2.2 Flashcards
Most common ion of Zinc
Zn 2+
Most common ion of silver
Ag +
What does the suffix ‘ate’ tell us
The compound usually includes oxygen
Nitrate
NO3 -
Carbonate
CO3 2-
Sulfate
SO4 2-
Hydroxide
OH-
Ammonium
NH4 +
Why does oxygen tend to make 2 bonds
It’s in group 6
What does the suffix ‘ium’ tell us
The compound is very likely to have a positive charge
Bromide
Br -
Chloride
Cl -
Sulfide
S 2-
Nitride
N 3-
Phosphate
PO4 3-
Oxide
O 2-
Phosphide
P 3-
What does the suffix ‘ide’ usually tell us
The ion has no other elements in it
How many electrons does each shell hold
Up to 2n^2 where n = shell number or quantum number
What are the 4 sub shells in electrons
s
p
d
f
Which sub shell is in the first electron shell
s
Which sub shells are in the second electron shell
s and p
Which sub shells are in the third electron shell
s, p and d
Which sub shells are in the fourth electron shell
s, p, d and f
How many electrons does the ‘s’ sub shell hold
2
How many electrons does the ‘p’ sub shell hold
6
How many electrons does the ‘d’ sub shell hold
10
How many electrons does the ‘f’ sub shell hold
14
How many orbitals are in the ‘s’ sub shell
1
How many orbitals are in the ‘p’ sub shell
3
How many orbitals does the ‘d’ sub shell have
5
How many orbitals does the ‘f’ sub shell have
7
Atomic orbitals
A region within an atom, that can hold up to 2 electrons, which have opposite spins
Order of electron arrangement
Smallest to the largest
Electrons -> orbital -> sub shell-> shell
Shape of ‘s’ orbital
Spherical
Hund’s rule
Electrons prefer to occupy orbitals on their own, and only pair up when no empty orbitals of the same energy are available
Why does 4s fill up before 3d
4s has a lower energy than 3d
How can we explain the electron configuration for chromium and copper
Sub-shells like to be full or half full. As 4s and 3d are so close in energy, electrons can move between them easily, and rearrange to make full and half full sub shells
How are elements classified by blocks
The block refers to the the highest type of sub shell occupied by electrons e.g. Groups 1 and 2 are in the ‘s’ block because all of those elements highest sub shell is s
Which electrons are lost when a cation is formed
The highest energy electrons but the Sc - Zn elements are an exception
Why are the Sc - Zn elements an exception to the ion rule
4s and 3d energy levels are very close but once 4s is full, it’s energy is above 3d so empties first as well
How to write short structure electron configuration
Put noble gas closest to it then continue writing structure
E.g. manganese = [Ar] 4s^2 3d^5
Electronegativity
Ability to attract a bonding pair of electrons in a covalent bond
The greater the electronegativity of an atom …
… the more it attracts electrons towards it
Factors affecting electronegativity
Atomic charge
Distance from the nucleus
Electron shielding
The greater the difference in electronegativity between the bonded atoms …
… the greater the permanent dipole
Electronegativity across a period
Atomic radius decreases
Greater nuclear charge
Stronger attraction between bonding pair of electrons and nucleus
Increased electronegativity
Electronegativity down a group
Decreases due to increased shielding and distance from nucleus
Less attraction between nucleus and bonding pair of electrons
Electronegativity up a group
Increases
Basic rules of VSEPR theory
All single-bonded and lone pairs arrange themselves as far apart in space as possible
Lone pairs repel more strongly than bonding pairs/regions
Shape of molecule if there are 2 bonding pairs and 0 lone pairs
Linear
Shape of molecule if there are 2 bonding pairs and 1 lone pair
Non linear
Angle of linear shape
180
Angle of non linear shape
117.5
Shape of molecule with 3 bonding pairs and 0 lone pairs
Trigonal planar
Angle of trigonal planar shape
120
Ionic bonding
Electrostatic attraction between positive and negative ions
Covalent bonding
The strong electrostatic bond between a shared pair of electrons and the nuclei of the bonded atoms
Shape of 4 bonding pairs and 0 lone pairs
Tetrahedral
Angle of tetrahedral
109.5
Shape of 3 bonding pairs and 1 lone pair
Trigonal pyramid
Angle of trigonal pyramid
107
