Module 2.2 electrons, bonding and structure Flashcards
1
Q
Principal quantum numbers
A
The numbers given to shells
2
Q
shells further away from the nucleus have
A
higher energy (larger principal quantum number) than shells closer to the nucleus
3
Q
types of sub shell with number of orbitals and maximum electron capacity
A
s sub shell. 1 orbital. max e- = 2
p sub shell. 3 orbitals. max e- –> 3 x 2 = 6
d sub shell 5 orbitals. max e- –> 5 x 2 = 10
f sub shell 7 orbitals max e- –> 7 x 2 = 14
4
Q
how many electrons can one orbital carry
A
2
5
Q
shell, what sub-shells it contains and total no,. of electrons
A
1st: 1s total e- = 2
2nd: 2s 2p total e- : 2+(3x2) = 8
3rd: 3s 3p 3d total e- : 2 + (3x2)+ (5x2) = 18
4th: 4s 4p 4d 4f total e-: 2+ (3x2) + (5x2) + (7x2) = 32
6
Q
1st shell: sub shells and total number of electrons
A
1st
sub-shell: 1s
total e- : 2
7
Q
2nd shell: sub shells and total number of electrons
A
2nd
sub-shells: 2s 2p
total e- : 2 + (3x2) = 8
8
Q
3rd shell: sub shells and total no. of electrons
A
3rd
sub-shells: 3s 3p 3d
total e- : 2 + (3x2) + (5x2) = 18
9
Q
4th shell: sub shells and total no. of electrons
A
4th
subshells: 4s 4p 4d 4f
total e- : 2 + (3x2) + (5x2) + (7x2) = 32
10
Q
total no. of electrons in 1st shell
A
2
11
Q
total no of e- in 2nd shell
A
8
12
Q
total no of e- in 3rd shell
A
18
13
Q
total no of e- in 4th shell
A
32
14
Q
sub-shell s: no. of orbitals and max no. of e-
and shape
A
s
orbitals = 1
max e- = 2
shape = spherical
15
Q
sub shell p: no of orbitals and max no. of e-
and shape
A
p
orbitals = 3
max e- = 6
shape= three dumbbell shaped orbitals at right angles to one another
16
Q
sub-shell d: no. of orbitals and max no. of e-
A
d
orbitals = 5
max e- = 10
17
Q
sub-shell f: no. of orbitals and max no. of e-
A
f
orbitals = 7
max e- = 14
18
Q
no. of e- in sub-shell s
A
2
19
Q
no. of e- in sub-shell p
A
6
20
Q
no. of e- in sub shell d
A
10
21
Q
no. of e- in sub-shell f
A
14
22
Q
what is an orbital
A
is a bit of space in which the electrons move in. orbitals within the same sub shell have the same energy.
23
Q
spin of electrons in orbitals
A
orbitals contain 2 electrons each.
they spin in opposite directions
spin pairing
24
Q
order in which shells fill
A
1s - 2s - 2p - 3s - 3p - 4S- 3D - 4p - 4d - 4f
main point: 4s fills before 3d
25
why does 4s fill before 3d
because 4s has a lower energy level than 3d even though its principle quantum number is bigger.
26
how do electrons fill orbitals in a sub shell
electrons fill orbitals singling before they start sharing
27
how are noble gas symbols used in sub shell notation
noble gas symbols like argon (Ar)
which e- configoration is: 1s(2) 2s(2) 2p(6) 3s(2) 3p(6)
are sometimes used .
e.g calcium: 1s(2) 2s(2) 2p(6) 3s(2) 3s(2) 3p(6) 4s(2) can be written as: [Ar] 4s(2)
28
define ionic bond
is an electrostatic attraction between two oppositely charged ions
29
electrostatic attraction
is the force that holds -ve and +ve ions together
30
sodium chloride structure:
giant ionic lattice
basic repeat unit of Na+ and Cl-
it foms becacuse each ion is electrostatically attracted in all directions to ions of opposite charge.
cine shape
31
melting point of sodium chloride
is extremely high (801 oc) because of the very strong ionic bonds
32
3 principles of ionic structure that explains behaviour of ionic compounds
1. ionic compounds conduct electricity when they're
molten or dissolved.
2. ionic compounds have high melting and boiling points.
giant ionic lattices are held together by strong
electrostatic forces which take a lot of energy to
overcome.
3. ionic compounds dissolve in water.
water molecules are polar and so interact with the
charges of the ions.
33
define covalent bonding
the strong electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atoms.
34
drawing covalent bonding
like ven diagram with shared electrons in the middle
35
boron trifluoride BF(3) is a special case in covalent bonding as
boron forms 3 covalent bonds with the 3 fluorine so that each fluorine has a full outer shell. however boron only ends up with 6 electrons in outer shell.
36
sulfur hexafluoride SF(6) is a special case in covalent bonding as
sulfur forms a covalent bond with 6 fluorines so that each have full outer shell (8 e-).
to do this sulfur EXPAND THE OCTET
so that it has 12 electrons in its outer shell.
37
average bond enthalpy
measures the energy required to break a covalent bond
| the stronger a bond is the more energy required and so greater bond enthalpy.
38
double - triple covalent bonds
multiple covalent bonds share more than one pairs of electrons between two atoms. e.g
O2 which is a double covalent bond and shares 2 sets of electrons.
N2 has 5 outer electrons so forms tripple bond to complete shell
CO2 carbon has 4 and needs another 4. each oxygen atoms share two electrons.
