1. Reference Table: Table S & PT
CHEMISTRY: PACKET #3
Atomic Concepts
Reference Table: Table S & PT

2. The Atom Atoms are the building blocks of all matter.
The atom is composed of 3 SUBATOMIC PARTICLES: protons, neutrons, and electrons.
Protons and electrons have equal but opposite charges. A neutral atom must contain equal numbers of both. Each atom has a nucleus with an overall positive charge. Within the nucleus are the protons and the neutrons. Surrounding the nucleus are negatively charged electrons.
An atoms identity is defined entirely by the number of protons in the nucleus; the number of protons of any given element NEVER changes.

3. Atom = NEUTRAL #protons = #electrons

4. Outside Nucleus (Orbital)
SUBATOMIC PARTICLES
Particle
Charge
Mass
Location
Proton +1
1 amu
(atomic mass unit)
Nucleus
Neutron
Electron -1
1/1836 or 0 amu
Outside Nucleus (Orbital)

5. Atomic Charge
Nuclear Charge
Vs.

6. Atomic Mass Unit (AMU or u)
The mass of an atom is so extremely small that the atomic mass scale replaces grams as the unit used to describe the masses of atoms
The nucleus of a carbon atom containing 6 protons and 6 neutrons is taken as the standard mass for this scale
The amu is defined as 1/12 the mass of a Carbon atom or 1.66x10-24 g

7. The Atomic Model
The discovery of the atom as we know it today was a progression, like the discovery of DNA or any other major scientific discovery. Many scientists contributed to the development of present day atomic theory. Each proposed model of the atom was based on the models developed prior to it. With each new discovery dealing with the nature of the atom, a new atomic model was constructed.

8. History of Atomic Theory
John Dalton
( )
Also supported the notion of the atom
He defined an atom as: The smallest particle of an element that retains the chemical identity of that element.
Came up with 4 postulates that outlined his theory about atoms

9. Dalton’s Atomic Theory (1700s)
All matter is made up of tiny particles called atoms
All atoms of a given element are alike; while atoms of different elements are different because they have different masses
Compounds are formed when atoms of different elements combine in fixed proportions
Atoms are indivisible; atoms can neither be created nor destroyed in chemical reactions. A chemical reaction involves a rearrangement of atoms, not a change in the atoms themselves

10. J.J. Thompson (1890s) Cathode Ray Tube Experiment
Discovered Electrons, particles with a negative charge.
Cathode Ray Tube is a Sealed Glass Tube that contains a gas and has separated metals plates connected to external wires that pass an electrical current through the tube

11. Plum Pudding Model
The atom was a hard sphere of positive charge with electrons embedded in it.
Think of it as a chocolate chip cookie; positive - cookie & negative -chocolate chips (spaced apart about 50/50).

12. Further Development of the Atomic Model Ernest Rutherford
Ernest Rutherford (1910): disproved Thomson’s model
Gold Foil Experiment
VERY IMPORTANT!

13. Further Development of the Atomic Model Ernest Rutherford
Gold Foil Experiment - Observations
MOST of the alpha particles (+) went straight through
SOME of the alpha particles (+) deflected back (bounced back) to the source.
Rutherford’s Explanation
Most of the α particles pass directly through the gold foil because the gold foil atoms are composed of mostly empty space.
The particles that bounced back did so because they “directly hit” the positive center of the atoms (like charges deflect)

14. Opposite Charges Attract
Like Charges Repel

15. Bohr Model (1911)
Small, dense, positively charged nucleus surrounded by electrons in circular orbits. (a.k.a. the “planetary model”)
Bohr’s model is fundamentally incorrect because electrons do not move in fixed orbits around the nucleus like the planets do around the sun.

16. Bohr’s “Planetary” Atom Model

17. Wave-Mechanical Model (Modern Atomic Theory)
Energy and matter are now viewed as acting as both waves and particles
Small, dense, positively charged nucleus surrounded by electrons moving in "electron cloud".
"Orbitals" are areas where an electron with a certain amount of energy is most likely to be found.