Shape of 2 bonding pairs and 2 lone pairs
Bent
Angle of bent shape (2 bonding, 2 lone)
104.5
Shape of 6 bonding pairs and 0 lone pairs
Octahedral
Angle of octahedral
90
Shape of 5 bonding pairs and 1 lone pairs
Distorted square pyramid
Angle of distorted square pyramid
89
Shape of 4 bonding pairs and 2 lone pairs
Square planar
Angle of square planar
90
When does covalent bonding occur
Between atoms of the same element (N2, O2)
Between atoms of diff. elements on the RHS of the table
When one of the elements is in the middle of the table
With head-of-the-group elements with high ionisation energies
Permanent dipole
A small charge difference across a bond due to a difference in the electronegativities of the bonded atoms
Polar covalent bond
A covalent bond with a permanent dipole
Polar molecule
A molecule with an overall dipole, when you take into account any dipoles across the bonds
Not symmetrical
Sizes of atoms increases
Down a group
Decrease across period
How does nuclear charge affect electronegativity
More protons
Stronger attraction between nucleus and bonding pairs of electrons
How does atomic radius affect electronegativity
The smaller the radius
Closer electrons are to nucleus
Stronger attraction between nucleus and bonding pair of electrons
How does shielding affect electronegativity
Less shells
Less shielding
Stronger the attraction between the nucleus and bonding pair of electrons
When does a molecule have London Forces
When it is not polar
When does a molecule have hydrogen bonds
When it is polar and has N-H/ O-H/ F-H bonds
When does a molecule have permanent dipole-dipole forces
When it is polar and doesn’t have any N-H/ F-H/ O-H bonds
Hydrogen bonds
Strong dipole-dipole attraction between electron-deficient H atoms and a lone pair of electrons in a highly electronegative atom on a different molecule
Relative strength of forces
Ionic/ covalent - 1000
Hydrogen bonds - 50
Dipole-dipole - 10
London forces - 1
How do london forces arise
Electrons in shells are continually moving and fluctuating
Creates uneven distribution of electrons
Sets up instantaneous dipole
Induces dipole in neighbouring atoms
Creates weak forces of attraction
Anomalous properties of water
Ice is less dense than water - molecules in ice held apart by hydrogen bonds
Ice has a relatively high mp- hydrogen bonds are stronger than other intermolecular forces
Bond angle in CH4
109.5
Bond angle in NH3
107
Bond angle in H2O
104.5
How does lone pairs affect the mp
The more the lone pairs, the higher the melting point
What colour is anhydrous copper sulfate
White
Blue when hydrated
Dative covalent bond
Both electrons have been donated by one atom
Why does iodine have a higher bp than fluorine
More electrons
Greater LF’s
More energy needed to overcome
All lattices are …
Giant
Graphite
Giant covalent lattice w/ layers
Has delocalised electrons
Soft - layers can slide over each other and has LF
Why does pure water not conduct
No free charge carriers
Bonding in NH4 +
4 bonding pairs (one dative)
Tetrahedral
Why is the 1st IE lower for Al
e- released from 3p which has a higher energy than 3s
Why is the 1st IE lower for S
Pair of e- removed from 3p
Repulsion between paired e-
Graphene
One layer of graphite
How are atoms vaporized before going into a mass spectrometer
Dissolved in a volatile solvent
How does a mass spec work
High voltage supply
Vaporized atoms accelerate through capillary tube
Removes electrons to form ions
The ones with a lowest m/z are identified firs
Bigger current = more abundance
Shape of 5 bonded pairs
Bipyramidal
Angle in trigonal bipyramidal
90 and 120
Drawing trig. bipyramidal
Three on the plane
One doing back into the plane
One coming out of the plane
Polar bonds
C-O
C-N
N-O
S-F
What happens to the bond angle when you add a lp
Decreases by 2.5
s block elements
Gp 1 and 2
p block elements
Gp 13 to 18
d block elements
Transition metals
f bock elements
Lanthanoids and actanoids
Why is there a drop in IE between Be and B
2p has a higher energy than 2s so outermost e- is easier to remove
Why is there a drop in IE between N and O
Both outermost e- in 2p
In O paired e- in 2p repel each other so its easier to remove the outermost e-
