I want this.
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My muse
Only with this cliched vapid little moron, do I feel like the genius I was bound to become.
Pitbull ft Ne-Yo
(via ivanakuntin)
oh I am lol
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(Source: staypozitive)
(Source: staypozitive)
What book is this from? Classic!
Irish phrases anyone?
Chemical Reactions
Reactants and products - New substances are formed with chemical reactions. The chemical substances which are changed are called the reactants and the new substances formed are called the products.
Collision Theory - For a chemical reaction to happen, the reactants must collide. The collision between the molecules provides the energy needed to break the bonds so that new bonds can form. The molecules must collide in a particular way, the place on the molecule where the collision must take place is called the reactive site.
Sometimes a collision may occur but not enough kinetic energy is available to be transfered. Raising the temperature can increase kinetic energy enabling bonds to become broken during collisions. The energy you supply to get a reaction going is called the activation energy.
Exothermic Reactions: Releasing heat - In this type of reaction, heat is released when you go from reactants to products. The reactants are in a higher energy state than the products.
Endothermic reactions: Absorbing heat - In this type of reaction, heat is absorbed. The reactants are in a lower energy state than the products.
Different types of chemical reactions occur based on; the identity of the reactants and products and which bonds are broken and made. Some common reactions include - combination, decomposition, single displacement, double displacement and reduction-oxidation (redox).
Combination reactions - two or more reactants form one product. e.g sodium and chlorine to form sodium chloride.
Decomposition reactions - Here a single compound breaks down into two or more simpler substances. These reactions are opposite of combination reactions. e.g the decomposition of water into hydrogen and oxygen gases.
Single Displacement reactions - A more active element displaces another less active element from a compound.e.g Putting zinc metal into copper sulfate solution, the zinc displaces the copper.
Double displacement reactions - Here two species (normally ions) are displaced. Usually an insoluble solid or water will be formed.
Things to remember; 1) If a compound is soluble, it will not react at all and can be represented by (Aq). 2) If a compound is insoluble, it will precipitate (form a solid).
Neautralization reactions - A reaction between an acid and a base occurs forming water.
Combustion reactions - When a compound (usually containing carbob) reacts with oxygen gas in air resulting in burning.
Redox reactions - reactions where electrons are exchanged. These types of reactions are found in combustion, photosynthesis and respiration.
Balancing chemical equations - the law of conservation of mass states matter is neither created nor destroyed. i.e atoms in an equation cannot be gained or lost - just combined differently. The same number should be present on both sides of the equation, it should be balanced.
Its a good idea to wait until the end to balance hydrogen and oxygen atoms.
Example 1 - N2(g) + H2(g) → NH3(g)
Lets start with the nitrogen atoms. There are 2 on the left and 1 on the right. In order to balance them, place a 2 infront of the NH on the left.
N2(g) + H2(g) → 2NH3(g)
For the hydrogen atoms, there are 2 on the left and 6 on the right (2x3). Put a 3 infront of the H2 on the left. which gives N2(g) + 3H2(g) → NH3(g)
Bonding 2
Covalent Bonds - hydrogen has 1 electron in its orbit. It wishes to gain another one so its 1s can be filled. If two hydrogen atoms meet and the first hydrogen atom gave its one electron away to the second hydrogen atom. The second one would be complete (H-) but the first atom would be unstable with 0 electrons. As a result, they can share electrons known as covalent bonding. It is now H2, a diatomic (two atom) compound.
Other diatomic forms include - oxygen, nitrogen, fluorine, chlorine, bromine, iodine.
Differences between the two;
Ionic - Metal and Non-metal Covalent - Two non-metals
Ionic - Usually solid at room temperature Covalent - Can be solid, liquid or gas
Ionic - Melting point is higher than colvalent Covalent - Melting point is lower
Ionic - Tend to be electrolytes Covalent - Tend to be non-electrolytes.
Multiple Bonds - In diatomic molecules, only one electron is shared. However in other covalent bonds - the atoms share more than one electron pair.
Nitrogen (N2) is in the VA family meaning it has 5 outer valence electrons and needs three more to complete its octet (8 electrons). A nitrogen atom can fill its octet by sharing three electrons with another nitrogen atom - forming 3 covalent bonds. e.g N (with 5 in its outer shell) plus N (with 5 in its outer shell) = :N:::N:
Co2 is another example of this. Carbon has four valence electrons and oxygen has 6. Carbon can give away two of its valence electrons to each oxygen atoms to form Co2 (a double bond).
Structual Formulas - the molecular formula used in ionic compounds is enough to identify the compounds but not enough to identify covalent compounds. For instance C2H6O could be more than one compound. The only way to identify it is to look at how the atoms are bonded.
Compounds with the same molecular formula but different structures are called isomers to each other.
Basic Bonds - 1)Find the central Atom.The central atom is usually the one which forms more than one covalent bond when trying to fill their valence energy level. e.g oxygen, nitrogen, phosphorus, sulfur, silicon, carbon.
In the case of H2O, oxygen is the central atom and the two hydrogen atoms are bonded to it.
2)Count the valence electrons. each hydrogen has 1 electron and oxygen has 6 valence electrons = 8 electrons. a)Work out how many valence electrons are needed. (usually 8, unless its hydrogen its 2)The number of valence electrons available. e.g boron trifluoride (BF3). Boron - 5 (3 valence electrons), Flouorine - 9 (7 valence electrons). 9x3 as there are 3 fluorine particles = 21. 21+ 3 = 24 valence electrons total.
so for H2O; a)8 valenece plus 2 for each of the hydrogen atoms = 12 b) 6 valence for oxygen plus 1 for each of the hydrogens = 8.
c)12-8 = 4. 4 bonds shared in water. 4/2 = 2 which means there are two bonds.
Distribute four electrons to account for the bonds. Then use the remaining electrons to distribute around the other atoms.
Electronegativites -Atoms may share electrons through bonds but that doesn’t mean they share equally. Electronegativity is the strength an atom has to attact a bonding pair of electrons to itself. Electronegativities increase from left to right in a period and decrease top to bottom in a family.
A bond where an electron share is equally paired is known as a nonpolar covalent bond. A bond where the electron pair is shifted toward one atom is known as polar covalent bond.
The atom which attracts the bonding electron pair is slightly more negative.e.g in hydrogen chloride (Hcl) - hydrogen had an electonegativity of 2.1 and chlorine has an electronegativity of 3.0. Because chlorine has a larger value, the electron pair bondng Hcl shifts toward the chlorine atoms. If the atoms have very different electronegativities - they are likely to form ionic bonds instead of covalent.
Guide to types of bonds - Electronegativity differences - If its between 0.0-0.2 = nonpolar covalent. Between 0.3 and 1.4 = polar covalent. Anything over 1.5 is likely to be ionic.
Polar covalent molecules can act as weak electrolytes because the bond allows the substance to act as a conductor.
Polar molecules are dipoles (one end having a partial negative charge and the other end having a partial positive charge). Attraction and force between molecules is known as intermolecular force. There are three different types of force;
London force - a weak force occuring between non covalent molecules.
Dipole-Dipole interaction - when the positive end of one dipole is attracted to the negative end of another dipole.
Hydrogen bond - An extremely strong dipole-dipole interaction when hydrogen is bonded to one of the three extremely electromagnetic elements -O, N or F.