18. Wave-Mechanical Model

19. Orbit vs. Orbital
An orbit describes a particular path that an object follows as it travels around another object
Electrons do not follow a particular path around the nucleus. The exact path of an electron in this area is not known
Instead, an orbital describes the areas around the nucleus where an electron is most likely to be found (probability of location)

20. Atomic Theory from Past to Present

22. Atomic Number
Located on the lower left hand in the box of the individual element on the Periodic Table. The atomic number is equal to the number of protons (and in a neutral atom, the # of electrons).
An element is identified by its number of protons which NEVER CHANGES, and therefore an elements atomic number NEVER CHANGES!!

23. Represents the number of protons in the nucleus; the NUCLEAR CHARGE

24. ISOTOPES
Atoms of the same element that have different numbers of neutrons.
There are two methods of identifying isotopes.

25. Mass Number
Whole number on the upper left hand corner identifying the number of protons and neutrons located in the nucleus (they have mass).
To calculate the #of neutrons, you subtract the Atomic Number from the Mass Number (12-6 = 6 neutrons)
You’ll notice that most of the elements on the PT have a number with a decimal, indicating an average. This number is known as the Average Atomic Mass, and is defined as the average of all of the naturally occurring isotopes of a given element.

26. These are three ISOTOPES of carbon with MASS NUMBERS OF 12, 13, and 14
Represents the average of all the naturally occurring isotopes of that element.
Number of Neutrons =
These are three ISOTOPES of carbon with MASS NUMBERS OF 12, 13, and 14

27. Calculating Average Atomic Mass
Most elements occur in nature as mixtures of isotopes:
The average atomic mass of an element can be determined by taking into account the masses of all the naturally occurring isotopes of a particular element found in nature as well as their relative abundances.
Because the abundances for each isotope is different, YOU CANNOT JUST ADD AND DIVIDE BY THE # OF ISOTOPES!!!!!!!!

28. Calculating Average Atomic Mass
Change the abundance (%) to a decimal by moving the decimal point two places to the left.
Multiply that number by the Atomic Mass of the Isotope (more specific number). Do the same for every isotope given.
ADD everything together to calculate the Average Atomic Mass.

29. Convert % to decimal. 75.76% becomes .7576 & 24.24% becomes .2424
Step #1
Convert % to decimal % becomes & 24.24% becomes .2424

30. Steps #2 & 3 (34.97u) (.7576) + (36.97u) (.2424) = 35.45u
Multiply the abundances (now in decimal) by the Atomic Mass of the Isotope (the more specific number). Do this for every isotope provided. Finally add all the values together.
(34.97u) (.7576) + (36.97u) (.2424) = 35.45u
Average Atomic Mass

31. Calculating Average Atomic Mass

32. Calculating Average Atomic Mass

33. Calculating Average Atomic Mass

34. Electron Configuration
Located on the lower left corner, below the atomic number.
Bohr’s model describes electrons in terms of energy levels.
**Electrons that are close to the nucleus have a low energy level and electron farther from the nucleus have a higher energy level.**
Based on Bohr’s model we get Principle Energy Levels. The energy level shows how far the electrons are from the nucleus. The first energy level is closest to the nucleus and the others are further away. Electrons in the first level have the lowest energy and the energy of the electrons increase as the levels increase.

35. Electron Configuration
Represents the number of electrons and which Principle Energy Level (PEL) they’re located in. In a neutral atom, the total number will equal the number of protons.

36. FIRST PRINCIPLE ENERGY LEVEL: holds a maximum of 2 electrons.
SECOND PRINCIPLE ENERGY LEVEL: holds a maximum of 8 electrons.
THIRD PRINCIPLE ENERGY LEVEL: holds a maximum of 18 electrons.
FOURTH PRINCIPLE ENEGY LEVEL: holds a maximum of 32 electrons.

37. Valence Electrons - The number of electrons in the last principle energy level. According to the octet rule, there can be no more than 8 valence electrons. These electrons affect chemical properties of the element.
Non-Valence Electrons - All other electrons in an atom other than the last level (valence)
In this example, there are 6 VALENCE ELECTRONS and 10 NON- VALENCE ELECTRONS. The valence electrons are the most important because they have the most energy (farthest from the nucleus).

