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Chapter 11 (Modern Atomic Theory) parts 1-6 review
11.1 Rutherford’s Atom
Section Bullets
•
Rutherford was the man that bombarded the metal foil with alpha particles.
•
He is the person that recognized that an atom has a nucleus composed of positive
particles, protons, and neutral particles, neutrons.
•
Rutherford found that the nucleus is very small compared to the rest of the atom.
•
The electrons account for the rest of the atom.
•
He could not explain why the electrons are orbiting the nucleus and why they are
not attracted to it.
11.2 Energy and Light
Vocabulary:
•
Electromagnetic Radiation: the scientific term for light
•
Wavelength: (symbolized by the Greek letter Lambda) The distance between two
consecutive wave peaks. The top of one peak to the top of another
•
Frequency: (symbolized by the Greek letter nu) indicates how many wave peaks
pass a certain point per given time period.
•
Photon- A tiny packet of energy/light
Section Bullets
•
Energy is transmitted from one place to another via electromagnetic radiation
(like a light bulb being warm, and burning wood being hot)
•
Different types of electromagnetic radiation (from longest to shortest in
wavelength)
o Radio
o Micro
o Infrared
o Visible light (white light)
o Ultraviolet
o X
o Gamma
•
Electromagnetic radiation is transmitted through waves
•
3 properties of a wave:
o Wavelength –check vocabulary
o Frequency- check vocabulary
o Speed – indicates how fast a given peak travels through water
•
Electromagnetic radiation can have the characteristics of a certain particle- photon
•
Different photons carry different amounts of energy.
o A red photon is less energized than a blue photon
o Based on the wavelength of the electromagnetic radiation the photon is
carrying
11.3 Emission of Energy by Atoms
Section Bullets
•
Different elements’ energy changes correspond to different photons
o Lithium energy changes correspond with red photons
o Copper energy changes correspond with green photons
o Sodium energy changes correspond with yellow/orange photons
•
When atoms get energized, they want to release that excess energy, so they
release the energy in the form of a photon.
•
The smaller the wavelength of light in the photon, the more energized the atom is
11.4 The Energy Levels of Hydrogen
Vocabulary:
•
To Quantize: to only allow certain values
Section Bullets
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Ground state- lowest possible energy state
•
When a hydrogen atom absorbs energy from an outside source, it uses that energy
to enter an excited state
•
It goes back to the ground state by releasing the energy
•
Only certain types of photons are produced when you have a large sample of
energized hydrogen atoms.
•
Because only certain photons are emitted, we know that only certain energy
changes are occurring.
•
Hydrogen has certain discrete energy levels, therefore, they are quantized
•
Energy levels of all atoms are quantized
11.5 The Bohr Model of the Atom
Section Bullets
•
Bohr constructed a model of the hydrogen atom with quantized energy levels
•
He pictured electrons in specific circular orbits corresponding to different energy
levels
•
It was at first promising, but it did not work with most other atoms
•
It is learned that most electron orbitals are not orbiting a nucleus like planets
around a sun
11.6 The Wave Mechanical Model of the Atom
Vocabulary:
•
Wave Mechanical Model: Bohr’s model except without two dimensional,
circular orbitals. (It contains the s, p, d, etc. orbitals), also makes electrons waves
Section Bullets
•
Louis Victor de Broglie and Erwin Schrodinger suggested that because light
seems to have both wave and barticle characteristics, the electron might also
exhibit both characteristics.
•
Although everyone had assumed that the electron was a tiny particle, these
scientists said that it might be a wave
•
Schrodinger invented the Wave Mechanical Model, where orbitals are nothing
like orbits.
•
Schrodinger ran experiments with fireflies as being similar to electrons and
created probability maps. (figures 11. 18 and 19)
For practice problems, look at the problems at the end of the chapter that are under the
same sections that this guide covers. Also, you might want to look at the figures up to
11.19 to use as reference.
Chapter 11 Part 2- 11.7-11.11
Harris Gurny
Orbital- Location and movement of electrons
The probability of finding an electron decreases as you get further away from the
nucleus.
Electrons have separate energy levels called principal energy levels. These are divided
into sublevels. The number of sublevels increases as the principal energy level increases.
In order to label orbital’s,
1. The number tells you the principal energy level.
2. 2. The letter tells the shape of the orbital. The letter s means a spherical orbital;
the letter p is a two-lobed orbital. S has one orbital, while p has who, and s is
always there first sublevel in a principal energy level, while p is always the
second. After s and p comes d and f.
The s sublevel has one orbital. The p sublevel contains 3 orbitals. The d sublevel has 5,
and the f orbital has 7.
Pauli exclusion principle is when an atomic orbital can hold a max of 2 electrons (with
opposite spins)
This means that each orbital holds two electrons. S orbital hold 2, a p holds 6, d
holds 10, and f holds 14.
Principal Components of the Wave Mechanical Model of the atom
1. Atoms have a series of energy levels called principal energy levels, which are
designated by whole numbers symbolized by n; n can equal 1,2,3,4… Level 1
corresponds to n=1, level 2 corresponds to n=2, and so on.
2. The energy of the level increases as the value of n increases.
3. Each principal energy level contains one or more types of orbitals, called
sublevels.
4. The number of sublevels present in a given principal energy level equals n. For
example, level 1 contains one sublevel (1s); level 2 contains two sublevels (two
types of orbital’s), the 2s orbital and the three 2p orbital’s; and so on.
5. The n value is always used to label the orbitals of a given principal level and ids
followed by the letter that indicates the type of the orbital. For example, the
designation 3p means an orbital in level 3 that has two lobes.
6. An orbital can be empty or it can contain one or two electrons, but never more
than tow. They must have opposite spins.
7. The shape only indicates the probability distribution.
4d
5s
4p
3d
4s
3p
3s
2p
2s
1s
Valence electrons- the electrons in the outermost principal energy level of an atom, core
electrons are the inner atoms.
The atoms of elements in the same vertical group have the same number of electrons in a
given type of orbital.
Metals are found at the left and center of the periodic table. The most chemically active
metals are found in the lower left-hand corner. The most chemically active nonmetals are
located in the upper-right hand corner.
For the representative elements, atomic size increases going down a group but decreases
going from the left to right across a period.
Ionization energy, the energy required to remove an electron from a gaseous atom,
decreases going down a group and increases going from left to right across a period.
Questions:
1. What is the electron configuration for phosphorus?
2. What is the number of unpaired electrons in an oxygen atom?
3. What is the maximum number of electrons allowed in a p orbitals?
4. What is the maximum number of electrons allowed in the 4th energy level?
5. Order the elements S, Cl, and F in terms of increasing ionization energy.
Magnesium vs sodium: who has a larger…
6. Atomic size
7. Number of valence electrons
8.ionization energy
9. Which has the smallest atomic radius (F, N, S, Br, Cl)?
10. Sodium has how many electrons in its outermost principal energy level?
Answers:
1. P=1s^2 2s^2 2p^6 3s^2 3p^3
2. 2
3. 2
4. 32 (not 18)
5. S, Cl, F
6. Sodium has a larger atomic size than magnesium.
7. Mg has more
8. Sodium’s is lower
9. F
10. 1
Reference:
Zumdal, Steven S. Zumdal, Susan L. DeCoste, Donald J. World of Chemistry. Boston,
MA: Houghton Mifflin Company, 2002, Print