GIVE ME TEN POINTS!!!
HERES UR GROUPS....
Group 0 Elements
The elements in this group have a full outer electronic shell and thus, the elements in this group have no tendency to lose, gain or share electrons. Thus, the elements in the group are chemically inert.
All the elements in this group are gaseous. Because of their chemical inertness, the elements in this group are called the Nobel Gases :
Helium
Neon
Argon
Krypton
Xenon
Radon
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Group I Elements
The elements in this group have one electron in their outer electronic shell. Thus, each element in this group has a tendency to lose a single electron, to form a singly charged positive ion, which has the stable electronic configuration of its neighbouring Nobel Gas element in the periodic table.
The elements in the group are chemically reactive.
The group is divided into two sub-groups.
Sub-Group Ia : The Alkali Metals
Lithium
Sodium
Potassium
Rubidium
Caesium
Francium
Sub-Group Ib : Transition Metal Elements
Copper
Silver
Gold
Hydrogen is included in this group because it has a single electron in its outer electronic shell. However, hydrogen has none of the metallic properties of the alkali metals.
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Group II Elements
The elements in this group have two electrons in their outer electronic shell. Thus, each element in this group has a tendency to lose two electron, to form a doubly charged positive ion, which has the stable electronic configuration of its neighbouring Nobel Gas element in the periodic table.
The elements in the group are chemically reactive.
Sub-Group IIa : The Alkaline Earth Metals
Beryllium
Magnesium
Calcium
Strontium
Barium
Radium
Sub-Group IIb : Transition Metal Elements
Zinc
Cadmium
Mercury
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Group III Elements
The elements in this group have three electrons in their outer-most electronic shell. The electronic configurations of the n th orbitals (i.e. outer-most orbitals) are ns2 np1. Thus, each element in this group has a tendency to lose three electron, to form a triply charged positive ion, which has the stable electronic configuration of its nearest neighbouring Nobel Gas element in the periodic table.
The elements in the group are chemically reactive.
Sub-Group IIIa : Transition Metal Elements
Scandium
Yttrium
Lanthanum
The following elements also have the same electronic configuration as lanthanum (i.e. 4s2 4p1) in their outer-most electronic orbitals, while the inner 3d orbitals are being filled, from going from element to element.
Cerium
Praeseodymium
Neodymium
Promethium
Samarium
Europium
Gadolinium
Terbium
Dysprosium
Holmium
Erbium
Thulium
Ytterbium
Lutecium
These are the transition elements, and are also called the Lanthanides.
Actinium
The following elements also have the same electronic configuration as actinium (i.e. 5s2 5p1) in their outer-most electronic orbitals, while the inner 4d orbitals are being filled, on going from element to element.
Neptunium
Plutonium
Americium
Curium
Berkelium
These are the inner transition elements, and are also called the Actinides.
Sub-Group IIIb : Main Group Elements
Boron
Aluminium
Gallium
Indium
Thallium
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Group IV Elements
The elements in this group have four electrons in their outer electronic shell. Thus, each element in this group has a tendency to share these four electrons, to form covalent compounds, thereby gaining the stable electronic configuration of its neighbouring Nobel Gas element in the periodic table.
The elements in the group are chemically reactive.
Sub-Group IVa : Transition Metal Elements
Titanium
Zirconium
Hafnium
Thorium
Sub-Group IVb : Main Group Elements
Carbon
Silicon
Germanium
Tin
Lead
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Group V Elements
The elements in this group have five electrons in their outer electronic shell. Thus, each element in this group has a tendency to gain three two electron, to form a triple charged negative ion, which has the stable electronic configuration of its nearest neighbouring Nobel Gas element in the periodic table.
The elements in the group are chemically reactive.
Sub-Group Va : Transition Metal Elements
Vanadium
Niobium
Tantalum
Protactinium
Sub-Group Vb : Main Group Elements
Nitrogen
Phosphorus
Arsenic
Antimony
Bismuth
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Group VI Elements
The elements in this group have six electrons in their outer electronic shell. Thus, each element in this group has a tendency to gain two electron, to form a doubly charged negative ion, which has the stable electronic configuration of its nearest neighbouring Nobel Gas element in the periodic table.
