Mr.+Jeff+Astor's+Textbook

My textbook will take you through every step of this year's chemistry curriculum.

__Unit 1 Theme: Death by Element__
flat **__Lesson 1.1 Summary:__** media type="custom" key="23784802" Our Friend the Atom Parts of the Atom: There are three subatomic particles that exist within the atom:
 * 1) //** Protons **//
 * 2) //** Neutrons **//
 * 3) //** Electrons **//

Protons Neutrons Electrons
 * Protons are positively charged subatomic protons that exist within the nucleus of the atom
 * They weigh 1 amu.
 * They are attracted to other subatomic particles (neutrons) by the strong nuclear force).
 * Particles that have a neutral charge that exist within the nucleus of the atom.
 * They weigh 1 amu
 * Negatively charged subatomic particles that exist outside the nucleus of the atom in the electron cloud
 * Weight Virtually nothing

**__Lesson 1.2 Summary:__**
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**Dead White Guys and the Atom** Models of the Atom Daltons's, Thomson's, and Rutherford's model of the atom
 * [[image:astronautschemclass/Models of the atom align="left"]]

There are three central players to the development of the atom. They are J. J. Thomson, Ernest Rutherford, and Niels Bohr.

**J. J. Thomson's Experiment**

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 * Attached a battery to the cathode tube and saw a beam form. He concluded that the beam must contain subatomic particles because the cathode tube was a vacuum. He put a positive magnet next to the beam and it bent towards the magnet. Put a negative magnet next to the beam and it bent away from the magnet

Thomson had discovered the electron! The beam had to have been negatively charged because negative charges are attracted to positive magnets and repelled by negative magnets. //**Therefore, the beam contained negatively charged subatomic particles that he called electrons.**//

Thomson did NOT discover the location of these electrons, rather he merely discovered the existence of them. ||   || **Ernest Rutherford** media type="custom" key="23579908" Ernest Rutherford was a student of J. J. Thomson and he provided additional evidence to further develop the view of the atom. He provided additional evidence using the gold foil experiment.

Rutherford took positively charged alpha particles and shot them at a piece of gold foil that was as thick as tissue paper. Most particles went straight through the foil. Small amount of the particles were deflected, but still went through. An even smaller amount were fired back at the original source.

Most of the alpha particles went straight through because the atom is mostly empty space. Some alpha particles showed slight deflections because they came close to the negatively charged electrons. **//The few alpha particles that bounced back were coming into contact with the very small, but very dense positively charged nucleus.//**

Niels Bohr
media type="custom" key="23579930" Scientists knew that if you shot electricity through a tube of hydrogen gas, a bright color would be produced. If you looked through a special lens, you knew that only a few colors were being emitted.

Bohr knew that shooting electricity through the hydrogen excited the electrons. A color was produced when electrons returned to original state. There was a problem through: if electrons were free to roam, then we should get all sorts of colors, but we only get those four colors.

He did not solve this using an experiment. He solved it using only math. **//Electrons are not free to roam in the electron cloud, electrons are restricted to orbits or energy levels.//**

__ **Lesson 1.3 Summary**: __
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__//**Essential Point**//__: The periodic table is split into __**groups**__ (columns) and __**periods**__ (rows) tells us a few things about every element:
 * 1) __**Atomic Number**__
 * 2) __**Atomic Mass**__
 * 3) Element Name
 * 4) Element Symbol

• **__Atomic Number __** – The number of  protons in the nucleus of an atom.



__•Atomic Mass - __ The mass of an atom: # of protons + # of neutrons.

The element name and symbol are listed as well. The element symbol always starts with a capital letter, and will be either 1 or 2 letters. The 2nd letter of an element symbol is always lower case. 



__**Lesson 1.4 Summary:**__
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__//**Essential Point**//__: On the periodic table, all elements are presented in their neutral state: the number of protons equals the number of electrons. Because of this, the atomic number also tells us the number of electrons!



Knowing this helps us draw out what the atom looks like. We can use Niels Bohr's idea that electrons do not roam freely, but instead exist in specific energy levels called orbitals to draw out what an atom of any element should look like.

We know that the first orbital can only hold 2 electrons and every subsequent orbital holds 8 electrons at most. We also know that electrons don't fill up a new orbital until the previous one has been completely filled. Any scientist can look at the atomic number, determine the number of electrons in the atom, and fill up the orbitals accordingly. The ones on the very most outside orbital are called __**valence electrons**__.

media type="file" key="1.4.Lithium Bohr Electron Configuration Video.mp4" width="300" height="300"

Valence electrons play a role in every interaction between elements. The __**octet rule**__ is one of the most important governing guidelines in chemistry and states that atoms will gain or lose electrons in order to complete their outer octet and have 8 valence electrons (hence, the octet rule). If the outer shell is the first electron orbit, then having 2 electrons will satisfy the octet rule, because the outer orbital is full.

media type="file" key="8.31.2013.Bohr Electron Diagram for Magnesium after it completes the octet rule.mp4" width="300" height="300"

To find out how many valence electrons an element has, scientists only need to look at what group the element resides in.



Because all elements in the same group have the same number of valence electrons, they react very very similarly.

The only way to really satisfy the octet rule is to either gain or lose electrons from/to another atom. When electrons are gained or lost, the atom is no longer in its neutral state, giving the atom a charge. Any atom with an unequal number of protons and electrons is known as an //**__ ion __**//, and has either a positive or negative charge. Positive ions are known as __**cations**__ and negative ions are known as __**anions**__.

media type="file" key="8.31.2013.Bohr Diagram for Oxygen Ion.mp4" width="300" height="300"

Since all elements in the same group have the same number of valence electrons, they all form the same ions!



