1. UNIT 5: THE ATOMIC STRUCTURE
How has the model of the atom evolved over time?
How can we describe the modern model of the atom?
How can we describe where are electrons are located within the electron cloud?
How can we describe how electrons are arranged?
What is electromagnetic radiation?
How is light produced?

2. AIM # 1: How can we describe an atom?
The Theory of the Atom
________________, a famous Greek teacher who lived in the 4th Century B.C., first suggested the idea of the atom.
________ __________ came up with his atomic theory based on the results of his experiments.
The Atom
The smallest particle of an ________________ is an atom.
The atom is made up of three ________________ particles.
The electron was discovered in _______ by J. J. Thomson by using a cathode ray tube. The electron has a _______ charge. It’s mass is much smaller than the other 2 subatomic particles, therefore it’s mass is usually ______________.
Democritus
John Dalton
element
subatomic
1897
(−)
ignored

3. Cathode Ray Tube

4. (2) The proton has a ______ charge, and it was discovered in _________ by E. Goldstein.
(3) The neutron does not have a charge. In other words, it is ________. It was discovered in ____ by James Chadwick. The neutron has about the same _________ as the proton.
These three particles make up all the ____________________ in the Universe!
(+)
1886
neutral
1932
mass
visible matter
 Although Goldstein did observe protons. Goldstein did not recognize these particles to be a fundamental building block of atoms, nor could have related the protons to the nucleus, since the nucleus hadn't been invented yet. This is why Rutherford is usually associated with discovering the proton.

5. Nuclear Atomic Structure
The atom is made up of 2 parts/sections:
(1) The ______________ --- (in the center of the atom)
(2) The ____________ _________ --- (surrounds the nucleus)
nucleus
electron cloud
(p+ & n0)
e− cloud

6. The Nucleus Discovered by Ernest ________________ in ________.
He shot a beam of positively charged “alpha particles”, which are ___________ nuclei, at a thin sheet of ______ _____.
Rutherford
1911
helium
gold foil
99.9% of the particles went right on through to the ______________.
Some were slightly deflected. Some even ____________ ________ towards the source!
This would be like shooting a cannon ball at a piece of tissue paper and having it bounce off.
detector
bounced back

7. Rutherford’s Experiment

8. Conclusions about the Nucleus
(1) Most of the atom is more or less _______________ (2) The nucleus is very _________. (Stadium Analogy) (3) The nucleus is very ___________. (Large Mass ÷ Small Volume) (4) The nucleus is ______________ charged.
empty space
tiny
dense
positively

11. Counting Subatomic Particles in an Atom
The atomic # of an element equals the number of ____________ in the nucleus.
The mass # of an element equals the sum of the _____________ and ______________ in the nucleus.
In a neutral atom, the # of protons = # of ______________.
To calculate the # of neutrons in the nucleus, ______________ the ___________ # from the __________ #.
protons
protons
neutrons
electrons
subtract
atomic
mass

12. Important Definitions
Nucleons: particles in the nucleus (protons and neutrons)
Nuclear Charge: charge of the nucleus (same at # of protons)
Examples:
How many nucleons are there in an atom with an atomic number of 20 and 23 neutrons?
What is the nuclear charge of an Iron atom?
Answers 1) 43 2) 26

13. Practice Problems
Find the # of e-, p+ and n0 for sodium. (mass # = 23)
Find the # of e-, p+ and n0 for uranium. (mass # = 238)
3) What is the atomic # and mass # for the following atom? # e- = 15; # n0 = 16

14. Element Name
Symbol
Atomic #
AMU
# of Protons
# of Electrons
# of Neutrons K 19 5 16 23 10 48 25 14

15. AIM #2 : How can atoms of the same element be both similar and different at the same time?- Isotopes
An isotope refers to atoms that have the same # of ___________, but they have a different # of ___________.
Because of this, they have different _________ #’s (or simply, different ___________.)
protons
neutrons
mass
masses

16. (The # shown after the name is the mass #.)
Isotopes
Isotopes are the same element, but the atoms weigh a different amount because of the # of ______________.
Examples---> (1) Carbon-12 & Carbon-13
(2) Chlorine-35 & Chlorine-37
(The # shown after the name is the mass #.)
For each example, the elements have identical ___________ #’s, (# of p+) but different _________ #’s, (# of n0).
Another way to write the isotopes in shorthand is as follows:
neutrons
atomic
mass
The top number is the ________ #, and the bottom # is the __________ number. Calculating the # n0 can be found by _____________ the #’s!
mass
atomic
subtracting 12 C 35 Cl 6 17

17. Two isotopes of sodium.

18. More Practice Problems
Find the # e-, p+ and n0 for Xe-131.
Find the # e-, p+ and n0 for
3) Write a shorthand way to represent the following isotope:
# e- = # n0 = 0 # p+ = 1 63 Cu 29

19. Aim # 3 How can we calculate the average mass of isotopes?
Based on the relative mass of Carbon-12 which is exactly _______.
1 p+ ≈ __ atomic mass unit (amu) 1 n0 ≈ __ amu 1e- ≈ __ amu
The atomic masses listed in the Periodic Table are a “weighted average” of all the isotopes of the element.
Formula: 12 1 1

20. Average Atomic Mass
Practice Problems: (1) 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? (2) Oxygen has 3 isotopes: O-16 (99.76%) O-17 (0.037%) O-18 (0.2%)

21. Average Atomic Mass
(3) Copper has an average atomic mass of amu.  It contains only two natural isotopes, which are Cu-63, with an isotope mass of and Cu-65 with an isotope mass of   What are the percent of the two isotopes in naturally occurring copper?

