Brainscape's Knowledge GenomeTM

Browse over 1 million classes created by top students, professors, publishers, and experts.


See full index

Chemistry 10003 > Inorganic Chemistry > Flashcards

Inorganic Chemistry Flashcards

1
Q

Lewis Acids and Bases

A

Lewis acid = electron acceptor (e.g. protons, BF3, SiF4)

Lewis base = electron donator (e.g. NH3, CN-, OH-, H2O)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
2
Q

Bronsted-Lowry Acids and Bases

A

Bronsted-Lowry acid = proton donator
(e.g. CHCOOH (formic acid), acetic acid)

Bronsted-Lowry Base = proton acceptor
(e.g. NaOH, ammonia)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
3
Q

Representation of hydrated protons

A

Hydrogen ions do not exist as an ‘naked’ ion in aqueous solutions (it is hydrated, degree of hydration is unknown)

Hydrogen ions are both represented by H+ and H3O+ (hydronium)

All ions/molecules are surrounded by solution to form a SOLVATION SHELL

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
4
Q

Conjugate Acid-Base pairs

A

Conjugate base = acid that lost a proton/electron (paired with a acid)

Conjugate acid = base that gained a proton/electron

We can calculate Ka from the conjugate acid of the base (+ vice versa)
as Ka x Kb = Kw
Ka/Kb is the acid/base dissociation constant (strength of acid/base in a solution)
= [H+][OH-] thus pKa + pKb = 14 (only at 25 degrees)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
5
Q

Ionisation of water

A

Pure water undergoes self-ionisation where water can behave both like an acid or base (weak acid/base)

As it is a weak acid/base, the left side is favoured

Kw = 1 x 10^-14 at 25 degrees
(ionisation constant of water)

In pure water, concentration of hydrogen and hydroxide is equal
√ Kw = 10^-7 M
(solution is said to be neutral)

Ionisation constant of water is dependent on temperature
-increasing temperature decreases the ionisation constant of water (concentration of ions are lesser)
(adding substances may change concentrations oh H+ and OH- but product of concentrations are always constant at a given temperature)

e.g. of H+ = 10^-1 M then OH- = 10^-13 M

If H+ concentration is > 10^-7 = solution is acid
If OH- concentration is > 10^-7 = solution is basic

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
6
Q

pH and pOH scale

A

Since concentrations of hydrogen/hydroxide is very small, we express them in log scale

pH scale is in logarithmic scale indicate that hydrogen ions increase by a scale of 1,000

Units: mol/L, mol dm3 or M

pH = -log[H+]
pOH = -log[OH-]
(if H+ = 10^-1M then pH = 1, pOH = 13)

pH + pOH = 14

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
7
Q

Quantifying acid strength

A

Ka = acid dissociation constant
Ka = [A-][H3O+] / [HA]
(water does not get included as it is solution and concentration doesn’t change)

STRONG ACIDS = HIGH Ka AND SMALL pKa
large Ka = reaction lies to the right

pKa = dissociation in water (lower value = higher dissociation)
pKa = -log[Ka]
Ka = 10^-pKa
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
8
Q

Quantifying base strength

A 
Kb = base dissociation constant 
Kb = [BH+][OH-] / [HB]

STRONG BASES = HIGH Kb AND SMALL pKb

pKb = -log[Kb]
Kb = 10^-pKb
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
9
Q

Strong acids and bases

A

Strong acids/bases almost fully dissociate in water (we write one arrow in equation)

Strong acids = HCl, HBr, sulfuric acid (H2SO4), HClO4, HBF4
conjugate bases of strong acids are weak acids (anions)
e.g. HCl- or HBr-

We assume concentration of H+ = HA (acid conc at equilibrium) as it has fully dissociated
OH- = HB

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
10
Q

Weak acids and bases

A

Weak acids only go through small degree of ionisation (e.g. carboxylic acids)
Weak bases go through small degree of protonation (e.g. ammonia, amines)
-equilibrium lies to the left

In weak acids, conjugate base conc is assumed same as protons (conjugate acid) [H3O+/H+] = [A-]

