Chapter 4: The Structure of the Atom

Atoms are the fundamental blocks of matter…

Chapter Big Idea

Section 1: Early Ideas About Matter

• What are the similarities and differences of the

atomic models of Democritus, Aristotle, and Dalton?

• How was Dalton’s theory used to explain the

conservation of mass?

Section 1: Essential Questions & Vocabulary

• Dalton’s atomic theory

• Theory

Vocabulary

Section 1: Big Idea

The ancient Greeks tried to

explain matter, but the

scientific study of the atom

began with John Dalton in the

early 1800’s.

•How was science thousands of

years ago different from science

now?

• Lacked controlled experimentation &

tools for scientific investigations

• Intellectual thought as truth

Roots of Atomic Theory

• Philosophers – scholarly thinkers

• Speculated about the nature of matter & formulated their own

explanations based on their own life experiences

• Common Conclusions:

• Matter composed of things such as earth, water, air, and fire.

• Matter could easily be divided into smaller and smaller pieces.

Roots of Atomic Theory

• Greek philosopher

• First to propose that matter was NOT infinitely divisible.

• Atomos- Greek word meaning INDIVISIBLE.

• Matter is composed of atoms which move through empty space

• Size, shape, and movement of atoms determine the properties of matter

Democritus (460 BC -370 BC)

Democritus – Atomic Model & Analogy ~ 400 BC

Atomic Model Analogy Atoms are small, hard particles that are all made of the same material, but are formed into different shapes and sizes

Legos

• What holds the atoms together?

THINK – PAIR - SHARE

• Why do you think it was hard for Democritus to defend his ideas?• Lack of experimentation

• Ahead of his time

Democritus – Criticism

• One of the most influential Greek

Philosophers

• Rejected Democritus’ notion of atoms

because it contradicted his own ideas

about nature.

• Because he was so influential, this led to

Democritus’ atomic theory to be rejected.

Aristotle (384 – 322 BC)

Aristotle – Atomic Model & Analogy

~ 300 BC – 1800’sAtomic Model Analogy

• All matter was made of only four elements & four Properties• Fire, air, water, and earth• Hot, cold, dry, and wet

Death of Chemistry for 2000 years!!!

• Marks the beginning of modern atomic theory

• Revived Democritus’ idea of atoms based on the results of his scientific research • Studied numerous chemical reactions

• Determined the mass ratios of the elements involved in those reactions

John Dalton (1766 – 1844)

• Matter is composed of extremely small particles called atoms

which are indivisible and indestructible.

• Atoms of a given element are identical in size, mass, and

chemical properties. Atoms of a specific element are different

from those of another element.

• Different atoms combine in simple whole-number ratios to form

compounds.

• In a chemical reaction, atoms are separated, combined, or

rearranged.

Dalton’s Atomic Theory – 1803

Dalton – Atomic Model & Analogy 1803

Atomic Model Analogy • All matter is made of atoms. • All atoms of a given element are alike, atoms of

different elements are different.• Atoms combine in whole-number ratios.• In chemical reactions, atoms are separated,

combined, or rearranged.

Billiard Ball

• Law of conservation of Mass- mass is conserved in

any process.

• Dalton’s Atomic Theory – explains the conservation

of mass in chemical reactions as the result of

SEPARATION, COMBINATION, or REARRANGEMENT of

atoms.

• Atoms are NOT created, destroyed or divided in the

process.

Conservation of Mass

• Which of these reactions show Dalton’s Theory?

Dalton’s Theory : Practice Question

• Which of these reactions show Dalton’s Theory?

Dalton’s Theory : Practice Question Solution

•Six atoms of Element A combine with eight

atoms of Element B to produce six compound

particles.

• How many atoms of Element A does each compound

particle contain?

• How many atoms of element B does each compound

particle contain?

• Are all of the atoms used to form compounds?

Dalton’s Theory: Practice Question

Dalton’s Theory Practice Solution

Element A (6)

Element B (8)

Compound (6 units)

• Only have 6 A Elements – can have up to 6 compound units

• Although you have 8 B Elements – not enough for 2 per compound• Must have 2 leftover B Elements

Section 2: Defining the Atom

• What is an atom?

• How can the subatomic particles be distinguished in terms of relative charge and mass?

• Where are the locations of the subatomic particles within the structure of the atom?

Section 2: Essential Questions & Vocabulary

• Atom

• Cathode ray

• Electron

• Nucleus

• Proton

• Neutron

• Model

Vocabulary

Section 2: Main Idea

An atom is made up of a

nucleus containing

protons and neutrons;

electrons move around

the nucleus.

