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module 2 - foundations Flashcards

1
Q

how many electrons can each orbital hold?

A

2 electrons

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2
Q

what is the order of electronic configuration?
(what order do the orbitals fill up in?)

A

1s^2, 2s^2, 2p^6, 3s^2, 3p^6, 4s^2, 3d^10, 4p^6, 5s^2

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3
Q

what is pauli’s exclusion principle?

A

electrons must have different spin to their pair. (one up, one down)

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4
Q

what is “The Aufbau process”?

A

electrons fill up the lowest energy levels first

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5
Q

what is Hund’s rule?

A

we half-fill orbitals with electrons of the same spin

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6
Q

what shape are s orbitals?
and how to they arrange more than one orbital on an atom?

A

spherical,
the first orbitals are inside the sphere of the second orbital etc.

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7
Q

what is the “Heisenberg Uncertainty Principle”?

A

you cannot know where an electron is AND its speed and direction.

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8
Q

what shape are P-orbitals
and how do they arrange with each other on an atom?

A

like an infinity symbol or dumbbell
aligned along their perpendicular axis, eg. Px= horizontal along the x orbital, same with Py and Pz

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9
Q

at each pole/ end of the p orbitals, is there a high or low chance of the electron being there?

A

high chance at the poles and in the orbital, low chance where an orbital is not apparent

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10
Q

why are the electrons shared in covalent bonding and “donated” in ionic bonding?

A

in ionic bonding, there is a great difference in electronegativity causing the electron to be more attracted to one atom than the other, so much so that it is donated
whereas in covalent bonding, there is less difference in electronegativity

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11
Q

in covalent bonding, why might electrons “favour” one of the atoms to the other (not shared equally)

A

if there is a difference in electronegativity but not enough for the electron to be “donated”

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12
Q

even though NaCl and MgO have the same ionic structure, why does MgO have a stronger ionic bond?
what does this mean for the melting point?

A

-because in MgO, 2 electrons are donated compared to 1 electron in NaCl, therefore the electrostatic forces in the giant ionic lattice are stronger
- the melting point of MgO is higher because more energy is needed to overcome the attraction holding the bond together

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13
Q

put these in order of strongest to weakest bonds.
- small ions low charges
- small ions high charges
- big ions low charges
- big ions high charges
2. why is it in this order?

A
  1. small ions high charges
  2. big ions high charges
  3. small ions low charges
  4. big ions low charges

in this order because charge has a greater impact on bonds than the size of the ions

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14
Q

why is potassium a better conductor than sodium but has a lower melting point?

A

potassium is a larger molecule so has less electronegativity than sodium, meaning the delocalised electron is “more free” to carry charge - better conductor
potassium has a lower melting point as the attraction between the nucleus and outer shell electrons are weaker meaning less energy is needed to overcome the forces holding the bonds together

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15
Q

why, if iodine has extremely strong true covalent bonds, is its melting point low?

A

even though the covalent bonds are strong, the intermolecular forces are weaker, due to a lack of attraction between the nucleus and outer shell electrons, meaning less energy is needed to overcome these intermolecular forces.

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16
Q

why does iodine dissolve in hexane

A

because they’re both non-polar
(non-polar compounds tend to be more soluble)

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17
Q

why can iodine not conduct electricity?

A

there are no delocalised electrons or ions

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18
Q

what is dative covalent bonding

A

when one atom provides both electrons for a covalent bond

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19
Q

compared to normal covalent bonds and intermolecular forces, what is the strength of dative covalent bonds

A

weaker than normal covalent but stronger than intermolecular forces

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20
Q

what are polyatomic ions

A

groups of covalently bonded atoms that carry a net charge

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21
Q

why would electrons fill up 4s1 3d10
instead of 4s2 3d9

A

its more stable this way

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22
Q

define relative atomic mass

A

the weighted mean mass of an atom relative to 1/12th the mass of one atom of carbon-12

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23
Q

define relative isotopic mass

A

the weighted mean mass of an atom of an ISOTOPE of the element compared with 1/12th the mass of an atom of carbon - 12 (which is exactly 12)

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24
Q

what instrument is used to determine the isotopic abundance of an element?