39
dative covalent bonds
where both electrons come from one atom
e.g
ammonium NH4+
nitrogen atom donates a pair of electromns to H+
40
how are dative covalent bonds drawn
as an arrow potting away from donor atom
41
shapes of molecules
refer to real life flash cards
42
linear molecules
2 electron pairs around central atom
no lone pairs
bond angle: 180
43
trigonal planar
3 e- pairs around central atom
no lone pairs
bond angle: 120
44
tetrahedral
4 e- pairs around central atom
no lone pairs
bond angle: 109.5
45
trigonal pyramidal
4 e- pairs around central atom
1 lone pair
bond angle: 107
46
nonlinear / bent
4 e- pairs around central atom
2 lone pairs
bond angle: 104.5
47
trigonal bipyramidal
5 e- pairs around central atom
no lone pairs
bond angle: 90 / 120
48
octahedral
6 e- pairs around central atom
no lone pairs
bond angle: 90
49
angle size according to lone pair to bonding pair
lone pair / lone pair = angles are the biggest
lone pair/ bonding pair = second biggest
bonding pair/ bonding pair = angles are the smallest.
50
which type of electron pair repels more
lone pairs of electrons repel more producing larger angles
51
shape of a molecule depends on
type of electron pairs surrounding central atom
| number of electron pairs
52
electron pair repulsion theory of compounds which all have 4 pairs of electrons
4 bonding pairs = 109.5 o
1 lone 3 bonding = 107 o lone pair repels the bonding pairs effectively squeezing them together
2 lone 2 bonding = 104.5 o 2 lone pairs reduce bond angle even more
53
electron pair repulsion theory
theory that predicts the molecular shape of molecules
54
electronegativity
an atoms ability to attract the electron pair in a covalent bond is called electronegativity
55
Pauling scale
mesures electronegativity. the greater an element pauling value, the higher its electronegativity.
56
polar molecules
if two atoms have different electronegativities in a covalent bond the bonding electrons will be pulled towards the more electronegative atom. making the bond polar. as it causes a permanent dipole.
the greater the difference in electronegativity the more polar the bond is.
57
dipole
dipole is the difference in charge between two atoms caused by a shift in electron density in the bond.
58
carbon and hydrogen have similar
electronegativities and so are essentially non-polar
59
how does arrangement of bonds change with permanent dipoles
the arrangement of polar bonds determines whether the molecule will have an overall dipole.
symmetrical arrangement cause dipoles to cancel out. so non polar
uneven arrangement causes causes molecule to have overall dipole as they do not cancel out and are polar
60
the transition between ionic to convalent bonding is
gradual.
only bonds between atoms of the same element can ever be purely covalent e.g diatomic gases. because the electronegativity difference is 0.
very few compounds are completely ionic.
61
how can u predict type of bonding
by looking a electrongativity difference
the higher it is the more ionic
the lower it is the more covalent.
62
intermolecular forces are very
Weak
63
name the types of intermolecular forces from strongest to weakest
1. hydrogen bonding
2. peranent dipole- dipole interactions (e.g ionic and
polar)
3. induced dipole-dipole or london forces
64
what are induced dipole-dipole forces
found between all atoms and molecules
cause attraction.
electrons charge clouds are constantly moving causing temporary dipoles where an electron is nearer one atom than another. this dipole induces dipoles in the opposite direction on a neighbouring atom causing them to be attracted. domino effect.
they are created and destroyed constantly causes overall attraction
65
variations on london forces strength
- larger molecules have larger electron clouds causing
stronger induced dipole - dipole forces.
molecules with greater surface area also have stronger - london forces as they have bigger exposed electron
cloud.
66
boiling points and london forces
when boiling liquid you overcome inter molecular forces as they molecules turn to gas.
more energy needed to overcome stronger inter molecular forces the greater the boiling point
67
induced dipole- dipole forces hold lattices
responsible for holding iodine in lattice
iodine atoms are held together by strong covanet bonds.
but the molecules are held in a lattice by weak induced dipole -dipole interactions
68
how are permanent dipole - dipole interactions drawn
s+ and s- used to label the dipoles of the atom and dotted line used to show the weak electrostatic force of attraction i.e. week bond between s+ of one molecule and s- of another
69
permanent dipole-dipole interactions happen in addition to..
induced dipole-dipole interactions (london forces)
70
hydrogen bonding can only happen when hydrogen is covalently bonded to...
fluorine, nitrogen or oxygen.
71
how does hydrogen bonding work?
- Hydrogen has a high charge density because it is so
small
- fluorine, nitrogen and oxygen are very electronegative
- the bond is so polarised that a weak bind forms
between hydrogen of one molecule and a lone pair of
electrons on the F, N, O in another molecule.
72
molecules with hydrogen bonding usually contain which groups?
-OH
-NH
water and ammonia both have hydrogen bonding
73
hydrogen bonding effect on properties of a substance
- they are soluble in water
| - higher boiling and freezing points than molecules of similar size
74
why is ice less dense than water
because hydrogen bonds are relative long.
75
explain why simple covalent compounds have __ melting and boiling points interms of inter molecular forces
they have low melting and boiling points because:
- inter-molecular forces are weak so don't need much
energy to break.
- they are often liquids or gases at room temperature
76
expain why polar molecules are _______ in water
polar molecules are soluble in water because:
- water is polar so only tends to dissolve other polar
substances.
- substances with hydrogen bonds can form hydrogen
bonds with water molecules, so will be soluble
- molecules that only have induced dipole-dipole forces
will be insoluble
77
simple covalent compounds don't conduct ________ because
simple covalent compounds don't conduct electricity because:
| - even though some have permanent dipoles overall covalent molecules are uncharged.
78
when you melt or boil a simple covalent bond you have to overcome the
inter molecular forces that hold the molecules together
