than or less than 1:1 for “small” atoms the isotope is unstable (smaller than Ca)

• if ratio is greater then 1:1.5 for “large” atoms the isotope is unstable

The Subatomic Particles

THE ELECTRON • e- (negative electrical charge)

• located in the electron cloud which is divided into energy levels, sublevels, orbitals, and spins

• relative mass = 0 atomic mass units

• mass = 9.11 x 10-28 grams

• equal to the number of protons if atom is neutral

• atom becomes a charged ion if electrons are gained or lost

• positive ion = CATION

• formed by the loss of electron, happens to metals

• negative ion = ANION

• formed by the gain of electron, happens to nonmetals

Location of Electrons

• Energy Levels

• Discovered by Niels Bohr

• # electrons = 2n2

• “n” is the energy level

• 1st level can hold 2 e-

• 2nd level can hold 8 e-

• 3rd level can hold 18 e-

• (eight if the outside energy level)

• 4th level can hold 32 e-

• (eight if the outside)

• The outside level is called the valance level and can never hold more than 8 electrons.

NUCLEAR SYMBOLS

mass number ion charge

23 +1

Na p+ = 11

11 n0 = 12

atomic number e- = 10

name symbol atomic

number

mass

number

ion

charge

number

of

protons

number

of

neutrons

number

of

electrons

atomic

mass

calcium 20 42 +2 40.08

19 -1

F

9

10 10 10 20.18

238 0

U

92

name symbol atomic

number

mass

number

ion

charge

number

of

protons

number

of

neutrons

number

of

electrons

atomic

mass

potassium 19 40 +1 39.098

amu

15 -2

O

8

18 22 18 39.948

amu

56 0

Fe

26

Objective

• Use isotopic composition to

calculate the average atomic

mass of an element.

Mass Number vs. Atomic Mass

• mass number is given for an individual atom

• mass number is given in nuclear symbols

• atomic mass is an average mass for all isotopes for the element

• atomic mass is the number on the periodic table

• if you round the average atomic mass you will have the mass number of the most common isotope

Average Atomic Mass