TOPIC 2 - Chemical Bonding and Structure

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  • Created by: Go
  • Created on: 25-10-15 13:48
What are the typical 5 properties of metals?
1. High melting temperatures. 2. Good electrical conductivity. 3. Good thermal conductivity. 4. Malleability. 5. Ductility
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Do metals have high or low ionisation energies?
Low
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The electrical conductivity of of a metal generally increases as the number of _____ _____ _________ increases.
Outer-shell electrons
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What is metallic bonding?
The strong electrostatic attraction between the nuclei of metal cations and delocalised electrons.
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In order to melt a metal what do you need to overcome? And why? (3)
The electrostatic forces of attraction (1) between the nuclei of the cations and the delocalised electrons (1) so that the cations are free to move arond the structure. (1)
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Do group 1 metals have a higher or lower melting temperature than group 2 metals?
Lower
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Do metals in the d-block typically have high melting temperatures, + why?
Yes, they typically have high melting temperatures bc they have more delocalisd electrons per ion.
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The smaller the cation = the closer the delocalised electrons to the nucleus of the cation. Does this cause an increase or decrease in the forces of attraction?
Increase
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What 2 factors contribute to the abiliy of metals to transfer heat energy?
1. The free-moving delocalised electrons passing KE along the metal. 2. The closely packed cations passing KE from one cation to another.
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What is ionic bonding confned to? (3)
Solid materials (1) consisting of a regular array of oppositely charged ions (1) extending throughout a giant lattice network (1).
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How are the ions in an ionic bond arranged?
They're arranged in such a that the electrostatic attractions between the the oppositely charged ions are greater than the electrostatic repulsions between ions with the same charge.
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For ions of the same charge the smaller the ions the MORE/LESS energy is required to overcome the electrostatic interactions between the ions and seperate them.
More
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Define: Isoelectronic
Atoms with the same number of electrons + so have the same electronic configuaration.
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What are the 4 physical properties of ionic compunds?
1. High melting temperatures. 2. Brittleness. 3. Poor electrical conductivity when solid but good when molten. 4. Often soluble in water.
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Why do ionic compounds have high melting temperatures? (Mention 3 things)
1) They consist of a giant lattice network of oppositely charged ions. 2) There are many ions in the lattice, so comined electrostatic forces among all the ions is large 3) A large amount of energy is required to overcome the forces of attracton.
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Why are ionic compounds brittle? (Mention 2 things)
1) If a stress is applied to an ionic solid than the layers of ions slide over each other. 2) Ions of the same charge are now side by side + they repel each other
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Why do solid ionic compounds NOT conduct electricity? (2)
No delocalised electrons + no ions that are free to move.
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Why will molten ionic compounds conduct electricity?
The ions are mobile
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Define: First Ionisation Energy
The energy required to remove 1 mol of electrons from 1 mol of atoms in gaseous state.
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With successive ionisation energies, why do big increases in energy occur between electron shells? (2)
The shell closer to the nucleus will have greater attractions for electrons (1) + less shielding from complete shells (1)
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Ionisation energy trends can be explained by what 3 things?
1. The distance from the nucleus. 2. The nuclear charge / attraction. 3. The amount of shielding.
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Why does the 1st ionisation energy always decrease going down the group, even for the increase in nuclear charge? (2)
The electrons is further from the nucleus (1) + had more shielding (1).
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Going across the period...why is there an overall increase in ionisation energy?
Due to the increase in nuclear charge for the same distance from the nucleus.
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Going across the period...why is there a drop from Be to B?
Due to shielding from the full 2s orbital
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Going across the period...why is there a drop from N to O?
Due to electron replsion when p electrons pair up (easier to remove).
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Define: Ionic Bond
The attraction between ions.
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Are the strongest ionic bonds (most reactive elements) formed between the top or bottom of group 1 and 2?
Bottom
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Are the strongest ionic bonds (most reactive elements) formed between the top or bottom of group 6 and 7?
Top
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Do cations become slightly smaller or bigger than their neutral atoms?
Slightly Smaller
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Do anions become slightly smaller or bigger than their anions?
Slightly Bigger
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Define: Covalent Bond
The sharing of a pair of electrons.
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Why must the electron pair lie between the nuclei in a covalent bond?
So the attraction forces outweigh the repulsion between the n.
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What is a covalent bond formed by?
The overlap of 2 atomic orbitals each containing a single electron.
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Both an end-on overlap of 2 s-orbitals and an end-on overlap of 2 p-orbitals leads to the formation of...
A sigma bond