How do we know if it is a weak acid/base though?
5% RULE
-if dissociation is less than 5%, we classify it as a weak acid/base
[H+] x 100 / acid concentration to get dissociation percentage

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
11
Q

Buffers

A

Buffer solution = solution that resist pH change when small amounts of acid/base are added

Buffers can be formed by:
-weak acid + conjugate base
-weak base + conjugate acid
(good buffers will have equal proportions of the acid/base and its conjugate pair)
(also can through adding enough strong acid/base to neutralize half of weak acid/base)

Qualitative explanation of buffers

  • a buffer from acetic acid and its conjugate base has protons added to it, the conjugate base are able to pick up free protons to form acetic acid
  • protons will not be free in solution, causing pH to not change as much

Buffer strength is greatest when pH = pkA
In buffers, pKa range is +/- 1 but when pH = pKa then we can add acid and base to solution

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
12
Q

Henderson-Hasselbach Equation

A

pH = pKa + log [A-] / [HA]

pH = pKa as the concentration ratio between acid and conjugate base is 1, where log(1) = 0

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
13
Q

Acid-base titration

A

Equivalence point = equal amount of acid + base added (pH = 7; for a strong acid-base reaction)

With a weak acid/base + strong acid/base reaction
Buffer region = lots of acid/base has to be added to see small pH change
Equivalence point would be higher/lower than pH = 7
(e.g. adding strong base to weak acid gives equivalence point of pH = more than 7 indicating the conjugate base of weak acid is basic)

Halfway of equivalence point = pH (pKa of weak acid)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
14
Q

Ionic vs Covalent Bonds

A

Ionic bonds = cations and anions held together by electrostatic forces (bonds held together by the electrostatic forces of oppositely charged ions)

  • non directional bonds
  • 400-4000 kJ/mol

Covalent bond = bonds held together by the electrostatic force between a positively charged nuclei and shared electrons
-directional bonds
-150-550 kJ/mol
Valence electrons = outer shell electrons are shared between nuclei (bonding will occur if the energy of the molecule if lower than each separate atom)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
15
Q

Lewis Structures

A

bonding model that demonstrates the arrangement of valence electrons in a molecule

OCTET RULE = many stable compounds have atoms that achieve an octet of electrons
-INTRINSIC (independent of amount present) stability related to having 8 valence electrons (group 18 = noble gases are inert)
PERIOD 2 + 1 ELEMENTS CANNOT BREAK OCTET RULE

Formal charge = charge assigned to an atom assuming all the electrons are shared equally
Equation = [number of valence electrons] - [number of lone pair electrons] - [0.5 x electrons in bonds]

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
16
Q

Resonance structures

A

2 Equivalent lewis structures result in resonance
-lewis structures do not define structure accurately, but defines length is the average length between resonance structures

Ozone (O3) = 2 resonance structures
Nitrate ion (NO3-) = 3 resonance structures

Isoelectronic molecules = same number of valence electrons

RESONANCE AND BREAKING OF OCTET RULE SHOWS LIMITATIONS OF LEWIS STRUCTURES

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
17
Q

Coodinate (dative) bond + Radicals

A

Structures where molecule has less than 8 valence electrons (e.g. BF3, it is a lewis acid as it is electron deficient)

Coordinate bonds form when a covalent bond between an atom where electrons are from 1 atom

Radicals = molecules that have unpaired electrons (are very reactive)
e.g. NO, OH・, NO2

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
18
Q

VSEPR Theory

A

Valence Shell Electron Pair Repulsion - demonstrates molecular shape (allowing us to understand chemical and physical properties of compounds)

Assumptions: atoms in a molecule are held together by a pair of electrons, bonding pairs (BP)

  • some atoms within a molecule may not have pairs of electrons that were involved in bonding, lone pairs (LP)
  • electron pairs are negatively charged and repel each other (electron pairs adopt positions to position themselves as far as possible)

Electron geometry = shape of molecule regarding electron pairs
Shape = actual shape of molecule

We fill out equitorial positions before axial (lone pairs are further from groups at axial positions compared to equitorial)

