Section 4-1

Section 4.1 Early Ideas About Matter

• Compare and contrast the atomic models of Democritus, Aristotle, and Dalton.

theory: an explanation supported by many experiments; is still subject to new experimental data, can be modified, and is considered successful if it can be used to make predictions that are true

• Understand how Dalton's theory explains the conservation of mass.

Section 4-1

Section 4.1 Early Ideas About Matter (cont.)

Dalton's atomic theory

The ancient Greeks tried to explain matter, but the scientific study of the atom began with John Dalton in the early 1800's.

How do we know atoms exist?

• How do we know atoms even exist?

– Atoms are too small to see

– Indirect Evidence

– For example, does wind exist?

Section 4-1

Greek Philosophers (cont.)

• Many ancient scholars believed matter was composed of such things as earth, water, air, and fire.

• Many believed matter could be endlessly divided into smaller and smaller pieces.

Section 4-1

Greek Philosophers (cont.)

• Democritus (460–370 B.C.) was the first person to propose the idea that matter was not infinitely divisible, but made up of individual particles called atomos.

• Aristotle (484–322 B.C.) disagreed with Democritus because he did not believe empty space could exist.

• Aristotle’s views went unchallenged for 2,000 years until science developed methods to test the validity of his ideas.

Section 4-1

Moving past the Greek Philosophers

• John Dalton revived the idea of the atom in the early 1800s based on numerous chemical reactions.

• Dalton’s _____________________easily explained conservation of mass in a reaction as the result of the combination, separation, or rearrangement of atoms.

• The invention of the chemical balance allowed for Dalton’s theories.

Section 4-1

1. Matter is composed of extremely small particles called atoms.

2. Atoms are indivisible and indestructible

3. Atoms of a given element are identical in size, mass and chemical properties.

Dalton’s Atomic Theory

4. Atoms of a specific element are different from those of another element.

5. Different atoms combine in simple whole-number ratios to form compounds

6. In a chemical reaction, atoms are separated, combined, or rearranged.

Dalton’s theory needed only slight changes

• Atoms can be divisible (rule 2)– Protons, neutrons,

and electrons

• Atoms of the same element can have different masses (rule3)– Isotopes – different forms

of the same elements

Carbon-13, Uranium -235

Daltons Theory needed small revisions over time

Section 4-2

Section 4.2 Defining the Atom

• Define atom.

model: a visual, verbal, and/or mathematical explanation of data collected from many experiments

• Distinguish between the subatomic particles in terms of relative charge and mass.

• Describe the structure of the atom, including the locations of the subatomic particles.

Section 4-2

Section 4.2 Defining the Atom (cont.)

atom

cathode ray

electron

nucleus

proton

neutron

An atom is made of a nucleus containing protons and neutrons; electrons move around the nucleus.

Section 4-2

The Atom

• The smallest particle of an element that retains the properties of the element is called an ____________.

• An instrument called the scanning tunneling microscope (STM) allows individual atoms to be seen.

Discovery of the electron

Use of the cathode ray tube

Different gases at very low pressure, with an electric current running through

Results

1)different colors produced by different gases

2)paddle wheel moves in ray = Mass

3)deflected away from (-) magnet = (-) charge

The Atom

Section 4-2

The Cathode Ray Tube

• When an electric charge is applied, a ray of radiation travels from the cathode to the anode, called a ___________________.

• Cathode rays are a stream of particles carrying a negative charge.

• The particles carrying a negative charge are known as _______________.

Section 4-2

The Electron (cont.)

• This figure shows a typical cathode ray tube.

Section 4-2

The Electron (cont.)

J.J. Thomson measured the effects of both magnetic and electric fields on the cathode ray

Could calculate charge / mass ratio of electron

Ratio says electron smaller than Hydrogen, the smallest element

• Thomson received the Nobel Prize in 1906 for identifying the first subatomic particle—the electron

J.J. Thomson• Experiment

– Cathode Ray Tube

• Outcomes – Proof of subatomic

particles

– Evidence of (-) charge

– Charge / mass ratio

– Ratio was same for all metals in electrode and all gases in tube

Actual mass & charge of electron

• American physicist Robert Millikan, 1909

• Millikan’s Oil drop experiment

• Very small mass: 9.109 x 10-28 g

• 1/1840 the mass of a hydrogen atom

Actual Charge and Mass of Electron

Robert Millikan

• Experiment– Oil Drop Experiment

• Outcome– Electric charge from

an electron

Millikan’s Oil Drop Experiment

Section 4-2

The Electron (cont.)

• 1.602 10–19 coulombs, the charge of one electron (now equated to a single unit, -1).

• With the electron’s charge and charge-to-mass ratio known, Millikan calculated the mass of a single electron.

the mass of a hydrogen

atom

Questions left unanswered

• Atoms are neutral, so there must have a positive to balance electrons

• Because electrons are so much smaller in mass than atoms, some additional particles must be present to account for mass

• How are particles arranged in atom– J.J. Tomson proposed a plum pudding model

JJ. Thomson’s Plum Pudding (choc. chip) model of the atom

Ernest Rutherford

• Experiment– Gold Foil Experiment

• Outcome– Nucleus is discovered – Evidence shows atoms

are mostly empty space

Section 4-2

The Nucleus

• In 1911, Ernest Rutherford studied how positively charged alpha particles interacted with solid matter.

