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=Welcome to the Chemistry-A Periodic Table wiki!= **​** =​Hey EveryONE!!!!! Listen UPPP. Chemistry is one the best subjects you could think of!!!. I know sometimes it might seem boring but you have to you your mind and think outside the box, I did.. and look how I turned out... =) wonderfuLL. So lets explore the wonders of Chemistryyy. WELCOME AND I HOPE YOU ENJOY THIS PAGE!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!=

Purpose of this wiki:

 * Allow all students in Wolosin's Chemistry A classes to access and contribute to the current unit of study (The Periodic Table).
 * Organize topics of study so that students can more effectively keep track of what they are supposed to be learning in Wolosin's Chemistry A class.

=Mendeleev's Periodic Table:= Mendeleev's Periodic Table contained more than 60 elements and was made by Mendeleev writing the properties of different elements on note cards and rearranging them until the order made sense. His periodic table was arranged in order of increasing atomic mass. However, this periodic table was created before scientists knew about the structure of atoms and therefore did not contain any information about the protons. This is not the periodic table currently in use because protons of the atoms are now taken into consideration rather than just the atomic mass he used. Mendeleev was also able to predict a few elements that had not been discovered yet by leaving gaps in the table where new elements could be placed. Mendeleev was also able to predict a few elements he names ekaboron, ekaaluminium, ekamanganese, and ekasilicon which are very close to the elements known today as scandium, gallium, technetium, and germanium. The atomic masses he predicted were very close to the actual ones.

Mendeleev's Periodic Table =The Periodic Law:= When elements are arranged in order of increasing atomic number, there is a periodic repetition of their physical and chemical properties.
 * Name || Predicted Atomic Mass || Name || Atomic Mass ||
 * Ekaboron || 44 || Scandium || 44.96 ||
 * Ekaaluminium || 68 || Gallium || 69.72 ||
 * Ekamanganese || 100 || Technetium || 98 ||
 * Ekasilicon || 72 || Germanium || 72.61 ||

=Coulomb's Law:=
 * As you go from left to right across the periodic table, the atomic number increases and the physical and chemical properties change. Properties include physical state, low or high melting points, etc.
 * In each period, each of the elements have the same number of orbitals. As you go from the 1st period down on the periodic table, the more orbitals each element has, meaning there are more electrons in each element.
 * In each group, the elements have more electrons in the outer shell as you go from the top on the group to the bottom of the group.
 * Periodic repetition is the repeating nature of the properties of the elements.

This is the equation for coloums law. F equals the force that one tiny particle exerts on another, where q1 and q2 are the charges on the particles. d is the distance between the particles in pm. Coulombs law is used to describe the electrostatic force between electric electric charges. It is this force that holds electrons in their orbitals. For this law If you increase the charge than the force will get bigger. If you increase the distance than the pull will get stronger.

Coulomb's law may be stated in scalar form as follows: //The magnitude of the electrostatic force between two point electric charges is directly proportional to the product of the magnitudes of each of the charges and inversely proportional to the square of the total distance between the two charges.// =Metals, Nonmetals, and Metalloids:= = =
 * Coulomb's law** is an equation describing the electrostatic force between electric charges. It was studied and first published in 1783 by French physicist Charles Augustin de Coulomb and was essential to the development of the theory of electromagnetism. Nevertheless, the dependence of the electric force with distance (inverse square law) had been proposed previously by Joseph Priestley[1] and the dependence with both distance and charge had been discovered, but not published, by Henry Cavendish, prior to Coulomb's works.

__Metals-__ A metal is an element, compound, or alloy characterized by high electrical conductivity.
- //This picture is an example of a Metal//
 * Metals are usually inclined to form cations through electron loss.
 * Metals are shiny and good conductors of heat and electricity. They have high densities, melting points, and boiling points.
 * Metals lose their electrons very easily
 * Metals form oxides that are basic, like in the sodium demonstration the pink water was basic
 * [[image:http://www.rsalloys.in/yahoo_site_admin/assets/images/Electrolytic_Manganese_Metal_Flakes.68131133.jpg width="203" height="124" align="center"]]

__Nonmetals-__ Used when classifying chemical elements
 * Nonmetlas are dull and bad conductors of heat and electricity. They have low densities, melting points, and boiling points.
 * Nonmetals also gain or share their electrons easily and they form oxides that are acidic



- //This picture is an example of a nonmetal// __Metalloids-__ Used when classifying chemical elements, however, a few elements with intermediate properties are referred to as **metalloids.**

 * Metalloids are elements that have properties of both metals and nonmetals.



- Th//is picture is an example of a metalloid//

There are also transition metals on the periodic table
. Transition metals are less reactive than groups one and two. They have higher melting points and densities than groups one and two, and are very strong. Transition metals are used to make some object and buildings.