With multiple bonds, electron-electron repulsion means double/triple bonds take up more space than single bonds (excluding resonance structures)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
19
Q

VSEPR Theory Values

A

2 bonded pairs, 0 lone pairs = linear (180 degrees)

3 bonded, 0 lone = trigonal planar (120 degrees)

2 bonded, 1 lone = bent/v-shaped (119 degrees)

2 bonded, 2 lone = v-shaped, bent (104.5 degrees)

3 bonded, 1 lone = trigonal pyramidal (107 degrees)

3 bonded, 2 lone = T-shaped (87.5, 175 degrees)

4 bonded, 0 lone = tetrahedral (109.5 degrees)

4 bonded, 1 lone = seesaw (90, 120, 180 degrees)

5 bonded, 0 lone = trigonal bipyramidal (90, 120, 180 degrees)

5 bonded, 1 lone = square pyramidal (90, 180 degrees)

4 bonded, 2 lone = square planar (90, 180 degrees)

6 bonded, 0 lone = octahedral (90, 180 degrees)

2 bonded, 3 lone = linear

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
20
Q

Bond Polarity

A

Electronegativity = power of an atom in a molecule to attract electrons to itself (the greater the electronegativity difference, the more polar the bond)
Electronegativity increases across a period, decreases down a group
Differences in electronegativity can cause a ‘dipole moment’ (travels from positive end to negative end)

Dipole moment - degree of polarity, measured in Debye (D) (3.356 x 10-30 C m)
(the greater the electronegativity difference, the greater the dipole moment)
(bonds can be polar, but molecule does NOT have to be)

A symmetrical molecule with polar bonds may be non polar (e.g. carbon dioxide) OR unsymmetrical molecule but polar (e.g. water)

Electron affinity = amount of energy required to add 1 mole of electrons in gaseous state (opposite of ionisation energy, first electron affinity is always negative, second is positive)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
21
Q

Valence Bond Theory

A

Valence bond theory assumes that electrons are either localised in bonds between two atoms or localised on a single atom

We use valence bond theory when describing organic molecules (2s and 2p orbital present), purely covalent solids (e.g. isotopes of carbon)

Hybridisation:
sp3 = 2s orbital and 3/3 p orbital
sp2 = 2s orbital and 2/3 p orbital
sp = 2s orbital and 1/3 p orbital
2px --> 2py --> 2pz

LIMITATIONS:

  • does not explain trigonal bipyramidal/octahedral molecules
  • does not explain magnetic properties
  • incorrectly assumed localisation of electrons (concept of resonance must be added)
  • does not handle unpaired electrons well
  • does not give information about bond energy
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
22
Q

Molecular Orbital Theory (MOT)

A

Molecular orbitals are obtained by the linear combination of atomic orbitals

e.g. Hydrogen (H2) has 1 atomic orbital each atom
Electrons have wave-like properties that can overlap constructively or destructively

Contructively = increase of electron density (forms a bonding molecular orbital)
Destructively = 0 electron density between two atoms (node) (forms an antibonding molecular orbital)

Nodal plane makes an atom lie higher in energy compared to the bonding molecular orbital

Atomic orbitals that DO NOT overlap and remaind the same energy = non-bonding orbitals
(number of molecular orbitals = atomic orbitals)

BOND ORDER = indicator of stability of bond
(number of electrons in bonding orbitals) - (# antibonding electrons) / 2

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
23
Q

MOT and molecular ions

A

Example: H2+ or H2-
H2 + only has one atom contributing one electron meaning there is only one electron present in molecular orbitals
bond order = 1/2
H2- has 1 atom (1 electron) and another atom (2 electrons) combining with a total of 3 electrons
bond order = 1/2 yet it is weaker than h2 as the bond lies further apart

MOT demonstrates that 1 and 3 electron bonds are allowed (compared to valence theory)

With an He2 molecule, the bond order becomes 0 = it does not exist

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
24
Q

MOT and linear combination of p orbitals

A

As p orbitals have 3 orthogonal directions thus overlapping have different positions

