1. Chapter 4 notes
Atomic Structure

2. Section 4.1 Early Ideas About Matter
Compare and contrast the atomic models of Democritus, Aristotle, and Dalton.
Understand how Dalton's theory explains the conservation of mass.
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
Section 4-1

3. Greek Philosophers (cont.)
Many ancient scholars believed matter was composed of such things as earth, water, air, and fire.
Another common conclusion was that matter could be endlessly divided into smaller and smaller pieces.
Section 4-1

4. 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. (in Greek atomos means “indivisible”)
Section 4-1

5. Greek Philosophers (cont.)
Section 4-1

6. Aristotle (484–322 B.C.) rejected Democritus’ ideas because they didn’t agree with his own. He believed empty space could not exist. .

7. Greek Philosophers (cont.)
Democritus’ ideas weren’t “science” because there were no controlled experiments.
Aristotle’s views went unchallenged for 2,000 years until science developed methods to test the validity of his ideas
Section 4-1

8. Greek Philosophers (cont.)
John Dalton revived the idea of the atom in the early 1800s based on numerous chemical reactions. .
Section 4-1

9. atoms of a given element were identical (because isotopes exist)
Dalton’s atomic theory easily explained conservation of mass in a reaction as the result of the combination, separation, or rearrangement of atoms.
He was wrong about:
atoms of a given element were identical (because isotopes exist)
Atoms are indivisible (because atoms can be divided into smaller parts)
Section 4-1

10. Section 4.2 Defining the Atom
Define atom.
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.
model: a visual, verbal, and/or mathematical explanation of data collected from many experiments
Section 4-2

11. the world’s population in 2006 was 6.5 x 109 people
The Atom
The smallest particle of an element that retains the properties of the element is called an atom.
Note :
the world’s population in 2006 was 6.5 x 109 people
A normal penny would contain 2.9 x 1022 atoms
That would be ____________ atoms per person!
Placing 6.5 billion atoms side by side would be less than a meter long!
Section 4-2

12. An instrument called the scanning tunneling microscope (STM) allows individual atoms to be seen.
The field of nanotechnology promises to build machines atom-by-atom. This is molecular manufacturing!

13. The Electron
Cathode ray tubes (CRT’s) are glass tubes with a vacuum inside and electricity is passed through. They are affected by magnets. They allow researchers to study the relationship between mass and charge. The cathode is the negative electrode; The anode is the positive electrode.

14. Sir WilliamCrookes discovered cathode rays.
The Electron
Sir WilliamCrookes discovered cathode rays.
Cathode rays are a stream of radiation particles that travel from the cathode to the anode carrying a negative charge. The particles carrying a negative charge are known as electrons.
Cathode ray technology goes into television sets
2 important points of cathode rays:
they are a stream of charged particles
they are negatively charged (electrons!)
Section 4-2

15. This figure shows a typical cathode ray tube.
The Electron (cont.)
This figure shows a typical cathode ray tube.
Section 4-2

16. The Electron (cont.)
J.J. Thomson measured the effects of both magnetic and electric fields on the cathode ray to determine the charge-to-mass ratio of a charged particle, then compared it to known values.
The mass of the charged particle was much less than a hydrogen atom, then the lightest known atom.
Thomson received the Nobel Prize in 1906 for identifying the first subatomic particle—the electron
Section 4-2

17. The Electron (cont.)
In the early 1910s, Robert Millikan used the oil-drop apparatus shown below to determine the charge of an electron.
Section 4-2

18. the mass of a hydrogen atom
The Electron (cont.)
Charges change in discrete amounts—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
Section 4-2

19. Matter is neutral; to explain this there are models.
The Electron (cont.)
Matter is neutral; to explain this there are models.
J.J. Thomson's plum pudding model of the atom states that the atom is a uniform, positively changed sphere containing electrons.
Section 4-2

20. The Nucleus (cont.)
In 1911, Ernest Rutherford studied how positively charged alpha particles interacted with solid matter
Although most of the alpha particles went through the gold foil, a few of them bounced back, some at large angles.
Section 4-2