The elements in the group are chemically reactive.
Sub-Group VIa : Transition Metal Elements
Chromium
Molybdenum
Tungsten
Uranium
Sub-Group VIb : Main Group Elements
Oxygen
Sulphur
Selenium
Tellurium
Polonium
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Group VII Elements
The elements in this group have seven electrons in their outer electronic shell. Thus, each element in this group has a tendency to lose an electron, to form a singly charged negative ion, which has the stable electronic configuration of its nearest neighbouring Nobel Gas element in the periodic table.
The elements in the group are chemically reactive.
Sub-Group VIIa : Transition Metal Elements
Manganese
Technetium
Rhenium
Sub-Group VIIb : Main Group Elements
Fluorine
Chlorine
Bromine
Iodine
Astatine
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Group VIII Elements
This group contains three triads of elements, in the center of the periodic table.
Iron
Cobalt
Nickel
Ruthenium
Rhodium
Palladium
Osmium
Iridium
Platinum
These elements have the typical properties of metals, metallic luster, tensile strength, and rightly.
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Group Number
The group number is the number assigned to the vertical columns of the structured list of all known elements in the periodic table. Elements within the same group have the same number of electrons in their outer electron shells. Thus, all elements in the same group have similar chemical properties.
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Groups in the Periodic Table
The elements are arranged in the sequence of their increasing atomic numbers into the periodic table, which is arranged in rows and columns, so that elements with similar chemical properties are in the same vertical column.
The elements which are in the same columns are said to be in the same group, and they have similar chemical properties.
HERES UR TRENDS
Note: These are general periodic trends of elements. There are many exceptions to these general rules.
Review
Period - a row of elements on the periodic table. Remember that sentences are written in rows and end with a period.
Group - a column of elements on the periodic table. Remember that group is spelled group and groups go up and down.
Atomic Radius - Atomic radius is simply the radius of the atom, an indication of the atom's volume.
Period - atomic radius decreases as you go from left to right across a period.
Why? Stronger attractive forces in atoms (as you go from left to right) between the opposite charges in the nucleus and electron cloud cause the atom to be 'sucked' together a little tighter.
Group - atomic radius increases as you go down a group.
Why? There is a significant jump in the size of the nucleus (protons + neutrons) each time you move from period to period down a group. Additionally, new energy levels of elections clouds are added to the atom as you move from period to period down a group, making the each atom significantly more massive, both is mass and volume.
Electronegativity - Electronegativity is an atom's 'desire' to grab another atom's electrons.
Period - electronegativity increases as you go from left to right across a period.
Why? Elements on the left of the period table have 1 -2 valence electrons and would rather give those few valence electrons away (to achieve the octet in a lower energy level) than grab another atom's electrons. As a result, they have low electronegativity. Elements on the right side of the period table only need a few electrons to complete the octet, so they have strong desire to grab another atom's electrons.
Group - electronegativity decreases as you go down a group.
Why? Elements near the top of the period table have few electrons to begin with; every electron is a big deal. They have a stronger desire to acquire more electrons. Elements near the bottom of the chart have so many electrons that loosing or acquiring an electron is not as big a deal. This is due to the shielding affect where electrons in lower energy levels shield the positive charge of the nucleus from outer electrons resulting in those outer electrons not being as tightly bound to the atom.
Ionization Energy - Ionization energy is the amount of energy required to remove the outmost electron. It is closely related to electronegativity.
Period - ionization energy increases as you go from left to right across a period.
Why? Elements on the right of the chart want to take others atom's electron (not given them up) because they are close to achieving the octet. The means it will require more energy to remove the outer most electron. Elements on the left of the chart would prefer to give up their electrons so it is easy to remove them, requiring less energy (low ionization energy).
Group - ionization energy decreases as you go down a group.
Why? The shielding affect makes it easier to remove the outer most electrons from those atoms that have many electrons (those near the bottom of the chart).
Reactivity - Reactivity refers to how likely or vigorously an atom is to react with other substances. This is usually determined by how easily electrons can be removed (ionization energy) and how badly they want to take other atom's electrons (electronegativity) because it is the transfer/interaction of electrons that is the basis of chemical reactions.