A

mass spectrometer

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25
common ionic charges: nitrate ion - carbonate - sulfate - hydroxide - ammonium - metal ions: zinc - silver -
nitrate - NO^3- carbonate CO3^2- sulfate - SO4^2- hydroxide - OH- ammonium - NH4+ zinc - Zn ^2+ silver - Ag+
26
define bond energy
the amount of energy required to break a covalent bond
27
what does the bond length tell us?
the distance between the two nuclei
28
how does the size of ions affect the bond length? and how does the bond length affect the strength of the bond?
the bigger the size of the ion, the longer the bond length the longer the bond length, the weaker the bond
29
what is the bond length and strength like in double and triple bonds compared to single bonds?
double and triple bonds are stronger and shorter
30
when can an ionic solid dissolve in water? (refer to bond energy)
If the energy released when the ions interact with the water molecule compensates for the bond energy (energy needed to break the bond holding the ion together) AND the energy needed to separate the water molecules so that the ions can be inserted into the solution
31
what acronym do we use to tell the shape of molecules and secondly, explain what theory this refers to
Valence Shell Electron Pair Repulsion Electron pairs repel each other whether they are in chemical bonds or lone pairs. Valence electron pairs are oriented to be as far apart as possible to minimize repulsions.
32
when are electron pairs bonding and when are they non-bonding
bonding - one electron from each atom non-bonding - both electrons from one atom and not bonded to another atom
33
why do electron pairs repel each other as far as possible?
like charges repel
34
what shape ion has 2 bonding pairs and what angle can the electrons be furthest away from each other
linear - 180 degrees
35
what shape has 3 bonding pairs and what angle can the electrons be furthest away from each other
trigonal planar - 120 degrees
36
what shape has 4 bonding pairs and what angle can the electrons be furthest away from each other
tetrahedral - 109.5 degrees
37
what shape has 5 bonding pairs and what angle can the electrons be furthest away from each other
trigonal bipyramidal - some have 90 degrees, some have 120 degrees
38
what shape has 6 bonding pairs and what angle can the electrons be furthest away from each other
octahedral - 90 degrees
39
how do double and triple bonds affect the shape of molecules?
they do not, they count as one bond
40
put these in order of greatest to weakest repulsion 1. two bonding pairs 2. two lone pairs 3. one bonding pair, one lone pair
greatest - two lone pairs middle - one lone one bonding weakest - two bonding pairs
41
define relative molecular mass
the average mass of a molecule relative to 1/12th the mass of carbon 12. It is the sum of the relative atomic masses of each atom in a molecule.
42
define molecular formula what is the molecular formula of glucose?
an expression that defines the number and type of atoms present in a molecule of a compound C6 H12 O6
43
how do you find the molecular formula of a compound when given percentages and the molecular molar mass. (6 steps) 1. change 2. convert 3. divide 4. multiply 5. divide 6. multiply
1. change % to grams 2. convert grams to moles - moles = mass/mr 3. divide by smallest no of moles (empirical formula) 4. multiply moles by mr to find empirical molar mass 5. divide molecular molar mass by empirical molar mass 6. multiply this number by empirical formula
44
why do lone pairs of electrons provide more repulsion?
they're more compact as they are closer to the nucleus than bonding pairs
45
what is the shape name and the angle of ammonia (3 bonding, 1 lone) (and how much is the bond angle reduced by each lone pair?)
pyramidal shape angle reduced by 2.5degrees (from 109.5 to 107 degrees) - as repelled by lone pair
46
what do we call the shape of molecules with 2 bonding pairs and 1 lone pair? and what is the bond angle?
non-linear (bent or angular) bond angle- 117.5 degrees
47
what shape are molecules with 3 bonding pairs and 1 lone pair? and what is the bond angle?
pyramidal bond angle- 107 degrees
48
what are molecules called with two bonding pairs and two lone pairs? and what is the bond angle?
non-linear (bent or angular) bond angle- 104.5 degrees
49
what shape are molecules with 4 bonding pairs and 2 lone pairs? and what is the bond angle?
square planar 90 degrees
50
how do you draw 3D diagrams of molecules on 2D paper
solid wedges - show bonds coming towards you - out of the page diagonal hash wedges - show bonds going away from you straight singular lines show bonds going across the page (parallel to you)
51
what is the shape and bond angle of a molecule with 3 bonding pairs and 1 lone pair and why is it different to a normal 4 pairs tetrahedral?
trigonal pyramidal 107degrees the lone pair repels the other electrons more, reducing the bond angle by 2.5degrees
52
what is the shape and bond angle of a molecule with 2 bonding pairs and 2 lone pairs?
bent (v-shape) 104.5degrees
53
what is the shape and bond angles of a molecule with 4 bonding and 1 lone pair?
trigonal pyramidal or (see-saw) 119degrees and 89degrees
54
what is is the name and bond angle of a molecule with 3 bonding pairs and 2 lone pairs
trigonal planar or t-shape 120 degrees or 89 degrees
55
what is the name and bond angle of a molecule with 5 bonding and 1 lone pair?
square pyramid - 89 degrees
56
what is the name and bond angle of a molecule with 4 bonding and 2 lone pairs?
square planar 90 degrees
57
why do lone pairs repel more than bond pairs?
lone pairs are more compact than bonding pairs
58
what are the 5 steps to working out shapes (single bonds only) 1. count 2. work out 3. how many 4. based on 5. take into account
1. count number of electrons on the central atom (if charged, remove or add electrons to account for the charge) 2. each atom forms one bond to the central atom, work out how many electrons are left after bonding pairs 3. how many electrons pair together? 4. what is the shape based on? (eg. 4 electron pairs are based on a tetrahedral) 5. take into account the lone pairs to work out the bond angle