• By aiming the particles at a thin sheet of gold foil, Rutherford expected the paths of the alpha particles to be only slightly altered by a collision with an electron.

Rutherford’s quote

• "It was as if you fired a 15-inch shell at a sheet of tissue paper and it came back to hit you."

Finding’s of foil experiment• Rutherford waited 2 years before proposing his answer

• ______________________ is positively charged, very dense, central portion of atom that contains nearly all of the mass of the atom

• 2 types of particles in the nucleus– Protons, 1.673 x 10-24 g– Neutrons, 1.675 x 10-24 g, discovered by James Chadwick in

1932– 1,836 times larger than the mass of an electron– 99.95% of the mass of hydrogen-1.

JJ Thomson’s Plum Pudding Model vs.

Rutherford’s nucleus

Section 4-2

The Nucleus (cont.)

• James Chadwick received the Nobel Prize in 1935 for discovering _______________, which are neutral particles in the nucleus which accounts for the remainder of an atom’s mass.

• Why were they discovered last?

• Subatomic particles:

Electron, e-

Proton, p+

Neutron, n0

Section 4-2

The Nucleus (cont.)

• All atoms are made of three fundamental subatomic particles: the electron, the proton, and the neutron.

• Atoms are spherically shaped.

• Atoms are mostly empty space, and electrons travel around the nucleus held by an attraction to the positively charged nucleus.

Section 4-2

The Nucleus (cont.)

• Scientists have determined that protons and neutrons are composed of subatomic particles called quarks.

• Read about the Higgs Boson particle

Section 4-2

The Nucleus (cont.)

Chemical behavior of atoms can be explained by considering only an atom's ___________________.

The vast majority of chemistry deals with the interactions of electrons.

Section 4-3

Section 4.3 How Atoms Differ

• Explain the role of atomic number in determining the identity of an atom.

• Define an isotope.

• Explain why atomic masses are not whole numbers.

• Calculate the number of electrons, protons, and neutrons in an atom given its mass number and atomic number.

Section 4-3

Section 4.3 How Atoms Differ (cont.)

atomic number

isotopes

mass number

The number of protons and the mass number define the type of atom.

periodic table: a chart that organizes all known elements into a grid of horizontal rows (periods) and vertical columns (groups or families) arranged by increasing atomic number

atomic mass unit (amu)

atomic mass

Section 4-3

Atomic Number

• Each element contains a unique positive charge in their nucleus.

• The number of protons in the nucleus of an atom identifies the element and is known as the element’s ________________________.

Atomic number, mass number, and atomic mass

• Atomic number– # of protons– If neutral, it is also the #

of electrons

• Mass number– # of protons and

neutrons

• Mass number is used to calculate the atomic mass of the isotope

• Since electrons are so much smaller than protons and neutrons, they are not a major factor in atomic mass

Section 4-3

Isotopes and Mass Number

• All atoms of the same element have the same number of protons but the number of neutrons in the nucleus can differ.

• Atoms with the same number of protons but different numbers of neutrons are called _____________________.

Isotopes

• 2 ways of writing isotopes

• Carbon -13

• 136 C

Isotopes of Hydrogen

• Protium, Hydrogen-1– 99.985% of all natural hydrogen

• Deuterium, Hydrogen-2– 0.015% of all natural hydrogen

• Tritium, Hydrogen-3– Radioactive– Spiderman 2

Section 4-3

Isotopes and Mass Number (cont.)

The relative abundance of each isotope is usually constant.

Relative abundance = How much of each isotope is present

• Isotopes containing more neutrons have a greater mass.

• Isotopes have the same chemical behavior.

• The ____________________is the sum of the protons and neutrons in the nucleus.

Section 4-3

Isotopes and Mass Number (cont.)

Section 4-3

Mass of Atoms

• One atomic mass unit (amu) is defined as 1/12th the mass of a carbon-12 atom.

• One amu is nearly, but not exactly, equal to one proton or one neutron.

Relative Atomic Mass

• Atomic mass is the mass of an atom expressed in a.m.u

• Hydrogen-1 has atomic mass of 1.007825 u– Very close to mass number

• Oxygen-16 has atomic mass of 15.994915 u– Very close to mass number

Section 4-3

Mass of Atoms (cont.)

• The _________________________of an element is the weighted average mass of the isotopes of that element.

Weighted average of the isotopes of an element

Copper-63, 69.17%, 62.939 a.m.u

Copper-65, 30.83%, 64.927 a.m.u

Average atomic mass

• Atomic Mass vs Molar Mass

– Both are the weighted average of the isotopes of that element.

– 1 atom of carbon weighs 12.011 a.m.u - on weighted average

– 1 mole of carbon atoms weigh 12.011 g - on weighted average