21. Rutherford concluded that: atoms are mostly empty space.
The Nucleus (cont.)
Rutherford concluded that:
atoms are mostly empty space.
almost all of the atom's positive charge and almost all of its mass is contained in a dense region in the center of the atom called the nucleus.
Electrons are held within the atom by their attraction to the positively charged nucleus.
Section 4-2

22. The Nucleus (cont.)
Like charges repel; the positively charged nucleus and positive alpha particles caused the deflections.
Section 4-2

23. If the nucleus were the size of a dot on a piece of paper, it would weigh approximately as much as 70 cars!
This tells you that that the nucleus is REALLY dense and has almost all the weight of an atom
If the nucleus were the size of a nickel, the atom would be about the size of two football fields!
This tells you how small the nucleus is in the atom

24. The Nucleus (cont.)
Rutherford refined the model to include positively charged particles in the nucleus called protons.
James Chadwick received the Nobel Prize in 1935 for discovering the existence of neutrons, neutral particles in the nucleus which accounts for the remainder of an atom’s mass.
Section 4-2

25. Atoms are spherically shaped.
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

26. The Nucleus (cont.)
Scientists have determined that protons and neutrons are composed of subatomic particles called quarks.
Section 4-2

27. Chemistry can be explained by considering only an atom's electrons.
The Nucleus (cont.)
Chemistry can be explained by considering only an atom's electrons.
Section 4-2

29. The element’s atomic number
(symbol Z) is defined as the number of protons in the nucleus.
This number identifies the element. It is always a whole number on any periodic table.
Section 4-3

30. It is always a whole number on any periodic table
It is always a whole number on any periodic table. For example, look at a periodic table and find atomic number 80. That means you have mercury and mercury has 80 protons. ONLY mercury has 80 protons!

31. Atoms are always neutral because in any atom, the number of protons is also the number of electrons [because (+) = (-) to be neutral]

32. Isotopes containing different neutrons have different masses.
Isotopes and Mass Number
Isotopes are atoms with the same number of protons but different numbers of neutrons .
Isotopes containing different neutrons have different masses.
Isotopes have the same chemical behavior; you can’t tell them apart in a chemistry lab.

33. The mass number (symbol A) is the sum of the protons and neutrons in the nucleus. It is always a whole number

34. Mass number is not to be confused with atomic mass which is the decimal number on the periodic table. The atomic mass for an element is the average mass of all the isotopes of the elements.

35. Isotopes have different mass numbers.
Ex: C-12 and C-14 are two isotopes of carbon with mass numbers of 12 and 14.
The atomic mass of carbon is on the periodic table and is

36. Ions
ions are charged atoms and they are formed by a gain or loss of electrons. Many ions will result from an atom reacting to gain or lose electrons so that they look like a noble gas (column 18)
Section 4-3

37. Negative ions (anions) have gained electrons
They are formed from NONMETALS (right of stair like line on periodic table) Ex

38. positive ions (cations) have lost electrons
positive ions (cations) have lost electrons. They are formed from METALS (left side of periodic table)
Ex: calcium atom loses two electrons to form calcium ion
The charge of an ion is the sum of the charged particles [sum of protons (+) and electrons (-)]

39. may be represented as U – 238
How to calculate the # of protons, electrons, and neutrons:
Protons = Z
electrons = Z if atom is neutral; negative ions have extra electrons; positive ions have fewer electrons.
neutrons = A - Z
may be represented as U – 238
238 is the Mass number (A)
238 is the Mass number (A)
92 is the Atomic number (Z)

40. Ways to represent a “nuclide” (a nuclide is any atom with a specific number of protons and neutrons)
Pa – may be represented as
it has __ protons, ___, electrons & _____ neutrons
Notice how the “231” can be written in two different ways; it is the mass number (A). The “91” may or may not be written; it is the atomic number (Z) and can be found on any periodic table. Since this is neutral (no charge), number of electrons will equal number of protons. Neutrons = A - Z

41. has __protons, __ electrons & __ neutrons
Os – 195 has ____ protons, ____ electrons, ____ neutrons
Os – 188 has ____ protons, ____ electrons, ____ neutrons
has __ protons, ___ electrons, ___ neutrons
Notice how this is an ion; the 2- means it has 2 extra electrons
The number positive charges (protons) plus the number of negative charges (electrons) must equal the overall charge = -2
has _ protons, _ electrons, ___ neutrons
Notice how this ion has a 1+, meaning it has 1 less electron.
= +1

42. there are three isotopes of hydrogen, with special names
Protium is H-1 Deuterium is H-2 Tritium is H-3
How many protons, electrons, neutrons are in each?

