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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
© Addison-Wesley Publishing Company, Inc.
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