22. PRACTICE: Determine the average atomic mass of the following mixtures of isotopes
80% 127I, 17% 126I, 3% 128I
2. 50% 197 Au, 50% 198 Au

23. AIM# 4: How has the model of the atom evolved over time
AIM# 4: How has the model of the atom evolved over time?- Modern Atomic Theory
Atomic Models
(1) Model:
a ball of (+) charge containing a number of e-
no ________________
often described as the “________ _______________” atom.
(2) Model:
a ____________ of (+) charge surrounded by a number of e-
no _____________ and no e- orbitals
Thomson
nucleus
plum pudding
Rutherford
 Although Goldstein did observe protons. Goldstein did not recognize these particles to be a fundamental building block of atoms, nor could have related the protons to the nucleus, since the nucleus hadn't been invented yet. This is why Rutherford is usually associated with discovering the proton.
nucleus
neutrons

24. Atomic Models
(3) Model:
a nucleus of (+) charge that also contains ______________
nucleus is encircled by e-’s located in definite orbits (or paths).
e-’s have ___________ energies in these orbits
e-’s do not lose energy as they orbit the nucleus
(4) Mechanical Model ( Wave Mechanical Model)
no definite ____________ to the e- path (“fuzzy” cloud)
orbits of e-’s based on the _________________ of finding the e- in the particular orbital shape.
Bohr
neutrons
fixed
Quantum
shape
probability

25. Bohr Atomic Model

26. Bohr Atomic Model

28. Quantum Mechanical Model

29. Quantum Mechanical Model

30. Aim # 5 What do we call an atom that is not neutral?
An atom can gain or lose electrons to become electrically charged.
Cation = (___) charged atom created by ___________ e-’s.
Cations are ______________ than the original atom.
_____________ generally form cations. (****MELPS)
Anion = (___) charged atom created by _____________ e-’s.
Anions are ____________ than the original atom.
_______________ generally form anions.
Practice Problems: Count the # of protons & electrons in each ion.
a) Mg p+ = _____ e− = ______
b) F− p+ = _____ e− = ______ +
losing
smaller
Metals −
gaining
larger
Nonmetals

31. AIM #6 : How can we describe where electrons are located
AIM #6 : How can we describe where electrons are located? - Energy Levels
Diagram:
The energy levels in an atom are sort of like _________ of a ladder.
The more energy an electron has, the __________ away from the nucleus it usually will be.
The energy levels are not evenly spaced. They get ___________ together as you travel farther away.
To move from one “rung” to another requires a “____________” of energy.
rungs
farther
closer
quantum

33. continuous energy levels quantized energy levels
Figure 11.15: The difference between continuous and quantized energy levels.
continuous energy levels
quantized energy levels

34. Quantum Numbers
Describe the ______________ of the e-’s around the nucleus.
Quantum #’s are sort of like a home _____________ for the electron.
This information about the location of the e-’s in an atom can be used to:
(1) determine chemical & physical _____________ for the elements.
(2) show how the _______________ __________ is organized.
(3) show _____ and _____ elements combine to form compounds.
location
address
properties
Periodic Table
how
why

35. Electron Configurations
Practice Problems:
Write the electron configuration notation for each of the following
atoms: H C Fe Br Ne

36. Electron Configurations & Properties
How do electron configurations relate to the chemical and physical properties of an element?
All elements with the _________ outer shell e- configurations have ________ properties.
This means that elements in the same ____________ group have similar properties.
Examples: (1) Li, Na, K, Rb, and Cs all have __ lone e- for their last orbital This makes all of them ___________ reactive. They all react with __________ to produce hydrogen gas.
(2) Ne, Ar, Kr, Xe, and Rn all have the outer energy level completely __________ with electrons... This makes all of them ______________. They do not produce ______________!
same
similar
vertical 1
very
water
filled
inert
compounds

37. Aim #7 How Light is Produced? Excited vs. Ground State
When atoms get hit with energy (by _____________ them with electricity or by ____________ them up), the electrons absorb this energy and __________ to a higher energy level. Figure (a)
As they immediately fall back down to the “____________ state”, they give off this energy in the form of a particle of ___________ (or other types of electromagnetic radiation) called a _____________. Figure (b)
zapping
heating
jump
ground
light
photon

38. How Light is Produced
Each photon emitted has a specific ___________ (or frequency).
The color of the light that is given off depends on how _____ the electron _______ (which depends on how big of a jump it originally made.) The farther the fall, the ___________ energy the photon has.
color
far
fell
greater

39. Figure 11.6: Photons of red and blue light.

40. How Light is Produced energy
Since electrons are located only in certain __________ levels (or orbitals) around the nucleus, only certain specific _________ of light are emitted.
Scientists use a _________________ to separate these colors into bands of light. These bands of color look like a ______ code of color which is characteristic of that element. No two elements produce the same ______________ of colors. This can be used to distinguish one element from another contained in a sample. (See Fig )
color
spectroscope
bar
spectrum

41. Emission Spectrum
Hydrogen Spectrum
Neon Spectrum

42. SUMMARY https://www.youtube.com/watch?v=FSyAehMdpyI
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