chapter 3 Notes Atomic Theory

(1)

(2)

Aristotle

Aristotle did not believe in

atoms. He believed that all

matter was continuous,

meaning that matter could

not be divided into a

(3)

Early Atomic Theory

•

The Greek

philosopher

Democritus is

(4)

•

Democritus thought

that all matter was

made up of tiny

indivisible particles.

•

He called these

particles “atomos”

which in Greek

(5)

Important early discoveries in

chemistry

By the late 18

th

century and 19

th

century some significant

(6)

Antoine Lavoisier

Lavoisier is often called the

father of modern chemistry

.

He had many important

discoveries including the

elements oxygen and sulfur,

came up with the 1

st

extensive

(7)

Antoine Lavoisier

Maybe his most important

discovery was in 1774, the

law of conservation of

matter

– that matter cannot

be created or destroyed in a

chemical reaction!

Lavoisier was eventually

beheaded by the guillotine in

(8)

Joseph Proust

Joseph Proust, also a

Frenchman discovered

the law of definite

(9)

Law of definite proportions

– a chemical

compound will contain the same elements in

exactly the same proportions by mass

regardless of the size of the sample.

Example – NaCl is always 60.6% sodium

(10)

Law of multiple proportions

This law discovered by John Dalton in 1804

states that some elements will form more

than one type of compound, that the ratio of

the elements in a compound that is forms

can vary:

(11)

Dalton’s atomic theory

•

Then in 1803 an

English

schoolteacher

named John Dalton

proposed his own

atomic theory based

on his own

(12)

Dalton’s Atomic Theory

•

All matter is made up of tiny, indivisible

particles called atoms

•

Atoms of a given element are identical

in size, mass and properties

•

Atoms combine in whole number ratios

to form chemical compounds

•

In a chemical reaction, atoms cannot be

created or destroyed, they are just

(13)

Discovery of the Electron

In 1897, J.J. Thomson used a cathode ray tube to deduce the presence of a negatively charged particle.

(14)

(15)

Conclusions from the Study of

the Electron

 Cathode rays have identical properties regardless of

the element used to produce them. All elements must contain identically charged electrons.

 Atoms are neutral, so there must be positive

particles in the atom to balance the negative charge of the electrons

 Electrons have so little mass that atoms must

(16)

Thomson’s Atomic Model

(17)

Rutherford’s Gold Foil Experiment

 Alpha particles are helium nuclei

 Particles were fired at a thin sheet of gold foil  Particle hits on the detecting screen (film) are

(18)

(19)

Rutherford’s Findings

 Most of the particles passed right through  A few particles were deflected

 VERY FEW were greatly deflected

(20)

Conclusions of Rutherford

1. The nucleus of the atom is small,

densely packed with protons -

positively charged! The electrons

are outside the nucleus

(21)

Can you really touch atoms?

(22)

Modern Atomic Theory

 Atoms of an element have a characteristic

average mass which is unique to that element.

 Atoms cannot be subdivided, created, or

destroyed in ordinary chemical reactions.

However, these changes CAN occur in nuclear reactions!

• All matter is composed of atoms

 Atoms of any one element differ in

(23)

The atom

 Most of the mass of the

atom is in the nucleus (protons and neutrons)

 Electrons are found

outside of the nucleus (the electron cloud)

 Most of the volume of

the atom is empty space

Strong forces called

nuclear forces called hold protons and neutrons

(24)

How much empty space is in

an atom?

If you took a hydrogen atom, expanded the

proton to the size of a grapefruit,

(25)

Atomic Particles

Particle Charge Mass # Location

Electron -1 0 amu Electron cloud

Proton +1 1 amu Nucleus

Neutron 0 1 amu Nucleus

(26)

About Quarks…

Protons and neutrons are NOT fundamental particles. Protons are made of

two “up” quarks and one “down” quark.

Neutrons are made of one “up” quark and

two “down” quarks.

Quarks are held together by “gluons”

(27)

What are quarks made of?

The current theory is

(28)

Atomic Number

Atomic number of an element is the number of protons in the nucleus of each atom of that

element.

Example:

Element # of protons Atomic #

Carbon 6 6

Phosphorus 15 15

(29)

Protons and Electrons

• The number of

protons in an atom

will equal the

number of

(30)

Mass & Atomic Number

• The top number is

the

mass number

which is the sum of

protons and

neutrons in an atom

• The bottom number

is the atomic

(31)

• The atomic number

can also be found on

the periodic table

(32)

Finding the number of

neutrons

•

To find the number

of neutrons in an

atom, subtract the

number of protons

from the mass

(33)

Isotopes

Isotopes are atoms of the same element having

different masses due to varying numbers of neutrons.

Isotope Protons Electrons Neutrons Nucleus

Hydrogen–1 (protium)

1 1 0

Hydrogen-2 (deuterium)

1 1 1

Hydrogen-3 (tritium)

(34)

Finding subatomic particles

Fill in the table below using the information given and the periodic talbe.

Mass # = p+_{+ n}0

Nuclide

p

+

n

0

e

-

Mass

#

Oxygen - 10

- 33 42

- 31 15

8 8 18

18

Arsenic 75 33 75

(35)

Atomic mass

Atomic mass represents the

average mass

of all the

(36)

Atomic Masses

Isotope Symbol Composition of

the nucleus % in nature

Carbon-12 12C 6 protons

6 neutrons

98.89%

7 neutrons

1.11%

8 neutrons

<0.01%

Atomic mass is the average of all the naturally

(37)

How to determine atomic mass:

Convert the percentages to the decimal equivalent,

multiply by the mass associated with that

percentage then find the sum:

(38)

(39)

The Mole

1 dozen =

1 gross =

1 ream =

1 mole =

12

144

500

(40)

Avogadro’s Number

6.02 x 1023 is called “Avogadro’s Number” in

honor of the Italian chemist Amadeo Avogadro (1776-1855).

Amadeo Avogadro

(41)

Calculations with Avogadro’s

number.

Converting moles to atoms:

# of moles given

x

6.02 x 10

23

atoms/mole

Example:

Convert 3.5 moles of lithium to atoms

3.50 moles of lithium x 6.02 x 10

23

=

(42)

Converting atoms to moles

# of atoms given

÷

6.02 x 10

23

atoms/mole

Example:

Convert 8.50 x 10

23

atoms of C to

moles of C

8.50 x 10

23

atoms C ÷ 6.02 x 10

23

atoms/mole=

(43)

Molar Mass

The mass of one mole of an element is

equal to the atomic mass of that element

in grams.

Examples:

1 mole carbon = 12.011 grams

1 mole of lithium = 6.941 grams

(44)

Converting moles to grams

To convert moles to grams:

# of moles given x molar mass

To convert grams to moles:

(45)

Moles to grams

Convert 2.5 moles of

boron into grams:

(46)

Grams to moles

Convert 150 grams of chlorine to

moles

:

(47)

What is Avogadro’s favorite

type of music?