•Smallest particle of an element that retains the properties of the element.

•How big is an atom?• ~1.3 x 10-10 m

What is an Atom?

If you could increase the size

of an atom to make it as big as

an orange. In this new scale, the

orange would be as big as Earth!!

WHAT IS AN ATOM LIKE?

Now that scientists were convinced on the existence of atoms, a new set of

questions arises!!!

• When an electric charge is applied, a ray of radiation

travels from the cathode to the anode, called a cathode

ray.

• Cathode rays are a stream of particles carrying a negative

charge.

Cathode Ray Tube

Link to video on Cathode Ray Tubes by clicking on the image

• Completed a series of cathode-ray tube experiments at Cambridge University.

• Studied mathematics and physics

• Won the Nobel Prize in 1906

J.J Thomson (1856-1940)

JJ Thomson: Discovery of the Electron

1897 Measured the effects of both magnetic and

electric fields on the cathode ray to determine

the charge to mass ratio.

Particles that compose cathode rays are negatively

charged

Charge to mass ratio of the cathode-ray is always the

same.

Concluded that all cathode rays are composed of

identical negatively charged particles called

ELECTRONS

Experiments revealed the electron has a very large

charge for its tiny mass

PROOF – THERE

MUST BE A

PARTICLE SMALLER

THAN THE

ATOM!!

JJ Thomson – Atomic Model & Analogy

1897Atomic Model Analogy

• Discovered the presence of a negative particles in the atom.

• Atoms are made from a positively charged substance with negatively charged electrons scattered about, like raisins in a pudding or chocolate chips in a cookie.

Plum PuddingOr

Chocolate Chip Cookie

• American Physicist

• Determined the charge of an electron

• Nobel Prize Winner in 1923

Robert Millikan (1868 – 1953)

• Oil Droplets are dropped

• X-Ray knocks electrons from air which then attach to the oil droplets.

• Millikan could vary the electric field strength to makes the oil droplets move more slowly, rise, or become suspended.

• Could calculate the charge on the droplets based on their rate of fall.

• The magnitude of charge on the drops always changed by a discrete amount • Smallest common denominator = 1.602 x

10 -19 C

Oil Drop Experiment – Charge of an Electron

~1910

• JJ Thomson - charge to mass ratio of the electron.

•Millikan - charge of the electron.

•CAN THOMSON’S INFORMATION ALONG WITH MILLIKAN’S BE USED TO GET THE MASS OF THE ELECTRON?

The Electron – Mass???

Mass of electron = 9.1 x 10 -28 g

• Recap of what we know: • Matter is neutral!

• All matter is made up of atoms which have electrons (- charge)

• Electrons are much lighter than the lightest atom known!

• Questions:

• If electrons are part of all matter and they possess a negative charge, how can matter be neutral?

• If the mass of an electron is so small, what accounts for the mass of a typical atom?

Plum Pudding Model New questions arise

• JJ Thomson’s student

• Early work – discovered radioactive half-life

• Nobel Prize in Chemistry – 1908

• Father of nuclear physics

• Studied how alpha particles interacted with

matter

• Alpha particles – positively charged particles

Ernest Rutherford (1871-1937)

• Aimed a narrow beam of alpha particles at a thin sheet of gold

foil.

• A flash of light is produced when the particle strikes the gold foil

• WHAT DID HE EXPECT?

• Positive charge is evenly distributed

• Path of α – particle should not be altered

• α – should continue in a straight path

Rutherford’s Gold Foil Experiment

Rutherford’s Gold Foil ExperimentWhat actually happened!!

• Atoms consist of mostly empty space

• Almost all of the atom’s positive charge and mass is contained in a tiny dense region at the center of the atom

• He called this dense region – the NUCLEUS

• Electrons are held within the atom by their attraction to the positively charged nucleus• Opposite charges attract

Rutherford’s Gold Foil Experiment Conclusions

Rutherford – Atomic Model & Analogy

1911Atomic Model Analogy

• Atoms have a small, dense, positively charged center that he called NUCLEUS

• Nucleus is tiny compared to the atom because the atom is mostly empty space

Cherry with a Pit

• Worked with Rutherford (after WWI)

• Proved the existence of neutrons

• Elementary charges devoid of any electrical

charge.

• Nobel Prize in Physics – 1935

• Manhattan Project – Development of

Atomic Bomb

James Chadwick (1891-1974)1932 - Neutron

Bohr – Atomic Model & Analogy 1913

Atomic Model Analogy• Theorized electrons move in definite

orbits around the nucleus like planets circle the sun.