43. Isotopes and Mass Number (cont.)
Section 4-3

44. How many protons, electrons, and neutrons in each?
Pt – 195
Hg – 200
Hg – 201
Ca – 40 Ca
Ca – 39
As – 75 As
Which of the isotopes listed of mercury is the more common? Hg-201
Which of the calcium isotopes is more common? Ca-40
How do you know? Because the most common isotope has a mass number nearest the average atomic mass found on a periodic table.

45. Mass of Atoms
One atomic mass unit (amu or u) is defined as 1/12th the mass of a carbon-12 atom. One amu is nearly, but not exactly, equal to one proton and one neutron.
Section 4-3

46. Mass of Atoms (cont.)
The atomic mass of an element is the average mass of the isotopes of that element. It is usually the decimal number on any periodic table.
How to calculate atomic mass:
atomic mass = (percent)(mass) + (percent)(mass) + etc
[as many times as needed]
Example: If 75.78% of an element has a mass of amu and % has a mass of amu, what is the average mass? And what element is it?
Section 4-3

48. Section 4.4 Unstable Nuclei and Radioactive Decay
Explain the relationship between unstable nuclei and radioactive decay.
Characterize alpha, beta, and gamma radiation in terms of mass and charge.
element: a pure substance that cannot be broken down into simpler substances by physical or chemical means
Section 4-4

49. Radiation is the rays and particles emitted in radioactivity.
In a Nuclear reaction, one element changes into another element through a change in the nucleus.
radioactivity is spontaneously emitted radiation. Spontaneous means it happens on its own; no outside force is required. It was discovered in the late 1890s
Radiation is the rays and particles emitted in radioactivity.
A reaction that involves a change in an atom's nucleus is called a nuclear reaction.
Section 4-4

50. Radioactive Decay
radioactive decay is unstable nuclei losing energy by emitting radiation spontaneously; they form more stable elements – often lead.
Section 4-4

51. Carbon-14 used in radioactive dating
C-14 has a “half-life” of 5770 years
“Half life” is time it takes for ½ of a radioactive sample to “decay” and form other elements.
How much C-14 would be left of a sample after 17,310 years if the original amount was 2.00 grams?

52. Radioactive Decay (cont.)
Alpha radiation (the least damaging type of radiation) is deflected toward a negatively charged plate and is made up of positively charged particles called alpha particles.[note – these were the same particles used in the “gold foil” experiment covered earlier in this chapter]
Each alpha particle contains two protons and two neutrons and has a 2+ charge.
Symbol is or α [Greek letter alpha]
Section 4-4

53. Radioactive Decay (cont.)
A new element can be formed only in a nuclear equation. The following shows the radioactive decay of radium-226 to radon-222 and an alpha particle.
NOTE: The mass is conserved in nuclear equations (total mass on left of arrow = total on right)
Section 4-4

54. Radioactive Decay (cont.)
Beta radiation is deflected toward a positively charged plate. It is radiation that has a negative charge and emits beta particles.
Each beta particle is an electron with a 1– charge. Its symbol is β or
Section 4-4

55. Radioactive Decay (cont.)
Section 4-4

56. Radioactive Decay (cont.)
Gamma rays are high-energy radiation with no mass and are neutral.
Gamma rays account for most of the energy lost during radioactive decay.
Section 4-4

57. Radioactive Decay (cont.)
Stability of nuclei:
Atoms that contain too many or too few neutrons are unstable and lose energy through radioactive decay to form a stable nucleus.
Few radioactive isotopes exist in nature—most have already decayed to stable forms.
Section 4-4

58. Examples of nuclear reactions:
Alpha decay of Sm-146
Beta decay of Xe-152