P orbitals can overlap end-on-end:
2px orbitals (both left sides) = antibonding molecular orbital (sigma orbital)
2px orbitals (facing each other) = sigma bonding molecular orbital
P orbitals can overlap side to side:
2pz orbitals (opposite ends) = pi antibonding molecular orbital
2pz orbitals (same side) = pi bonding molecular orbital

x and z oriented p orbitals can not overlap as the bonding and antibonding interactions cancel each other out

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
25
Q

MOT and non bonding orbitals (HF)

A

1s orbital of hydrogen is much higher energy compared to 2s orbital of F therefore there is no interaction between the two

2s atomic orbital remains on its own = non bonding orbital

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
26
Q

MOT and relative energies

A

The closer in energy of atomic orbitals are, the better overlapping will occur (end on end is better causing big difference of energy between bonding and antibonding orbitals)

Side on overlap is not as stable compared to end on overlapping
HOWEVER
Antibonding side on overlapping is more stable compared to bonding side on overlapping, therefore it has more energy

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
27
Q

MOT and Oxygen’s diagram

A

Diagram of oxygen’s molecular orbital theory demonstrates that oxygen has 2 unpaired electrons, explaining the magnetic characteristic
-2 unpaired electrons in the pie antibonding molecular orbitals = oxygen is paramagnetic

Paramagnetic = substance/molecule is attracted to a magnetic field, attracting by an external magnetic field, causing an induced, internal magnetic field, due to unpaired electrons

Diamagnetic = where all electrons are paired, and repels magnetic field (creates an internal magnetic field in opposite direction)

28
Q

Metallic bonding

A

Band model = explains why some elements/molecules are metals, insulators or semi-conductors

(consider an array of atoms, with their atomic orbitals are overlapping from a large number of metal atoms = virtual continuum of energy levels = bands)

Consider Lithium:
As the atoms of lithium increases, more atoms interact with each other = increase of molecular orbitals + get closer and merging into a band

Conduction band = highest energy level of band with no electrons
Valence band = highest band with electrons

When energy is added, electrons can be promoted into conduction band (this characteristic allows metals to conduct electricity)
Application of an electric field causes negative electrons to move to positive side and positive holes to the negative side

Metals = band gap is very small
Semi-conductors = band gap is small (e.g. silicon, germanium)
Insulators = band gap is too large for electrons to jump up (thermal energy cannot promote electrons to conduction band)
29
Q

Trends in periodic table

A

Periodic table - arrangement of elements based on their electronic structure (valence electrons)

-the number of valence electrons determines their chemical behaviour

Groups = vertical lines
Period = horizontal lines
30
Q

Trends in periodic table (Atomic radius)

A

Atomic radius is dependent on the volume of space that electrons are allowed to occupy
(nature of electrons in orbitals means that atomic boundaries are ‘fuzzy’ and difficult to define)

ATOMIC RADIUS off an element is half the distance between the nuclei of neighbouring atoms in pure element (covalent radius)

Van Der Waals Radii = distance between two molecules that are not chemically bonded (halving the distance between two molecules)

Atomic size DECREASES IN SIZE ACROSS A PERIOD
-nuclear charge increases whilst number of shells remain the same
-electrons in outer shell are pulled in more tighter by increasing positive charge
(for far right elements, as the shell gets more crowded, due to electron-electron repulsion, atomic size becomes a little larger (seen in period 2 elements)

Atomic size INCREASES IN SIZE DOWN A GROUP

  • shell number increases (principal quantum number increases)
  • the larger the principal quantum number is = the larger amount of volume the orbitals take up + greater distance of outer electron from nucleus
31
Q

Size of ions

A

Cations (positive ions) are always smaller than their parent atoms
Anions (negative ions) are larger than their parent atoms

-negative ions have greater electron-electron repulsion = larger

32
Q

Size of Isoelectronic Species

A

Isoelectronic species = two species are isoelectronic when they have the same number of electrons

e.g. O2(2-) and F- are isoelectronic = 10 electrons

As the number of protons increase, the electrons are held more tightly thus atomic size is smaller
Since increasing number of positive charge is acting on same number of electrons, stronger positive charge = smallest size