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A sideways overlap of 2 p-orbitals (pi bond) leads to the formation of...
A pi bond.
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A sigma bond leads to the formation of...
A covalent bond between 2 atoms.
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Pi bonds cannot form until a _____ bond has been formed.
Sigma
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Why do pi bonds only exist between atoms that are joined by double or triple bonds?
Bc pi bonds can only form until a sigma bond has been formed.
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When a covalent bond forms...where is the highest electron density?
Between the 2 nuclei.
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Does a pi bond result in a high or low electron density both above and below the molecule?
A high electron density.
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In ethene, which is weaker, the pi or the sigma bond?
The pi bond.
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Why do alkenes have an increased reactivity compared with alkanes + why can alkenes readily undergo additional reactions.
Bc the pi bond in ethene is weaker than the pi bond.
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How many sigma and pi bonds is the triple bond in the nitrogen molecule made up of?
1 sigma bond and 2 pi bonds.
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Define: Bond Length
The distanace between nuclei of the 2 atoms that are covalently bonded together.
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Define: Bond Stength
The amount of energy required to break 1 mol of the bond in gaseous state.
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What is the general relationship between bond length + bond strength for bonds that are of similar nature?
The shorter the bond = the greater the bond strength.
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Explain why the shorter bond you have, the greater the bond strength.
This is a result of an increase in electrostatic attraction between the 2 nuclei + the electrons in the overlapping atomic orbitals.
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To form an ammonium ion by dative covalent bonding, only the hydrogen's _______ is transferred.
Nucleus.
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Electronegativity of elements in general INCREASES / DECREASES down a group of the periodic table.
Decreases
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Electronegativity of elements in general INCREASES / DECEASES from left to right across the period.
Increases
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Define: Electronegtivity
The ability of an atom to attract a bonding pair of electrons.
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What happens if 2 atoms of equal electronegativity bond together?
They both hve the same tendency to attract the bonding pair of electrons + so it will be found on average half way between the 2 atoms.
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Define: Pure Covalent Bond
Where the electrons are shared evenly between the 2 atoms.
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What happens if electron B is slighly more electronegaative than electron A.
That mwans that the B end ofthe bond has a higher electron density + so becomes slightly negative. The A end becomes slightly positive.
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Define: Polar Covalent Bond
A type of covalent bond between 2 atoms where the bonding electrons are unequally distributed. -SO- 1 atom carries a slight negative charge + the other a slight positive charge.
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What happens if electron B is a lot more electronegative than electron A.
The electron pair is dragged right over to B's end of the bond. Electron A has lost control of its electron, electron B has complete control of both electrons. Ions have been formed.
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The electronegativity of an atom depends on what 3 things?
1. The no. of protons in the nucleus 2. The distance from the nucleus 3. The amount of shielding of inner electrons.
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Which one is true... "the shorter distance = the higher electronegativity" OR "the bigger distance = the higher electronegativity".
The shorter distance = the higher electronegativity.
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Which is one is true... "more energy levels = higher electronegativity" OR "fewer energy levels = higher electronegativity".
Fewer energy levels = higher electronegativity.
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Why are symmetrical molecules not polar?
Because the dipoles cancel out.
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How are instantaneuous dipoles produced on non-polar molecules?
Due to the random movements of electrons - at any moment in time a very slight dipole can be created.
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What are the names of the 3 type of van der waal forces?
1. Permanent Dipole - Permanent Dipole Attraction 2. Permanent Dipole - Induced Dipole Attraction 3. Instantaneuous Dipole - Induced Dipole Attraction.
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Where does permanent dipole - permanent dipole attraction occur?
Between polar molecules.
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Where does permanent dipole - induced dipole attraction occur? And how does this occur?
Between a polar molecule and a non-polar molecule, whereby the polar molecule induces a dipole on the non-polar molecules.
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Where does an instantaneuous dipole -induced dipole attraction occur? And how does this occur?
Between 2 non-polar molecules, whereby a dipole is instantaneuously produced on one; inducing a dipole on the other.
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Card 2

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Do metals have high or low ionisation energies?

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Low

Card 3

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The electrical conductivity of of a metal generally increases as the number of _____ _____ _________ increases.

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Preview of the front of card 3

Card 4

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What is metallic bonding?

Back

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Card 5

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In order to melt a metal what do you need to overcome? And why? (3)

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