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52 Cards in this Set
- Front
- Back
2 Bonded Pairs |
Linear eg. BeCl2 180° |
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3 Bonded Pairs |
Trigonal Planar eg. BF3 120° |
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4 Bonded Pairs |
Tetrahedral eg. CH4 109.5° |
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5 Bonded Pairs |
Trigonal Bi-pyramidal eg. PCl5 90°&120° |
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6 Bonded Pairs |
Octahedral eg. SF6 90° |
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Electronegativity |
Tendency to attract a bonding pair of electrons in a covalent bond |
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Pauling Scale Trends |
-Left to Right = Increase in Electronegativity -Non-metals have higher Electronegativity -Top to Bottom = Decrease in Electronegativity |
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Factors Affecting Electronegativity |
- Nuclear Charge - More Protons = Stronger attraction between nucleus and bonding pair of electrons. - Atomic Radius - Closer to nucleus = stronger attraction between nucleus and bonding pair. - Shielding - Less shells of electrons between nucleus and electrons = Less shielding = Stronger attraction between nucleus and bonding pair. |
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Pure Covalent Bond |
No charge as they share electrons equally |
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Polar Covalent Bond |
Low partial charge as they share electrons unequally |
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Ionic Bond |
High full charge as there is a loss or gain of electrons |
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How do lone pairs affect shapes of molecules? |
Repulsion of the lone pairs is stronger then the bonded pairs and so -2.5 from angle |
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Ionic Bond |
Strong electrostatic attraction between a non-metal and a metal, two oppositely charged ions. |
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Covalent Bond |
Strong chemical bond between two non-metals includes sharing of electrons. |
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Giant Ionic Lattice |
-'Giant' because made up of same basic unit repeated over again -Forms because each ion is electrostatically attracted in all directions ions of opposite charge - Sodium Chloride packed in lattice and is cube shaped. |
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Physical Properties Of Ionic Bonds |
-High melting point so ions held by strong attraction. -Often soluble in water but not in non-polar solvents so particles are charged as ions are pulled apart by polar molecules. -Bad conductor of electricity when solid but good when molten or dissolved so ions are in fixed position when solid but free to move and carry charge when liquid. -Not malleable as repulsion between ions would be strong so ionic compounds break easily. |
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Trend In Ionic Radii |
-Ionic Radius decreases as you go down a group because electron shells are added. -Isoelectronic Ions are ions of different atoms with same number of electrons, Ionic Radius of a set of isoelectronic ions decreases as atomic number increases. |
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Patterns In Solubility |
-Highly polar solids such as ionic salts dissolve in water (polar solvent) but not in a hexane (non polar solvent). -Polar organic substances such as sucrose dissolve in water but not in hexane. -Non polar solids such as candle wax do not dissolve in water but dissolve in hexane. -Non polar liquids such as petrol mix completely. -Polar liquid and Non polar liquid such as water and petrol do not mix. |
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Ionic Solid In Water (IONIC SALTS) |
Anions and Cations held together by strong electrostatic forces and the ionic lattice may only be broken by using energy, this energy is called lattice energy. |
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Alcohol In Water (ETHANOL) |
Soluble as they have polar OH groups that can hydrogen bond to -OH groups in H2O. Solubility decreases with increasing carbon chain length of alcohol, this is due to a smaller proportion being polar. Carboxylic acid dissolves in water similarly |
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Non Polar Substances In Water (HEXANE) |
Not soluble and so floats on separate layer, this is because the forces between hexane molecules are weaker than the hydrogen bonds between water molecules and so hexane is unable to disturb the structure of water. |
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Non Polar Liquids (HEXANE AND OCTANE) |
Two non polar liquids that mix completely as both pure liquids contain weak van de waals forces. When mixed, the forces are extended. |
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Alkanes |
Tetrahedral structure |
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Alcohols |
H-C-H 109.5° C-O-H 104° |
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Alkenes |
More than one tetrahedral structure |
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Carboxylic Acid |
-COOH is planar Rest is tetrahedral carbon atoms Methanoic Acid is planar |
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Halogenoalkanes |
4 bonding pairs -Tetrahedral |
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Isoelectronic |
Same numbers of electrons or the same electronic structure
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3 Main Intermolecular Forces |
-Hydrogen Bonding -Van der Waals -Dipole-Dipole Interactions |
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Hydrogen Bonds |
Strongest bond. Must be covalently bonded to electronegative atom, limited to oxygen, nitrogen, fluorine, in order to hydrogen bond to molecules. Atom must have a lone pair of electrons. |
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Dipole-Dipole Interactions |
Molecules with permanent dipole have one end with slightly negative charge and the other slightly positive. Opposite ends will therefore attract one another. |
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Van Der Waals (INSTANTANEOUS DIPOLE-INDUCED DIPOLE FORCES) |
Weakest bond. Concentration of negative charge on one side can induce a dipole in molecule beside, as electrons are repelled. This creates electrostatic attraction between two molecules. Even though they have short life, they can result in chain reaction. The bigger the molecule, the greater the strength of the force as there are more electrons. |
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Why do branched chains have lower boiling points? |
Branched alkanes have weaker van der waals forces due to smaller area of contact between molecules. Less energy is needed to break the intermolecular force.Linear alkanes can line up more closer and therefore have stronger van der waals forces than unbranched alkanes. |
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Trends that affect Boiling Point |
-The relative strength of IM force: Hydrogen bonding > dipole dipole > Van der Waals dispersion forces. The influence of each of these attractive forces will depend on the functional groups present. -Boiling points increase as the number of carbons is increased. -Branching decreases boiling point |
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Whats an intra-molecular force? |
Covalent bond holding a molecule together. |
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What happens to the intermolecular when a substance melts or boils? |
Intermolecular Forces broken |
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What happens to the intermolecular when a substance condenses? |
Intermolecular Forces formed |
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Chlorine, Cl2 |
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Hydrogen Chloride, HCl |
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Oxygen, O2 |
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Water, H2O |
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Methane, C2H4 |
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Nitrogen, N2 |
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Carbon Monoxide, CO |
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Bond Enthalpy is related to Bond Length how? |
The greater the bond enthalpy, the shorter the length of the bond. |
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Bond Length |
Distance between two nuclei where attractive and repulsive forces balance each other out |
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Use Electron Pairs to predict shape of molecules? |
1. Find central atom. 2. Work out no. of electrons in outer shell. 3. Molecular formula tells us how many atoms the central atom bonded to. Work out how many electrons are shared. 4. Add up electrons and divide by 2. 5. Compare electron pair no. with no. of bonds to find no. of lone pairs. 6. Use no. of electron pairs and no. of lone pairs and bonding centres around central atom to work out shape. |
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What is lone pair of electrons? |
Electrons are not shared with another atom. |
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Why is water not linear? |
Due to presence of lone pair giving it a slight angle. |
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Properties Of Giant Covalent Bond |
-High melting point due to strong covalent bonds. -Variable conductivity, Diamond does not, Graphite does (free electrons), Silicone is semi-conductive. -Hard -Insoluble due to being attracted to neighbouring molecules in lattice more. |
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Fullerene |
Not giant covalent but giant molecular, carbon atoms form pentagons and hexagons |
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Giant Metallic Structure |
-Malleable and Ductile -Melting points considerably high -Good thermal conductors as delocalised electrons can pass on kinetic energy. -Insoluble due to strength |