33
Q

Trend of Ionisation Energy

A

Ionisation energy = energy required to remove the highest energy (most loosely held) electron from an atom in gaseous state

ALWAYS ENDOTHERMIC (requires energy)

Ionisation energy INCREASES ACROSS A PERIOD

  • as number of protons increases, the nuclear charge increases, increasing the difficulty to remove an electron from the same valence shell
  • less shielding as core electrons are less efficient in shielding outer electrons from nucleus (valence electrons are closer to nucleus)
  • large drop between group 8 and group 1 as most outer electron is in a NEW SHELL

Ionisation energy DECREASES DOWN A GROUP
-most furthest electron is further away from nucleus indicating it is easier to remove

High ionization energy = smaller size; low ionisation energy = larger size

34
Q

Trend of Electron Affinity

A

Electron affinity = energy change associated with the addition of an electron to an atom in gaseous state

X + E- –> X-

Generally, first electron affinity is exothermic = energy released by the attraction of nucleus to added electron is greater than the energy absorbed to overcome inter-electronic repulsion
Second electron affinity is endothermic = adding an electron to a negative ion (requires more energy to overcome the inter-electronic repulsion between both negatively charged species)

Generally, electron affinity INCREASES ACROSS A PERIOD
-has a more negative value, due to increasing nuclear charge

Electron affinity DECREASES DOWN A GROUP
-electron is being added further away from nucleus

(electron affinity is not always exothermic ; e.g. adding to an anion = endothermic)

35
Q

Trend of Electronegativity

A

Electronegativity = the ability of an atom to attract shared pairs of electron to itself (attract electrons in a covalent bond)

Assessed with pauling scale (0-4)
Using this scale: fluorine has value of 4.0 (group 17)
Cs = 0.9 (group 1)

Most electronegative elements are in the TOP RIGHT of periodic table (excluding noble gases)

Generally, electronegativity INCREASES ACROSS A PERIOD
Electronegativity DECREASES DOWN A GROUP

36
Q

Trend of metallic characteristics

A

Increase of metallic characteristics : DOWN A GROUP

Decrease in metallic characteristics : ACROSS A PERIOD

37
Q

Intermolecular forces (London dispersion)

A

Intermolecular forces = forces that attract molecules together

They are considered a type of secondary bonding
-comparatively weak compared to interactions INBETWEEN molecules (intramolecular forces)

London Dispersion forces (van Der waals forces)

  • weakest intermolecular forces that occurs between ALL molecules (both polar + non polar); 2kJ/mol
  • are temporary attractions due to instantaneous dipoles arising from fluctuations in electron distributions
  • dispersion forces arise when AN ELECTRON IN TWO ADJACENT ATOMS OCCUPIES A POSITION THAT FORMS TEMPORARY DIPOLES
  • instantaneous dipole in one atom can induce dipoles in neighbouring atoms/molecules
  • non-polar molecules attract each other as spherical distribution of electron density is not constantly maintained which can lead to instantaneous dipoles (can induce dipoles in other molecules)
  • dispersion forces are stronger for atoms that more easily polarised (easier to distort electron clouds)

(high polarization is associated with large numbers of electrons = stronger dispersion forces)

-dispersion forces are only effective at small distances
-shape can affect dispersion forces
(e.g. pentane and neopentane have same number of carbons and hydrogens yet pentane has a higher boiling point than neopentane
this is due to cylindrical shape of pentane = allowing all molecules to come into contact with another + closer
neopentane = spherical)

38
Q

Intermolecular forces (Dipole-Dipole)

A

Dipole-dipole interactions occur between polar molecules

Molecules have a net electrostatic attraction meaning dipoles align to the lowest energy state

Magnitude of dipole-dipole forces are found in 1% of all covalent/ionic bonds

39
Q

Intermolecular forces (Hydrogen bonding)

A

Special type of dipole-dipole bonding that is very strong (~10-60 kJ/mol)