• Energy levels are located at certain distances from the nucleus.

Solar System

Modern – Atomic Model & Analogy

Late 20th Century – 21st CenturyAtomic Model Analogy

• Schrodinger, Heisenberg, Einstein & many other scientists• Electrons move at high speeds in an electron could around

the nucleus• In the ELECTRON CLOUD, electrons orbit around the

nucleus billions of times in one second• Electron’s motion is dependent on the AMOUNT of ENERGY

they contain

Cotton Balls

Atomic Theory Ted Ed talk !!

• All atoms have 3 fundamental subatomic particles • Protons, neutrons, electrons

• Atoms are spherically shaped, with small, dense positively charged nucleus surrounded by negatively charged electrons.

• Most of an atom consists of fast moving electrons that are held within the atom by their attraction to the positively charged nucleus.

• Nucleus is composed of neutrons (neutral charge) and protons (positive charge)

• Scientists have determined that protons and neutrons are composed of “quarks”• Scientists unsure if and how “quarks” affect chemical behavior.

• Chemical behavior can be explained by considering only the atom’s electrons.

Completing the Model of the Atom

Section 3: How Atoms Differ

• His the atomic number used to determine the identity of an atom?

• What is an isotope?

• Why are atomic masses not whole numbers?

• Given the mass number and the atomic number, how are the number of electrons, protons, and neutrons in an atom calculated?

Section 3: Essential Questions & Vocabulary

Vocabulary• Atomic Number

• Isotope

• Mass Number

• Atomic Mass Unit (AMU)

• Atomic Mass

• ATOMIC NUMBER – the number of protons in the nucleus

of an atom

• the number of protons in an atom identifies it as an atom of a

particular element.

Atomic Number

• All atoms are neutral.

• So the number of

protons must be equal

to the number of

electrons.

Atomic Number Practice

Atomic Number Practice

Atomic Number Practice SOLUTIONS

Oxygen

Sodium

Chlorine

Uranium

6 6

8

11

92

17

• Sum of the number of protons and neutrons in the

nucleus

• Always a whole number.

• Not on the periodic table.

Mass Number

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Mass Number : Practice

Isotopes• Atoms of the same element with different mass

numbers.

C126Mass #

Atomic #

¨ Nuclear symbol:

¨ Hyphen notation: carbon-12

• If isotopes are the same element but have different

mass number, what is different in each isotope?

• Number of Neutrons

Isotope Question?

© Addison-Wesley Publishing Company, Inc.

• In nature, most elements are found as mixtures of

isotopes.

• Usually, the relative abundance of each isotope is

constant regardless of where the element is obtained.

Natural Abundance of Isotopes

C. Johannesson

Relative Atomic Mass

• 12C atom = 1.992 × 10-23 g

¨ 1 p = 1.007276 amu1 n = 1.008665 amu1 e- = 0.0005486 amu

© Addison-Wesley Publishing Company, Inc.

¨ atomic mass unit (amu)

¨ 1 amu = 1/12 the mass of a 12C atom

Atomic Mass• Weighted AVERAGE of all isotopes of that element

• Atomic mass found on the Periodic Table

• Round to 2 decimal places Write the relative

abundance percentage as

a decimal.

Average Atomic Mass : Practice

• EX: Calculate the avg. atomic mass of oxygen if its

abundance in nature is 99.76% 16O, 0.04% 17O, and 0.20% 18O.

Average Atomic Mass Practice Solution

• EX: Calculate the avg. atomic mass of oxygen if its

abundance in nature is 99.76% 16O, 0.04% 17O, and 0.20% 18O.

16O = 16 x 0.9976 = 15.9616

17O = 17 x 0.0004 = 0.0068

18O = 18 x 0.0020 = 0.0360+Average Atomic Mass for O = 16. 00 amu

Add all isotope masses

together!!

Average Atomic Mass : Practice• EX: Find chlorine’s average atomic mass if approximately 8 of

every 10 atoms are chlorine-35 and 2 are chlorine-37.

Average Atomic Mass : PracticeSolution• EX: Find chlorine’s average atomic mass if approximately 8 of

every 10 atoms are chlorine-35 and 2 are chlorine-37.1. Find the relative percentage of each isotope• Chlorine – 35 : 8 out of 10 atoms

= 80 % • Chlorine – 37 : 2 out of 10 atoms

= 20 % 2. Calculate the average atomic mass

35Cl = 0.80 x 35 = 28.0 37Cl = 0.20 x 37 = 7.4+

Average Atomic Mass = 35.4 amu