2.2 Foundations In Chemistry 5.1-6.4 Flashcards
Ionisation
A process in which atoms lose or gain electrons and become ions
First ionization energy definition
The energy required to remove one mole of electrons from one mole of gaseous atoms to produce one mole of gaseous cations
M(g) —> M^+ (g) + e^- (g)
Second ionization energy definition
Involved the removal of a second electron
M^+ (g) —> M^2+ (g) + e^- (g)
Evidence for energy levels
Plotting the successive ionisation energies of an element clearly shows the existence of different energy levels and the number of electron in each
Successive ionisation energies increase as more electrons are removed
Large jumps in the ionisation energy reveal where electrons are being removed from the next principal energy level
What are the 4 sub levels
S p d f
Max number of electrons for each sub level
S- 2
P- 6
D- 10
F- 14
Principal energy levels’ sub levels and max electrons
1- 1s -2
2- 2s 2p -8
3- 3s 3p 3d -18
4- 4s 4p 4d 4f -32
Order of energy sub levels
1s 2s 2p 3s 3p 4s 3d 4p 4d 4f
The Aufbau principle
the lowest energy sub levels are occupied first
Who developed the Aufbau principle
Niels Bohr
Why does the 4s sub level fill up before the 3d
It’s lower in energy
Hydroxide
OH-
Sulphate
SO4^2-
Carbonate
CO3^2-
Nitrate
NO3^-
Ammonium
NH4^+
Silver
Ag+
Electron orbitals
It is impossible to exactly locate the position of an electron within an energy sub level. By measuring the electron density around the nucleus, is possible to define regions where electrons are most likely to be found at any one time.
Each energy sublevel has one or more orbitals, each one can contain a maximum of two electrons
Name the sublevels, the number of orbitals and the maximum number of electrons within them
S – 1– 2
P – 3–6
D – 5–10
F – 7–14
The Pauli exclusion principle
States that each orbital may contain no more than two electrons
Spin
The Pauli exclusion principle introduces a property of electrons called spin, which has two states: up and down. The spins of electrons in the same orbital must be opposite
Spin diagram
Shows how the orbitals are filled. Orbitals are represented by squares, and electrons by arrows pointing up or down
Rules for filling electrons in spin diagrams
When two electrons occupy a p-sub level, they could either completely fill the same p orbital or half fill two different P orbitals
Hund’s rule states that the single electrons occupy all empty orbitals within a sub level before they start to form pairs in orbitals
What happens when two electrons are in the same orbital
There is repulsion between them due to their negative charges
The most stable configuration is with single electrons in different orbitals
Hund’s rule
States that single electrons occupy all empty orbitals within a sub level before they start to form pairs in orbitals
Why is the electron configuration for chromium and copper different
In each case the 4s orbital contains one electron
This is because the 4s and 3d sub levels lie very close together in energy, and the 3d being either half full or completely full is a lower energy arrangement
S orbital shape
Spherical
P orbital shape
Dumbbell
Electron relative mass
1/2000
Ionic bond definition
The electrostatic attraction between pos and neg ions
Holds tog cations (pos ions) and anions (neg ions) in ionic compounds
Ionic compounds structure and properties
Results in giant ionic lattice structure
High mp and bp
Strong electrostatic forces between opp charged ions
Higher for lattices w ions w larger ionic charges
Conducts elec when when melted or aqueous solution
Dissolved in polar solvents (e.g. water)
Covalent bonds of compound ions examples
Hydroxide ion
Sulphate ion
Nitrate ion
Carbonate ion
Covalent bonding diagram of hydroxide ion
OH-
Single bond
Oxygen has extra electron from unknown metal
Covalent bonding diagram of sulphate ion
SO4^2-
Two oxygen atoms two oxygen ions
Double binds with oxygen atoms. Single with ions
O ions have extra electron from Unknown metal
Covalent bonding diagram of nitrate ion
NO3^-
Double bond with one Oxygen
Single bind with oxygen that has extra electron from unknown metal
Single (dative) bond with oxygen (both electrons from nitrogen)
Covalent bonding diagram of carbonate ion
CO3^2-
Double bond with O
Single bind with two O that both have extra electron from unknown metal
Covalent bond definition
The electrostatic force if attraction between the shared pair of electrons and the 2 adjacent nuclei of the atoms
Covalent bond- orbital overlap
A covalent bond is the overlap of atomic orbitals, each containing one electron to give a shared pair of electrons
Sigma bond
Single covalent bond
S and S
S and P
P and P
Dative covalent bond
(Coordinate bond)
Covalent bond in which the shared pair of electrons has been supplied by one of the bonding atoms only. The shared electron pair was originally a lone pair of electrons on one of the bonded atoms
Average bond enthalpy
A measurement of covalent bond strength
Electron pair repulsion theory
The electron pairs surrounding a central atom determine the shape of the molecule or ion
The electron pairs repel one another so that they’re arranged as far apart as possible
The arrangement of electron pairs minimises repulsion and thus holds the blinded atoms in a definite shape
Different numbers of electron pairs result in different shapes
Shape of molecules and ions with 2 electron pairs around central atom
Linear
180
Shape of molecules and ions with 3 electron pairs around central atom
Trigonal planar (2D) 120
Shape of molecules and ions with 4 electron pairs around central atom
Tetrahedral
109.5
Shape of molecules and ions with 5 electron pairs around central atom
Trigonol bipyramidal
90
120
Shape of molecules and ions with 6 electron pairs around central atom
Octahedral
90
Distortion of basic shapes due to lone pairs- 3 electron pairs and 1 lone pair
Trigonometrische pyramidal
107
Distortion of basic shapes due to lone pairs- 2 electron pairs and 2 lone pairs
Bent
104.5
Distortion of basic shapes due to lone pairs- 4 electron pairs and 2 lone pairs
Square planar
90
Bonded pair and lone pair repulsion
A lone pair of electrons is slightly closer to the central atom and occupies more space than a bonded pair
Resulting in a lone pair repelling more strongly than a bonding pair
What’s the bond angle reduced by for each lone pair
2.5
Electronegativity
The ability of an atom to attract a bonding or non bonding pair of electrons
Electronegativity- what changes when the bonded atoms are different elements?