=Groups of Elements:= Alkali Metals**
 * The Representative Elements
 * **Alkali metal** || **Standard Atomic Weight (u)** || **Melting Point (K)** || **Boiling Point (K)** || **Density (g·cm−3)** || **Electronegativity (Pauling)** ||
 * Lithium || 6.941 || 453 || 1615 || 0.534 || 0.98 ||
 * Sodium || 22.990 || 370 || 1156 || 0.968 || 0.93 ||
 * Potassium || 39.098 || 336 || 1032 || 0.89 || 0.82 ||
 * Rubidium || 85.468 || 312 || 961 || 1.532 || 0.82 ||
 * Caesium || 132.905 || 301 || 944 || 1.93 || 0.79 ||
 * Francium || (223) || 295 || 950 || 1.87 || 0.70 ||


 * Alkaline Earth Metals**
 * Halogens**
 * **Halogen** || **Standard Atomic Weight (u)** || **Melting Point (K)** || **Boiling Point (K)** || **Electronegativity (Pauling)** ||
 * Fluorine || 18.998 || 53.53 || 85.03 || 3.98 ||
 * Chlorine || 35.453 || 171.60 || 239.11 || 3.16 ||
 * Bromine || 79.904 || 265.80 || 332.00 || 2.96 ||
 * Iodine || 126.904 || 386.85 || 457.40 || 2.66 ||
 * Astatine || (210) || 575 || 610 (?) || 2.20 ||


 * Noble Gases**

The noble gases are found in group 18 in the periodic table which is all the way to the right. The nobel gases have low boiling points because they have weak atomic forces. The nobel gases have their outer shell full with the most amount of electrons therefore they are all stable.

Table #1 Properties of The Nobel Gases
 * ~ Property[|[11]][|[24]] ||~ [|Helium] ||~ [|Neon] ||~ [|Argon] ||~ [|Krypton] ||~ [|Xenon] ||~ [|Radon] ||
 * [|Density] (g/[|dm³]) || 0.1786 || 0.9002 || 1.7818 || 3.708 || 5.851 || 9.97 ||
 * [|Boiling point] (K) || 4.4 || 27.3 || 87.4 || 121.5 || 166.6 || 211.5 ||
 * [|Melting point] (K) || 0.95[|[25]] || 24.7 || 83.6 || 115.8 || 161.7 || 202.2 ||
 * [|Enthalpy of vaporization] (kJ/mol) || 0.08 || 1.74 || 6.52 || 9.05 || 12.65 || 18.1 ||
 * [|Solubility] in water at 20 °C (cm3/kg) || 8.61 || 10.5 || 33.6 || 59.4 || 108.1 || 230 ||
 * [|Atomic number] || 2 || 10 || 18 || 36 || 54 || 86 ||
 * [|Atomic radius] (calculated) ([|pm]) || 31 || 38 || 71 || 88 || 108 || 120 ||
 * [|Ionization energy] (kJ/mol) || 2372 || 2080 || 1520 || 1351 || 1170 || 1037 ||
 * [|Allen electronegativity][|[26]] || 4.16 || 4.79 || 3.24 || 2.97 || 2.58 || 2.60 ||

table # 2 Electron configuration of the nobel gases ( z= atomic number)
 * ~ [|Z] ||~ [|Element] ||~ [|No. of electrons/shell] ||
 * 2 || helium || 2 ||
 * 10 || neon || 2, 8 ||
 * 18 || argon || 2, 8, 8 ||
 * 36 || krypton || 2, 8, 18, 8 ||
 * 54 || xenon || 2, 8, 18, 18, 8 ||
 * 86 || radon || 2, 8, 18, 32, 18, 8 ||
 * 118 || ununoctium || 2, 8, 18, 32, 32, 18, 8 ||

=Periodic Trends:= ====This needs to be understood as the charge felt by the outer shell electrons. Since for our purposes all outer shell electrons are at roughly the same distance from the nucleus, they all experience the same core charge and are pulled towards the nucleus because of it. For example, if an element has a core charge of +1 then ALL the outer shell electrons are affected equally by this +1 charge.==== We can calculate the core charge felt by the outer shell of an atom by considering the number of protons in the nucleus and how the electrons are arranged in their subshells. This is best demonstrated through examples, and elements in the second row of the Periodic Table are the best for this purpose.
 * Core Charge**

Example : Lithium If lithium has a total of 3 protons and electrons it has an overall pull of +3. But there not on the same shell since its electron configuration is 1s2, 2s1. So as the 2 electrons on the first shell are ' blocking out' the 1 electrons on the second shell, you subtract it ( 2-1 ) and you have a +1 charge. Heres an example of Lithiums Core Charge.