Metals
Period - reactivity decreases as you go from left to right across a period.
Group - reactivity increases as you go down a group
Why? The farther to the left and down the periodic chart you go, the easier it is for electrons to be given or taken away, resulting in higher reactivity.
Non-metals
Period - reactivity increases as you go from the left to the right across a period.
Group - reactivity decreases as you go down the group.
Why? The farther right and up you go on the periodic table, the higher the electronegativity, resulting in a more vigorous exchange of electron.
Ionic Radius vs. Atomic Radius
Metals - the atomic radius of a metal is generally larger than the ionic radius of the same element.
Why? Generally, metals loose electrons to achieve the octet. This creates a larger positive charge in the nucleus than the negative charge in the electron cloud, causing the electron cloud to be drawn a little closer to the nucleus as an ion.
Non-metals - the atomic radius of a non-metal is generally smaller than the ionic radius of the same element.
Why? Generally, non-metals loose electrons to achieve the octet. This creates a larger negative charge in the electron cloud than positive charge in the nucleus, causing the electron cloud to 'puff out' a little bit as an ion.
Melting Point
Metals - the melting point for metals generally decreases as you go down a group.
Non-metals - the melting point for non-metals generally increases as you go down a group.
HERES UR TYPES OF BONDS
Types of Chemical Bonds:
The forces of attraction that hold atoms together are called chemical bonds. The following is a list of different types of chemical bonds, each having a mini-tutorial and associated animation.
Intramolecular Bonds:
Intramolecular bonds refers to the forces of attraction that hold atoms together within a molecule. These types of bonds are considerably stronger than intermolecular bonds. We will study three types of intramolecular bonds - covalent, ionic and metallic bonds.
Covalent Bonds:The animation above is a schematic depiction of what happens in the formation of a covalent bond. The individual atoms are atoms of chlorine with only their valence electrons shown. Note that each chlorine atom has only seven valence electrons, but really wants eight. When each chlorine atom shares its unpaired electron, both atoms are tricked into thinking each has a full valence of eight electrons. Notice that the individual atoms have full freedom from each other, but once the bond is formed, energy is released, and the new chlorine molecule (Cl2) behaves as a single particle.
A covalent bond is typically formed by two non-metals. Non-metals have similar electronegativities. Consequently, neither atom is "strong" enough to steal electrons from the other.Therefore, the atoms must share the electrons.
*A covalent bond exists when two electrons are shared by two non-metallic atoms.*
Ionic Bonds:The animation below schematically shows the process that takes place during the formation of an ionic bond. The individual atoms are sodium and chlorine with only their valence electrons shown. Note that chlorine has seven valence electrons (it wants a full shell of eight), and that sodium has one valence electron (it also wants a full shell of eight). Because chlorine has a high electronegativity (3.5) compared to sodium (0.9), chlorine can easily steal sodium's one lonely valence electron. This makes chlorine very happy, as it now has eight valence electrons. This also makes sodium very happy, as it also now has eight valence electrons. Where did these eight valence electrons come from? They were already a part of the sodium atom, in the 2nd Principle Energy Level (hidden from your view in this animation to simplify matters). The one electron it lost was in the 3rd Principle Energy Level.
The transfer of the electron caused the previously neutral sodium atom to become a positively charged ion (a cation), and the previously neutral chlorine atom to become a negatively charged ion (an anion). The attraction for the cation and the anion is called the ionic bond.
An ionic bond is typically formed between a metal and a non-metal. Metals have low electronegativities (less than 2.0), while non-metals have high electronegativities (above 2.0). Consequently, the non-metal is "stronger" than the metal, and can steal electrons very easily from the metal. This results in the metal becoming a cation, and the non-metal becoming an anion.
*An ionic bond is the resulting attraction for an anion and a cation after an electron is transferred from the metal to the non-metal.*
Metallic Bonding: "SEA OF MOBILE VALENCE ELECTRONS":The animation you are looking at attempts to help you understand the nature of a metal bond. The gray spheres represent metal cations (positively charged ions), and the red moving spheres represent electrons. Metals have low ionization energies, thus they do not have a tight hold on their valence electrons. These outer electrons easily move around, as they do not "belong" to any one atom, but are part of the whole metal crystal. The negatively charged electrons act as a "cement" that hold the positively charged metal ions in their relatively fixed positions.
