1. The Building Blocks of Matter
Chapter 3 - Atoms:
The Building Blocks of Matter

2. From Philosophy to Science
There were two schools of thought of the composition of the cosmos…
is everything in the universe continuous and infinitely divisible
Or, is there a limit to how small you can get?
Particle theory was not the most popular early opinion, but was supported as early as Democritus in ancient Greece.
From Philosophy to Science

3. From Philosophy to Science
Democritus proposed that all the matter is composed of tiny particles called “Atomos”
These “particles” were thought to be indivisible
Aristotle did not accept Democritus’ atom, he was of the “matter is continuous” philosophy
Because of Aristotle’s popularity his theory was adopted as the standard
From Philosophy to Science

4. From Philosophy to Science
By the 1700’s nearly all chemists had accepted the modern definition of an element as a particle that is indivisible
It was also understood at that time that elements combine to form compounds that are different in their properties than the elements that composed them
However, these understandings were based on observations not empirical evidence
From Philosophy to Science

5. From Philosophy to Science
There was controversy as to whether elements always combine in the same proportion when forming a particular compound.
In the 1790’s, chemistry was revolutionized by a new emphasis on quantitative analysis because of new and improved balances
This new technology led to the discovery of some new scientific understandings
From Philosophy to Science

6. From Philosophy to Science
The Law of Conservation of Mass:
Proposed by Antoine Lavoisier
States that mass is neither created nor destroyed during ordinary chemical reactions or physical changes.
Which means the total mass of the reactants must equal the total mass of the products.
From Philosophy to Science

7. Law of Conservation of Mass+ 
Carbon, C
Oxygen, O
Carbon Monoxide, CO
Mass x
Mass y
Mass x + Mass y
Carbon Monoxide, CO
Mass x + Mass y +
Carbon, C
Oxygen, O
Mass x
Mass y 

8. From Philosophy to Science
The Law of Definite Proportions:
The fact that a chemical compound contains the same elements in exactly the same proportions by mass regardless of the size of the sample or the source of the compound
NaCl is NaCl no matter if it is table salt (small crystals) or rock salt (large crystals)
From Philosophy to Science

9. From Philosophy to Science
The Law of Multiple Proportions:
If 2 or more different compounds are composed of the same 2 elements, then the ratio of the masses of the 2nd element combined with a certain mass of the 1st element is always a ratio of small whole numbers
From Philosophy to Science

10. Law of Multiple Proportions+ =
Carbon
Oxygen
Carbon Monoxide, 1:1 1 1 +
Carbon
Oxygen =
Carbon Dioxide, 1:2 1 2

11. In 1808, John Dalton proposed an explanation for each of the proposed laws
He reasoned that elements were composed of atoms & that only whole #’s of atoms can combine to form compounds
His ideas are now called the Atomic Theory of Matter
Atomic Theory

12. All matter is composed of extremely small particles called atoms.
Atoms of a given element are identical in size, mass, and other properties; atoms of different elements differ in size, mass, & other properties.
ELEMENT 1
ELEMENT 2
ELEMENT 3
ELEMENT 4
Atomic Theory

13. Atomic Theory Atoms cannot be subdivided, created, or destroyed
atoms of different elements combine in simple whole # ratios to form chem compds
in chemical rxns, atoms are combined, separated, or rearranged + +
Atomic Theory

14. Through these statements, evidence could be gathered to confirm or discount its claims
Not all of Dalton’s claims held up to the scrutiny of experimentation
Atoms CAN be divided into even smaller particles
Not every atom of an element has an identical mass
Atomic Theory

15. Atomic Theory Dalton’s Atomic Theory of Matter has been modified.
What remains is…
All matter is composed of atoms
Atoms of any one element differ in properties from atoms of another element
One of the disputed statements of Dalton was that atoms are indivisible
Atomic Theory

16. In the 1800’s it was determined that atoms are actually composed of several basic types of smaller particles
it’s the number and arrangement of these particles that determine the atom’s chemical properties.
The def. of an atom that emerged was, the smallest particle of an element that retains the chemical properties of that original element.
Atomic Theory