 * Atomic Radius

Figure _ __: Measuring an Atomic Radius__**

Atomic radii vary in a predictable and explicable manner across the periodic table.

 * Group Trends
 * Periodic Trends
 * Ions and Atoms
 * [[image:http://www.chem.umass.edu/~botch/Chem111F04/Chapters/Ch8/AtomicRadii.jpg width="341" height="368" align="center" caption="Atomic Radius"]]

====Ion formation is the gain or loss of electrons from the outer valence level of an atom so that it acquires a negative or positive charge the value of which is determined by the number of electrons gained or lost.====
 * Ion Formation**
 * Group Trends[[image:http://www.drbateman.net/gcse2003/gcsesums/chemsums/bonding/bondin1.gif width="406" height="305" caption=" A example of Ion Formation"]][[image:webkit-fake-url://2DD01ACF-19FB-4E70-B682-CE1FF4D88D56/making_sodium_ion_c_la_784.jpg caption="making_sodium_ion_c_la_784.jpg"]]
 * Periodic Trends
 * Cations
 * Anions[[image:http://rds.yahoo.com/_ylt=A0WTb_htZFZLOhABtP2jzbkF/SIG=12c0rejk6/EXP=1264039405/**http%3A//www.atmosair.com/assets/ion-formation-diagram.jpg]]

====- Ionization Energy is most commonly used to refer to the energy required to remove the outermost electron in the atom or molecule when the gas atom or molecule is isolated in free space and is in its ground electronic state.====
 * Ionization Energy**

- The value, in kJ/mol (or formerly kcal/mol) - Should be called the "molar ionization energy" but is often just called "ionization energy"
The ionization energy, or ionization potential, is the energy required to completely remove an electron from a gaseous atom or ion. The closer and more tightly bound an electron is to the nucleus, the more difficult it will be to remove, and the higher its ionization energy will be. The first ionization energy is the energy required to remove one electron from the parent atom. The second ionization energy is the energy required to remove a second valence electron from the univalent ion to form the divalent ion, and so on. Successive ionization energies increase. The second ionization energy is always greater than the first ionization energy. Ionization energies increase moving from left to right across a period (decreasing atomic radius). Ionization energy decreases moving down a group (increasing atomic radius). Group I elements have low ionization energies because the loss of an electron forms a stable octet.
 * Group Trends
 * Periodic Trends
 * 1st, 2nd, 3rd Ionization Energy

====What is Electronegativity? it is a chemical property that describes the ability of an atom to attract electrons (or density) towards itself. An atom's electronegativity is affected by both its atomic weight and the distance that its valence electrons reside from the charged nucleus.====
 * Electronegativity**

Electronegativity cannot be directly measured and must be calculated from other atomic or molecular properties

 * Group Trends
 * Periodic Trends

Atoms with large electronegativities (such as F and O) attract the electrons in a bond better than those that have small electronegativities (such as Na and Mg). The electronegativities of the main group elements are given in the figure below.



 When the difference between the electronegativities of the elements in a compound is relatively large, the compound is best classified as //ionic//. Example: NaCl, LiF, and SrBr2 are good examples of ionic compounds. In each case, the electronegativity of the nonmetal is at least two units larger than that of the metal.


 * //NaCl// ||  ||   || //LiF// ||   ||   || //SrBr////2// ||   ||   ||
 * Cl || //EN// = 3.16 ||  || F || //EN// = 3.98 ||   || Br || //EN// = 2.96 ||   ||
 * Na || //EN// = 0.93 ||  || Li || //EN// = 0.98 ||   || Sr || //EN// = 0.95 ||   ||
 * || //[[image:http://chemed.chem.purdue.edu/genchem/topicreview/bp/ch8/graphics/delta.gif width="11" height="10" align="center"]]
 * || //[[image:http://chemed.chem.purdue.edu/genchem/topicreview/bp/ch8/graphics/delta.gif width="11" height="10" align="center"]]

EN// = 2.23 ||  ||   || //

EN// = 3.00 ||  ||   || //

EN// = 2.01 ||  ||

We can therefore assume a net transfer of electrons from the metal to the nonmetal to form positive and negative ions