__**Lesson 1.5 Summary:**__
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__//**Essential Point**//__: As mentioned before, it's important to remember that all elements in the same group have the same number of valence electrons and therefore exhibit similar reactivities.

We name these groups specific titles based on their characteristics!

__Group 1: The Alkali Metals__
 * Group 1/1A elements
 * Have 1 valence electron, so they're very likely to give away 1 electron to satisfy the octet rule.
 * //__Highly__// reactive with water

Elements: Lithium, Sodium, Potassium, Rubidium, Francium, Caesium

__Group 2: The Alkaline Earth Metals__
 * Group 2/2A elements
 * Have 2 valence electrons, so they're very likely to give away 2 electrons to satisfy the octet rule.
 * Highly reactive, but not as much as the alkali metals

Elements: Beryllium, Magnesium, Calcium, Strontium, Barium, Radium

__Groups 3-12: Transition Metals__
 * Have all of the similar properties of metals: conductive, malleable, solid at room temperature, etc. etc.

__Group 17: Halogens__
 * Group 7/7A elements
 * Has 7 valence electrons, so these are very likely to gain 1 to satisfy the octet rule.
 * The most reactive group, because they only need 1 electron to be satisfied.

Elements: Fluorine, Chlorine, Bromium, Iodine, Astatine

__Group 18: The Noble Gases__
 * Group 18/18A
 * Has 8 valence electrons
 * Non-reactive, because it doesn't need any more or less valence electrons to be satisfied

Elements: Helium, Neon, Argon, Krypton, Xenon, Radon

The Periodic table also separates elements into clusters according to their form. The three types are metals, non-metals and semimetals (metalloids).

__Metals:__
 * Shiny and Malleable
 * Solid at room temperature
 * Metals are great __**conductors**__ of heat and electricity as well

__Non-metals:__
 * Dull and not malleable[[image:astronautschemclass/iodine.png width="206" height="90" align="right"]]
 * Gases at room temperature
 * Very poor __**conductors**__ of heat and electricity, and are therefore known as __**insulators**__

__Semimetals:__
 * [[image:astronautschemclass/silicon semimetal.png width="247" height="84" align="right" caption="SIM cards are made of silicon!"]]Characteristics of both metals and non-metals
 * Look metallic, but are brittle
 * Neither good nor bad conductors of heat and electricity, so they're known as semiconductors.

The periodic table can actually tell us exactly which elements are metals, non-metals, or semimetals by finding the staircase.

On the left side of the "staircase" or semi-metal line are the metals of the periodic table and on the right side are non-metals. Any element touching the staircase on two sides (except aluminum) is a semimetal.

**//__ The one exception: ____ Hydrogen __//**
Hydrogen is located on the left side of the periodic table, but is still a non-metal



__**Lesson 1.6 Summary:**__
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Ionization Energy: || The negatively charged electrons are attracted towards the positively charged nucleus. This attraction plays a major role in the periodic trends that are observed for elements. Nuclear attraction is responsible for all the trends that are observed for element.
 * [[image:http://chemistryhungergames.weebly.com/uploads/8/8/4/9/8849208/2069571.gif?332 caption="Picture" link="http://chemistryhungergames.weebly.com/uploads/8/8/4/9/8849208/2069571_orig.gif?332"]]

Ionization energy is the energy that is required to remove an electron from an atom. This is the energy that is required to turn our neutral atom into an ion. This is different depending on the element. As you move down a group the ionization energy decreases. As you move across a period on the Periodic Table, the ionization energy increases.

As you go down a group more electron orbits are added. Ionization energy decreases as you go down a group because it requires much less energy to remove a valence electron due to the shielding.

As you move across a period you add more protons and electrons within the same electron orbit. The larger amount of protons in the nucleus and the larger amount of electrons show an increased attraction. results in it being much harder to pull away electrons as you add more protons within a period and a higher ionization energy as you go across a period. ||



__**Lesson 1.7 Summary:**__
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We can categorize atoms based on their ability to attract electron Electronegativity is the measure of the ability of an atom to attract electrons. We can also call this the strength of the atom. The priniciples of nuclear attraction that are observed in ionization energy are also observed in this periodic trend.
 * Electronegativity:

As you go down a group, the electronegativity of an element decreases. As you go across a period, the electronegativity of an element increases.

As you go down a group more electron orbits are added. Electronegativity decreases because there is a decreased ability of the positively charged nucleus to attract valence electrons as a result of this larger distance and a decrease in electronegativity as you go down a group.

As you move across a period you add more protons and electrons within the same electron orbit. The larger amount of protons in the nucleus and the larger amount of electrons show an increased attraction. This results in a much larger ability for the positively charged nucleus to attract valence electrons within a period and an increase in electronegativity as you go across a period. || || || The atomic radius is the distance from the center of the nucleus to the outermost edge of the electron cloud for an atom.
 * Atomic Radius:

Atomic radius increases as you go down a group on the Periodic Table. Atomic radius decreases as you go across a period on the Periodic Table.

As you add more electron shells to an element, the element becomes “bulkier” due to the increased amount of electron orbits. This means that as you go down a group, more orbits are added, so the atomic radius becomes bigger.

As you move across a period, more protons are added to the nucleus. Also, more electrons are added within the same electron orbit. This means there is a larger positive and negative charge, which results in a higher attraction and the atomic radius to decrease. ||