38. Ground State
Electrons fill in energy levels and orbitals starting with the one that requires the least energy and progressively move to those levels and orbitals that require increasing amounts of energy.
When all electrons are at their lowest possible energy, it is called the "ground state” (configuration on the PT)
Example: Fluorine (F) has an atomic number of 9, The electron configuration that describes it in the ground state is 2-7

39. Excited State
When the electron gains a specific amount of energy from heat, light, or electricity, it moves to a higher level and is in the "excited state”. Electrons are NOT being lost or gained, just reconfigured.
Look at the electron configuration for sodium (Na) 2-8-1
If Na had the configuration of 2-7-2, you would notice that the second PEL was not completely filled with 8 electrons before the third PEL began to fill. This is an example of an electron configuration in the excited state. Notice that regardless whether it is in the ground or excited state, the total number of electrons never changes. For the example above, the number of electrons remains 11.

40. Flame Test Bright Line Spectrum
In this class we will be performing an experiment called the “Flame Test”. We will be heating up metal powders in order to excite the electrons to jump from a lower PEL to a higher PEL.

41. Bright Line Spectrum
When an electron jumps from a lower PEL to a higher PEL it gets “excited” and requires energy to do this. When the electron returns from a higher energy state to a lower energy state, it emits a specific amount of energy usually in the form of light. This is known as a bright line spectrum, and can be used to identify an element like a fingerprint.

42. Bright Line Spectrum
While this light appears as one color to our eyes, it is actually composed of many different wavelengths each of which can be seen using a spectroscope
The energy that is given off when an excited electron falls to the ground state is separated into its component wavelengths.

43. Each atom has its own distinct pattern of emission lines (bright line spectrum) and these spectra are used to identify elements.

44. Ions
A CHARGED PARTICLE. This is no longer considered an atom, it is not neutral anymore. The number of protons and electrons are NOT EQUAL!
We can never change the number of protons! So to produce a charge electrons are either gained or lost.
Atoms lose or gain electrons in order to have a complete outer shell (to become more stable), and follow the octet rule (8 valence electrons)

45. Example: Sodium
Sodium has 11 protons and 11 electrons with a ground configuration of 2-8-1
The third PEL in incomplete with only 1 electron.
To become more stable Na will lose the 1 electron in the incomplete PEL.
When Na loses 1 electron, the configuration becomes 2-8 and now has 11 protons and only 10 electrons (-10) = +1
The sodium ion is now written as Na+1
A positively charged ion is called a cation

46. Na → Na+ (loses 1 electron)

47. Example: Chlorine
Chlorine has 17protons and 11 electrons with a ground configuration of 2-8-7
The third PEL in incomplete with only 7 electrons.
To become more stable, Cl will gain 1 electron in the incomplete PEL.
When Cl gains 1 electron, the configuration becomes and now has 17 protons and only 18 electrons (-18) = -1
The chlorine ion is now written as Cl-1
A negatively charged ion is called a anion

48. Cl → Cl- (gains 1 electron)

49. Try these …..
Magnesium
Sulfur
Nitrogen
Aluminum
Carbon
Argon

50. Lewis Dot Diagrams
Diagram representing the number of valence electrons in the last principle energy level (not all the electrons). There can only be 2 electrons on each side of the symbol of the element.

51. In your own words, define an atom. What is an atom composed of
In your own words, define an atom. What is an atom composed of? (hint, hint )

53. Determine the number of protons, neutrons, and electrons in this atom of copper and krypton .

59. What were the observations of Rutherford’s Gold Foil Experiment?
What were the conclusions based on those observations?

63. Draw the Bohr Model for an atom of calcium (Ca).
Example:

71. Element
% in Nature
Atomic Mass
Copper-63
Copper-65
69.17%
30.83%
62.94 amu
64.93 amu
Uranium-235
Uranium-238
0.720%
99.280%
amu
amu
Hydrogen-1
Hydrogen-2
99.985%
0.015%
amu
amu