=**Chemical Bonds** > > > > CHECK OUT THIS LINK ON FRANKIUM: > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > > =Chemical Bonding= Chemical compounds are formed by the joining of two or more atoms. A stable compound occurs when the total energy of the combination has lower energy than the separated atoms. The bound state implies a net attractive force between the atoms ... a chemical bond. The two extreme cases of chemical bonds are: [|Covalent bond]: bond in which one or more pairs of electrons are shared by two atoms. [|Ionic bond]: bond in which one or more electrons from one atom are removed and attached to another atom, resulting in positive and negative ions which attract each other. Other types of bonds include [|metallic bonds] and [|hydrogen bonding]. The attractive forces between molecules in a liquid can be characterized as [|van der Waals bonds]. > > > **I**onic bonds arise from elements with low electronegativity reacting with elements with high electronegativity. In this case there is a complete transfer of electrons. > > A well known example is table salt, sodium chloride. Sodium gives up its one outer shell electron completely to chlorine which needs only one electron to fill its shell. Thus, the attraction between these atoms is much like static electricity since opposite charges attract. > > = Covalent bonds = =Covalent bonds involve a complete sharing of electrons and occurrs most commonly between atoms that have partially filled outer shells or energy levels. Thus if the atoms are similar in negativity then the electrons will be shared. Carbon forms covalent bonds. The electrons are in hybrid orbitals formed by the atoms involved as in this example: ethane. Diamond is strong because it involves a vast network of covalent bonds between the carbon atoms in the diamond.=
 * A chemical bond **is an interaction between atoms or molecules and allows the formation of polyatomic chemical compounds. A chemical bond is the attraction caused by the electromagnetic force between opposing charges, either between electrons and nuclei, or as the result of a dipole attraction. The strength of bonds varies considerably; there are "strong bonds" such as covalent or ionic bonds and "weak bonds" such as dipole-dipole interactions, the London dispersion force and hydrogen bonding.**=
 * [[image:http://rds.yahoo.com/_ylt=A0WTb_5CZlZLGF4A5L.jzbkF/SIG=12nvl713o/EXP=1264039874/**http%3A//www.sd41.k12.id.us/Staff/kdunn/images/chemical%2520bonds.jpg]]
 * [[image:http://rds.yahoo.com/_ylt=A0WTb_5CZlZLGF4A5L.jzbkF/SIG=12nvl713o/EXP=1264039874/**http%3A//www.sd41.k12.id.us/Staff/kdunn/images/chemical%2520bonds.jpg]]
 * [[image:http://rds.yahoo.com/_ylt=A0WTb_5CZlZLGF4A5L.jzbkF/SIG=12nvl713o/EXP=1264039874/**http%3A//www.sd41.k12.id.us/Staff/kdunn/images/chemical%2520bonds.jpg]]
 * []




 * Periodic Table**



 There is a significant difference between the physical properties of NaCl and Cl2, as shown in the table below, which results from the difference between the ionic bonds in NaCl and the covalent bonds in Cl2. **//Some Physical Properties of NaCl and Cl//****//2//**


 * ||  || //N////aCl// ||   || //Cl// //2// ||
 * Phase at room temperature ||  || Solid ||   || Gas ||
 * Density ||  || 2.165 g/cm3 ||   || 0.003214 g/cm3 ||
 * Melting point ||  || 801°C ||   || -100.98°C ||
 * Boiling point ||  || 1413°C ||   || -34.6°C ||
 * Ability of aqueous solution to conduct electricity ||  || Conducts ||   || Does not conduct ||

//**he Difference Between Polar Bonds and Polar Molecules__**// The difference between the electronegativities of chlorine (//EN// = 3.16) and hydrogen (//EN// = 2.20) is large enough to assume that the bond in HCl is polar.

+ ||  || - ||   ||
 * [[image:http://chemed.chem.purdue.edu/genchem/topicreview/bp/ch8/graphics/deltasm.gif width="5" height="11" align="center"]]
 * H || [[image:http://chemed.chem.purdue.edu/genchem/topicreview/bp/ch8/graphics/em.gif width="22" height="9" align="center"]] || Cl ||  ||

Because it contains only this one bond, the HCl molecule can also be described as polar. The polarity of a molecule can be determined by measuring a quantity known as the **dipole moment**, which depends on two factors: (1) the magnitude of the separation of charge and (2) the distance between the negative and positive poles of the molecule. Dipole moments are reported is units of //debye// (//d//). The dipole moment for HCl is small: µ = 1.08 //d//. This can be understood by noting that the separation of charge in the HCl bond is relatively small (//EN// = 0.96) and that the H-Cl bond is relatively short. C-Cl bonds (//EN// = 0.61) are not as polar as H-Cl bonds (//EN// = 0.96), but they are significantly longer. As a result, the dipole moment for CH3Cl is about the same as HCl: µ = 1.01 //d//. At first glance, we might expect a similar dipole moment for carbon tetrachloride (CCl4), which contains four polar C-Cl bonds. The dipole moment of CCl4, however, is 0. This can be understood by considering the structure of CCl4 shown in the figure below. The individual C-Cl bonds in this molecule are polar, but the four C-Cl dipoles cancel each other. Carbon tetrachloride therefore illustrates an important point: Not all molecules that contain polar bonds have a dipole moment.