17. All atoms consist of 2 regions that contain the subatomic particles
The nucleus
The electron cloud around the nucleus
The nucleus is a very small region located near the center of the atom
In every atom the nucleus contains at least 1 proton, which is positively charged particle and usually contains 1 or more neutral particles called neutrons
Atomic Structure

18. The electron cloud is the region that surrounds the nucleus
This region contains 1 or more elec-trons, which are negatively charged subatomic particles
The volume of the electron cloud is much larger than the nucleus
Atomic Structure

19. Discovery of the Electron
The discovery of the first subatomic particle took place in the late 1800’s.
A power source was attached to two metal ends of an evacuated glass tube, called a cathode ray tube.
A beam of “light” appears between the two electrodes called a cathode ray.
Discovery of the Electron

20. cathode ray cathode anode
Electric Current
cathode ray
cathode
anode
Cathode Ray

21. Modern Application:
TV Picture Tube

22. Discovery of the Electron
Investigators began to study the ray and they observed that…
An object placed in the path of the ray cast a shadow on the glass
A paddle wheel placed in the path of the cathode ray began to spin
Cathode rays were deflected by a magnetic field
The rays were deflected away from a negatively charged object
Discovery of the Electron

23. Discovery of the Electron
The first 2 observations support the idea that the ray is composed of tiny individual particles traveling through the vacuum tube
The second set of observations support the evidence that the ray is composed of a substance that is negatively charged.
Discovery of the Electron

24. Discovery of the Electron
J.J. Thomson studied the rays and proved that they were tiny negative particles being emitted from the metal atoms.
Dubbed these tiny particles “electrons”
And it was later determined that the electrons were not part of the mass of the atom.
Discovery of the Electron

25. What can their work help us conclude about the atom?
atoms are composed of smaller particles, and one of these compo- nents is negatively charged
atoms are neutral, so there must be an opposing (+) charge
because electrons are essentially mass-less, an opposing substance makes up the mass of the atom
First Atomic Model

26. Discovery of the Proton
In 1886, E. Goldstein observed in the cathode-ray tube a new set of rays traveling in the opposite direction than the cathode rays
The new rays were called canal rays and they proved to be positively charged
And the particles mass were about X’s that of the electron
Discovery of the Proton

27. Discovery of the Neutron
In 1932, the English physicist James Chadwick discovered yet another subatomic particle.
the neutron is electrically neutral
It’s mass is nearly equal to the proton
Therefore the subatomic particles are the electron, proton, and neutron.
Discovery of the Neutron

28. Scientists still didn’t really understand how the particles were put together in an atom.
This was a difficult question to resolve, given how tiny atoms are.
Most thought it likely that the atom resembled Thomson’s model
Atomic Structure

29. Negative particles embedded in a sphere of positive plasma-like matter.
THINK…
Chocolate Chip Cookie Mg + -

30. In 1911, Ernest Rutherford et al
In 1911, Ernest Rutherford et al. provided a more detailed picture of the internal structure of the atom
In his experiment, Rutherford directed a narrow beam of alpha particles at a very thin sheet of gold foil.
Alpha particles (a) are He atoms that have been stripped of their electrons
Rutherford Model

31. According to Thomson’s model, the heavy, positive alpha particles should pass easily through the gold, with only a slight deflection
And mostly that’s how it happened.
However, they found 1 in every particles had actually been deflected back toward the source.
Rutherford Model

33. Rutherford suggested a new structural model of the atom.
He stated that all the positive charge and the mass is concentrated in a small core in the center of the atom, AKA nucleus
And that the atom is mostly empty space with electrons surrounding the positively charged nucleus like planets around the sun.
Rutherford Model

34. Rutherford Model

35. With the exception of Hydrogen, every nucleus contains 2 kinds of particles protons and neutrons
they make up the mass of the atom (Mass Number = Protons + Neutrons)
Atoms are neutral because they contain equal #’s of protons & electrons
Atomic Structure

36. The atoms of different elements differ in the # of protons in their nuclei and therefore in their positive charge
The # of protons the atom contains determines the atom’s identity, also known as atomic number.
Only Oxygen contains 8 protons
Only Fluorine contains 9 protons
Structure of the Atom