The fact that the electrons flow easily helps to explain some of the characteristics of metals:
- Metals are good conductors of heat and electricity. This is directly due to the mobility of the electrons.
- The "cement" effect of the electrons determines the hardness of the metal. Some metals are harder than others; the strength of the "cement" varies from metal to metal.
- Metals are lustrous. This is due to the uniform way that the valence electrons of the metal absorb and re-emit light energy.
- Metals are malleable (can be flattened) and ductile (can be drawn into wires) because of the way the metal cations and electrons can "flow" around each other, without breaking the crystal structure.
*Metallic bonds are best characterized by the phrase "a sea of electrons"*
Intermolecular Bonds:
Intermolecular bonds refers to the forces of attraction that hold molecules together. These bonds are considerably weaker than intramolecular bonds. We will study three types of intermolecular bonds - hydrogen bonds, van der Waals forces, and molecule-ion attractions.
Hydrogen Bonds: The animation above schematically shows how polar water molecules align themselves as liquid water crystallizes (freezes) into the familiar hexagonal shape of snowflakes. This crystalline structure is due to hydrogen bonding. A hydrogen bond exists between two highly polar molecules containing hydrogen. Hydrogen has a relatively low electronegativity (2.2), and when it is covalently bonded to an atom of either fluorine, oxygen, or nitrogen (electronegativities of 4.0, 3.5, and 3.1, respectively), the resulting bonds are highly polar. So when two such molecules are nearby, they will orient themselves so that the partially negative end of one molecule will face the partially positive end of another molecule.
*The attraction of the partially positive end of one highly polar molecule for the partially negative end of another highly polar molecule is called a hydrogen bond.*
van der Waals Forces: The animation above schematically shows how van der Waals forces (also called London Dispersion Forces) are induced. van der Waals forces are found in non-polar molecules, such as hydrogen gas (H2), carbon dioxide (CO2), nitrogen (N2), and in the noble gases (He, Ne, Ar, Kr, etc).
Watch the neon atom in the upper right corner of the animation screen. Notice that its electrons (blue) are randomly moving, and in one particular moment, the electrons get piled up on one side of the atom. This creates a temporary polarity in the atom (we'll call it atom #1). A temporary polarity in the neighboring atom (we'll call it atom #2) is also then induced when its electrons are repelled by atom #1. The temporarily induced polarity allows the two atoms to be attracted to each other very weakly, when the partially negative end of atom #1 is attracted to the partially positive end of atom #2.
At normal conditions of temperature and pressure, van der Waals forces are negligible. As the pressure increases (molecules are forced closer together), and/or the temperature decreases (molecules slow down), van der Waals forces are much more important. Only at conditions of high pressure and/or low temperature are the molecules able to participate in van der Waals forces noticeably---only at these conditions will these gases liquefy. What holds non-polar molecules together in the liquid state? You got it! van der Waals forces!
Another important concept to understand about can der Waals forces is this - the bigger the molecule, the stronger the van der Waals. Why? Bigger molecules have more electrons, which results in bigger electron pile-ups, which results in a bigger induced polarity. Of the noble gases He, Ne, Ar, and Kr, which do you think is the easiest to liquefy? The one that has the strongest van der Waals forces....which will be the biggest on that list...krypton.
*van der Waals forces are weak attractive forces that hold non-polar molecules together.*
Molecule-Ion Attractions: The animation above schematically shows how an ionic solid (indicated by the green and gray ion spheres) gets dissolved in water (shown by the red and yellow molecules). The polar water molecule orients itself so that its partially positive end faces the negative ion in the ionic solid. The molecule-ion attraction is stronger than the ions' attraction for each other, and the anion can be removed from the crystal lattice. Once dislodged, the anion is surrounded by more water molecules, all oriented with the positive end facing the anion. The animation continues on, to show the similar process where a cation is dissolved from the ionic solid. This time, however, it is the partially negative end of the water molecule that faces the cation.
*This "molecule-ion attraction" is appropriately named, as a molecule (water) is attracted to an ion.*