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topic 2 Flashcards

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1
Q

ionic bonding

A

strong electrostatic forces of attraction between oppositely charged ions

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2
Q

what makes ionic bonding stornger

A

it is stronger and has higher mp when the ions are smaller and/or have higher charges

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3
Q

ionic radii - positive ions

A

ions are smaller compared to their atoms as there is one less shell of electrons. Electrons are held more closely together

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4
Q

ionic radii - negative ions

A
  • ions formed from group 5 to 7 are larger than their atoms. Neg ions has more electrons than its atom but same number of protons. Pull of nucleus is shared over more electrons and the attraction is less, making the ion bigger
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5
Q

general trend in ionic radii down a group

A

increases as you go down - more shells of electrons

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6
Q

physical properties of ionic compounds

A
  • high melting points (strong forces of attraction between ions)
    • non-conductors of electricity when solid (ions held tightly and cant move)
    • conduct electricity when molten or in solution
    • brittle
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7
Q

covalent bonding

A

electrostatic forces of attraction between the shared pair of electrons and the two nuclei

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8
Q

Why do diamond and graphite have high mp and bp

A

contain many strong covalent bonds in a macromolecular structure - requires a lot of energy to break the many strong bonds

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9
Q

effect of multiple bonds on bond strength and length

A

double and triple bonds have a greater electron density between them - results in greater force of attraction between nuclei and electrons between them - results in shorted bond length and greater bond strength

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10
Q

dative covalent bonding

A

when the shared pair of electrons come from one of the bonding atoms

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11
Q

linear

A

2 bp
0 lp
180°

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12
Q

trigonal planar

A

3 bp
0 lp
120°

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13
Q

tetrahedral

A

4 bp
0 lp
109.5°

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14
Q

trigonal pyramidal

A

3 bp
1 lp
107°n

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15
Q

bent

A

2 bp
2 lp
104.5°

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16
Q

trigonal bipyramidal

A

5 bp
0 lp
120° and 90°

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17
Q

octahedral

A

6 bp
0 lp
90°

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18
Q

how to explain shape

A
  • state number of lone pairs and bonding pairs
  • state electron pairs try to repel and get as far apart as possible
  • if no lone pairs state all bonding pairs repel equally
  • if there are lp, state lp repel more than bp
  • state shape and bond angle
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19
Q

what repels more - lone pair or bonding pair

A

lone pairs repel more than bonding pairs so reduce bond angle by 2.5°

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20
Q

electronegativity

A

the relative tendancy of an atom in a covalent bond in a molecule to attract electrons in a covalent bond to itself

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21
Q

factors that effect electronegativity

A
  • electronegativity increases across a period as proton number increases and atomic radius decreases because electrons in the same shell are pulled in more
  • electronegativity decreases down a group because distance between nucleus and outer electrons increases and shielding increases
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22
Q

what makes a compound purely covalent

A

if it contains elements of similar electronegativity hence a small electronegativity difference

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23
Q

what bonds are elements with large electronegativity differences

A

ionic
if electronegativity is greater than 1.7

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24
Q

formation of permanent dipole/polar covalent bond

Study These Flashcards
A

forms when the elements in the bond have different electronegativities (approx. 0.3 to 1.7)p

25
polar covalent bond
has an unequal distribution of electrons in the bond and produces charge separation
26
symmetrical molecules
- all bonds are identical and there are no lone pairs - are not polar
27
eg of symmetrical molecule
CO2
28
London forces
occur between ALL molecular substances and noble gases. They do not occur in ionic substances much weaker than covalent bonds
29
how are induced dipoles formed
temporary dipoles cause dipoles to form in neighbouring molecules. Induced dipoles are always the opposite sign to the original one
30
main factor affecting size of London forces
the more electrons there are, the higher the chance of temporary dipoles forming. Makes London forces stronger between the molecules and more energy is needed to break them - shape of molecule also effects size of London forces
31
affect on London force as molecules get larger
London forces increase
32
permanent dipole-dipole forces
- occur between polar molecules - stronger than London forces so compounds have higher boiling points - polar molecules have permanent dipole - they are asymmetrical
33
hydrogen bonding
occurs in compounds that have one hydrogen atom attached to one of the three most electronegative atoms of nitrogen, oxygen and fluorine
34
what forms hydrogen bonds
alcohols - high bp and low volatility compared to alkanes with similar number of electrons water - can form 2 hydrogen bonds per molecule (has 2 lp) ice - molecules are held further apart by the hydrogen bonds than in water, hence why ice is dense
35
why is hydrogen bonding stronger than the other two types of intermolecular bonding
hydrogen bonding occurs additionally to London forces
36
ionic substances dissolving in water
bonds in lattice break and form new bonds between the metal ions and water molecules. Negative ion are attracted to the positive hydrogen on the polar water molecules and positively charged are attracted to negative oxygen molecules
37
solubility of simple alcohols
- smaller alcohols are soluble in water because they can form hydrogen bonds with water - the longer the hydrocarbon the less soluble the alcohol
38
solubility in non-aqeous solvents
- compounds which have similar intermolecular forces to the solvent will dissolve
39
metallic bonding
electrostatic force of attraction between positive metal ions and delocalised electrons
40
3 main factors that affect metallic bonding
1. number of protons more protons = stronger 2. number of delocalised electrons per atom more = stronger 3. size of ion smaller = stonger
41
why do metals have high melting points
because the strong electrostatic forces between positive ions and sea of delocalised electrons requires a lot of energy to break
42
why do metals conduct electricity well
because the delocalised electrons can move through the structure
43
why are metals malleable
positive ions in the lattice are all identical. Ions can slide easily over one another
44
giant lattices present in
ionic solids covalently bonded solids (diamond, graphite) solid metals
45
can diamond conduct electricity
no because all 4 electrons per carbon are involved in covalent bonds - they are localised and cannot move
46
structure of diamond
tetrahedral 4 covalent bonds per atom
47
can graphite conduct electricity
yes because one electron per carbon is free and delocalised, so electrons can move easily along layers
48
structure of graphite
planar 3 covalent bonds per atom 4th electron is delocalised
49
why do both macromolecular structures have high melting points
due to strong covalent bonds in the giant structure so it takes a lot of energy to break the many strong covalent bonds
50
graphene structure
one layer of graphite - 3 covalent bonds per atom and 4th delocalised electrons
51
can graphene conduct electricity
yes because one electron is delocalised
52
why is graphene very strong
it has high tensile strength because of the strong structure of covalent bonds
53
can carbon nanotubes conduct electricity
yes cos one delocalised electron per carbon
54
use of carbon nanotubes
potentially use as vehicles to deliver drugs to cells
55
solubility of ionic substances
- ionic substances dissolve in water due to the hydration of the ions - the ions break up the bonds in the lattice and form new bonds between the metal ions and water molecules
56
57
insolubility of compounds in water
- compounds that cannot form hydrogen bonds with water molecules, e.g. polar molecules such as halogenoalkanes or non-polar substances will be insoluble in water
58
what are giant lattices present in
- ionic solids - covalently bonded solids (diamond, graphite & silicon oxide) - solid metals

chemistry (28 decks)

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