59
why are some bonds polar?
there is a difference in electro negatives therefore the electron is more attracted to one atom in the covalent bond than the other. the more electronegative atom will have a slightly negative charge as the electrons are closer. permanent dipole - slightly positive side and slightly negative side of molecule
60
define "dipole"
charge separation across a bond with one atom having a slightly negative charge and the other having a slightly positive charge due to electrons being more attracted to one atom
61
where on the periodic table is the least and most electronegative?
group 1 metals are the least the lower down (more shells) the least non-metals are the most electronegative most the higher up (least shells) the most
62
which atom will have the slightly negative charge and which atom will have the slightly positive charge in a polar covalent bond
the atom with the larger negativity value will have the slightly negative charge the atom with the smallest negativity value will have the slightly positive charge
63
what electronegativity difference do 1. covalent bonds 2. polar covalent bonds 3. ionic bonds have?
1. 0 2. 0-1.8 3. >1.8
64
why is H2O polar but CO2 is non-polar even though each atoms in both compounds both have permanent dipoles?
in H2O, the two dipoles act in different directions but do not directly oppose each other in CO2 the two dipoles act in OPPOSITE directions and DIRECTLY oppose each other, causing the dipoles to cancel and the overall dipole is zero
65
how do polar substances dissolve in water?
water molecules attract the Na+ and Cl- ions, breaking down the ionic lattice, due to the polar charge of the water molecules.
66
what are the three types of intermolecular force? which is the weakest which is the strongest?
induced dipole-dipole (London dispersion forces) - weakest permanent dipole-dipole hydrogen bonding (strongest)
67
define intermolecular force
forces of attraction between molecules
68
where are intermolecular forces apparent?
ionic compounds (lattices)
69
what is a permanent dipole-dipole
the weak electrostatic attraction between polar molecules (due to difference in electronegativity and the shape of the molecule)
70
how do permanent dipoles induce dipoles in other non-polar molecules
pulling or pushing electrons across the other molecule
71
what are instantaneous dipoles? what molecules experience them?
slight charge caused by the movement of electrons around the molecule. eg. electron moves closer to the one side, giving it a negative charge for an "instant" all molecules, polar and non-polar experience them
72
what causes london dispersion forces?
when an instantaneous dipole induces a dipole in a nearby molecule
73
how will the type of, and amount, of intermolecular forces affect the property of a molecule?
melting/boiling points - the stronger, or more numerous, forces present, the higher the boiling point as the more energy required to break them solubility - "like dissolves like" - polar substances can easily dissolve in polar solvents as they can easily produce permanent dipole-dipole interactions
74
what is sublimation
when a molecule goes straight from a solid to a gas
75
why are ionic compounds with induced dipole-dipole interactions easily sublimated?
when heated at a low temperature, the weak intermolecular forces are easily broken, allowing molecules to escape the lattice causing the compound to sublimate rather than melt
76
what happens when an instantaneous dipole induces another molecule
a London dispersion force in between the molecules (not a bond) weakest form of intermolecular force
77
what three elements can form hydrogen bonds? why?
oxygen, nitrogen, fluorine because they are small atoms that have a lone pair in the 2p orbital
78
why do we use a dashed line to show hydrogen bonds?
to show that it's not a full bond, it is actually an intermolecular force of attraction between the hydrogen atom and the lone pair of electrons
79
what is special about hydrogen bonding in water?
the hydrogen bonds are the greatest as there are 2 hydrogen bonds and 2 lone pairs. therefore, every molecule forms 2 hydrogen bonds.
80
what happens to the arrangement of water molecules when frozen? why? (hydrogen bonding) what is the melting point of ice like?
forming four bonds, the molecules take a tetrahedral shape explaining why ice is strong and less dense than water (due to the molecules being held apart by hydrogen bonds) ice has a relatively high melting point due to strong hydrogen bonds
81
why is the boiling point of butanol higher than that of methylpropanol?
because butanol has stronger induced dipole-dipole interactions because it has a straight-chain structure
82
define electronegativity
the ability of an atom to gain electrons in a covalent bond
83
what is the formula to find out the maximum number of electrons in the shell
2 (n^2) where N is the number of shells
84
define ionisation energy describe the trend?
the energy required to remove one mole of electrons from one mole of gaseous atoms goes up across (left to right) of periodic table, as nucleus gets closer to outer shell electrons decreases down the group as number of electron shells increase (nucleus gets further from the outer electrons)
85
explain the trend in the attraction between the nuclei and the outermost electrons as you go across a period
going across the period, the number of electron shells (shielding) remains the same but the number of protons increases. Causing more attraction between the nucleus and outer shell electrons as the atomic radius decreases
86
what is the oxidation number?
equal to the charge of the ion
87
why is ethanol soluble in water
because they are both polar molecules, therefore the ethanol and water molecules form hydrogen bonds due to the hydroxyl group in ethanol
88
how do you work out oxidation numbers?
sum of oxidation number = total charge
89
in pure elements, what is the oxidation number
ALWAYS 0
90
define redox reaction
a reaction involving reduction (gain of electrons) and oxidation (loss of electrons)

chemistry (14 decks)

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