The nuclear charges are different
Atoms may be different sizes
The shared pair of electrons may be closer to one nucleus than the other
Electronegativity- ionic or covalent?
If the electroneg difference is large, one bonded atom will have will have a much greater attraction for the shared pair than the other bonded atom
The more electroneg atom will have gained control of the electrons and the bond will now be ionic rather than covalent
Non polar bonds
The bonded electron pair is shared equally between the bonded atoms
A bond will appear polar when:
- the bonded atoms are the same or
- the bonded atoms have the same or similar electronegitivity
Polar bonds
The bonded electron pair is shared unequally between the bonded atoms
A bond will be polar when the bonded atoms are different and have different electroneg values, resulting in a polar covalent bond
Dipole
The separation of opposite charges
Dipole in a polar covalent bond
Doesn’t change
Permanent dipole
Polar molecules
Asymmetrical
Dipoles don’t even each other out
3 intermolecular forces
Induced dipole-dipole interactions (London forces)
Permanent dipole-dipole interactions
Hydrogen bonding
Strongest force
Covalent bond
Induced dipole-dipole interactions (London forces)
Weak intermolecular forces that exist between all molecules, whether polar or non polar
Act between induced dipoles in diff molecules
Only temporary
Origin of induced dipole-dipole interactions (London forces)
Movement of electrons produces a changing dipole in a molecule
At any instance, an instantaneous dipole will exist, but it’s position is constantly shifting
The instantaneous dipole induces a dipole on a neighbouring molecule
The induced dipole induces further dipoles on neighbouring molecules, which then attract one another
Permanent dipole-dipole interactions
Act between the permanent dipoles in diff polar molecules
Properties of simple molecular substance
Low mp
Low bp
Weak intermolecular forces break when melted
Covalent bonds are strong and don’t break
Hydrogen bonds
Special type of permanent d-d interaction found between molecules containing:
- an electroneg atom with a lone pair of electrons (O, N, F)
- a hydrogen atom attached to an electroneg atom (H-O, H-N, H-F)
H bonds acts between a lone pair of electrons on an electroneg atom in one molecule and a H atom in a diff molecule
Strongest type of intermolecular attraction
Hydrogen bonds
Anomalous properties of water
The solid (ice) is less dense than the liquid (water) Water has relatively high mp and bp
Anomalous prop of H2O- ice less dense than water
H bonds hold water molecules apart in open lattice struc
The water molecules in ice are further apart than in water
Solid ice is less dense and floats
Anomalous prop of water- high mp and bp
Water has London forces between molecules (like all molecules, not just water)
H bonds are extra forces over and above the London forces
An appreciable quantity of energy is needed to break the hydrogen bonds in water, so water has much higher melting and bp than would be expected from just London forces
Without H bonds, water would have a bp of …
-75oC
Strength of London forces increases as..
Molecular size increases
Number of protons and electrons increase
Branches decrease
Why does the strength of London forces increase as molecular size increases
As atomic radius increases the outer electrons become further from the nucleus
They are attracted less strongly by the nucleus and so temporary dipoles are easier to induce
How do the number of protons and electrons effect the strength of London forces
The more pros and elecs the stronger
The greater the charges of the pos nucleus and neg elecs are so stronger force of attraction
How does branching in molecules effect London forces
The more branching the weaker the forces
As The less the molecules can pack together so there is less surface contact between molecules
If the molecule is polar then there are…
“ non polar “ …
Exception…
Permanent d-d attractions
London forces
Linear molecules with polar bonds (not polar molecule as symmetrical)
Strength of metallic bonds depends on
Charge of metal ions
Size of ions
Charge of metal ions relative to strength
Greater the charge the greater the attraction between the ions and the delocalised elecs and the stronger the metallic bonds
Size of metal ions relative to strength of metallic bonds
Smaller the metal ion the closer the positive nucleus is to the delocalised elecs
Meaning there is a greater attraction between the two, which creates a stronger metallic bond