37. The nucleus is composed of a densely packed cluster of protons, which are all electrically positive
Don’t like charges repel?
Why don’t they fly apart?
When 2 protons are in very close proximity, there is a strong force of attraction between them.
similar attraction exists when neutrons are close
Structure of the Atom

38. These short-range p+-n0, p+-p+, & n0-n0 forces hold the nuclear particles together, A.K.A strong nuclear forces.
When these nuclear forces are strong enough the atom is stable
If the forces are not strong enough the atom (heavier atoms) the atom is unstable and becomes radioactive.
Structure of the Atom

39. Ch 3.3: Atomic Number

40. # protons in an atom = # electrons
Elements are identified by the number of PROTONS they contain
The “atomic number” of an element is the number of protons in the nucleus
PROTONS IDENTIFIES AN ELEMENT!!!
# protons in an atom = # electrons
Because atoms are neutral!

41. Complete Symbol X
Mass
number
Atomic
Subscript →
Superscript →

42. Cl 35 17 # OF PROTONS + # OF NEUTRONS MASS NUMBER ATOMIC NUMBER
NUMBER OF PROTONS
# OF PROTONS +
# OF NEUTRONS

43. Mass Number
Mass number is the number of protons and neutrons in the nucleus of an isotope:
Mass # = p+ + n0
Element p+ n0 e-
Mass #
Oxygen - 10 - 33 42
- 31 15
Mass number 18 8 8 18
Arsenic 75 33 75
Phosphorus 16 15 31

44. Find the # of e-, p+ and n0 for sodium. (mass # = 23)
Practice Problems
Find the # of e-, p+ and n0 for sodium.
(mass # = 23)
Find the # of e-, p+ and n0 for uranium.
(mass # = 238)
Atomic # = 11 = # e- = # p+
# neutrons = = 12
Atomic # = 92 = # e- = # p+
# neutrons = = 146

45. Check for understanding:
If an element has 91 protons and 140 neutrons find the:
Atomic number
Mass number
number of electrons
element name 91
231 91
protactinium

46. (Isotopes: The # after the name is the mass #.)
An isotope refers to atoms that have the same # of protons, but a different number of neutrons.
Because of this, they have different mass #’s.
Examples---> (1) Carbon-12 & Carbon-13
(2) Chlorine-35 & Chlorine-37
(Isotopes: The # after the name is the mass #.)

47. Cl Cl 18 20 35 37 17 17 EXAMPLE OF AN ISOTOPE ATOMIC MASS
ATOMIC NUMBER 18
NEUTRONS 20
NEUTRONS

48. Br Question #1 80 35 Find each of these: Atomic number Mass Number
number of protons
number of neutrons
number of electrons Br 35

49. Question #2
If an element has an atomic number of 34 and a mass number of 78, what is the:
number of protons
number of neutrons
number of electrons
complete symbol

50. Units = atomic mass unit (amu) 12
Units = atomic mass unit (amu)
The atomic masses listed in the Periodic Table are a “weighted average” of all the isotopes of the element.

51. In chemistry, chlorine has 2 isotopes:
Weighted Average
Practice Problems:
In chemistry, chlorine has 2 isotopes:
Cl-35 (75.8% abundance) Cl-37 (24.23 % abundance)
What is the weighted average atomic mass of chlorine?
35 x = amu
37 x = amu
Add them up!!!
(2) Oxygen has 3 isotopes:
O-16 (99.76%) O-17 (0.037%) O-18 (0.2%)
Estimate oxygen’s average atomic mass.
Barely over 16.0 amu. +
amu

52. Relating Mass Numbers to Atoms
The Mole: the amount of a substance that contains as many particles as there are atoms in exactly 12 grams of carbon-12.
Avogadro’s Number: the number of particles in exactly one mole of a pure substance = x 1023.
Molar Mass: the mass of one mole of a pure substance. Units = g/mol

53. This is when we get to use dimensional analysis!
The conversion factors we need are:
and of course…molar mass

54. Gram to Mole Conversions
Mass of Element in Grams
Number of Moles of Element
Number of Atoms of Element

55. Practice Problem
ALWAYS USE PARANTHESES AROUND YOUR CONVERSION FACTORS!!
You have 3.50 mol of Copper. What is it mass in grams?