-when a hydrogen is bonded to a small electronegative ion, the interaction between H-X dipole moment and polar molecules is much greater for ordinary dipole-dipole interactions
(hydrogen has a high magnitude to sigma positive that can interact with other electronegative atoms)
-interaction depends on magnitude of X-H dipole
(if bond was slightly polar only, interaction would be weaker)

Strongest hydrogen bonds comes from: O, N and F

There is often asymmetry in the hydrogen bonds (one hydrogen would be closer/further to another); closer to linear (typically around 160-170 degrees)

40
Q

Structure of ice

A

Ice forms an open structure due to hydrogen bonds in between water molecules

This leads to ice having a lower density than water + reason why ice floats on water
(4 hydrogen bonds, 2 acceptors and 2 donators)

41
Q

Gas solubility in water

A

Generally, the solubility of gases decrease with an increase of temperature

Increasing pressure has little effect on the solubility for solids + liquids but increasing pressure, increases solubility of gases in water

HENRY’S LAW (describes the relationship between gas pressure and the concentration of gas dissolved)
-amount of gas dissolved in a solution is directly proportionate to the pressure of gas (at a given temperature)
s = k(h) x p
s = gas solubility
k(h) = henry’s constant
p = partial pressure of gas

The higher Henry’s constant = the greater the solubility of gas

(e.g. Lake Nyos = cold water, high pressure holds lots of carbon dioxde (ocean is a sink), where the layers switched, releasing carbon dioxide in the atmosphere = killed many)

42
Q

Solubility of ionic solids

A

Examples of ionic solids (ores) where their solubility can affect their availability in the earth’s crust

The equilibrium constant for a sparingly soluble ionic solid is SOLUBILITY PRODUCT (Ksp) (only based on aqeous ions, NOT SOLIDS)
(equation has to be stoichiometrically balanced)

COMMON ION EFFECT = solubility of a sparingly soluble salt is reduced due to presence of a common ion
e.g. in barium sulphate, if sulphate ions were originally present, the equilibrium concentration of barium ions will decrease to keep the solubility product between sulphate ions and barium ions the same

Common ion effect is used in BARIUM MEAL (radiology)
-BaSO4 (barium sulfate) is made up as a slurry (mixture of dense solid and less dense liquid) with sodium sulfate to reduce the solubility of Ba2+ ions (which are toxic) forming a chunky solution
(sodium sulfate pushes equilibrium to solid barium sulfate)
-when swallowed barium meal, it will coat the intestinal tract which x-rays can detect (barium is a heavy metal that absorbs x-rays)

Solubility of ionic solids in environment

  • Shells are made of calcium carbonate where it is slightly soluble allowing living organisms to continuously produce
  • tooth enamel contains Hydroxyapatite where OH- groups become vulnerable to attack in acidic conditions (drinking cola; pH=2) leading to decay
  • F- is isoelectronic to OH- where we can replace OH- group with F- (repairs micro faults in teeth which can be colonised by bacteria; fluorapatite is less soluble in acidic conditions)

PRECIPITATION:
Ion product (Q) = quotient of the ion products (concentrations do NOT have to be at equilibrium)
For precipitation to occur, Q > Ksp until Q = Ksp
If Q < Ksp , NO PRECIPITATION will occur

43
Q

pH and Solubility

A

In the case of magnesium hydroxide, adding more OH- will shift the equilibrium to the left = DECREASES solubility

If we add hydrogen, solubility will increase as it will reduce OH- ions to form water (equilibrium shifts right)

Milk of Magnesia is used to treat stomach issues where Mg(OH)2 in stomach will reduce acidity as water will be formed

In the case of calcium carbonate, it can be dissolved by rainwater that absorbs carbon dioxide (ground water = acidic)
In acidic conditions, equilibrium of calcium carbonate shifts right and more carbonate ions are formed
This is how caverns and caves are formed

44
Q

Formation of ionic solids

A

Consider the formation of NaCl

  • to make gaseous sodium and chloride from solid requires energy
  • the main driving force of forming NaCl is the favourable association of ions that is present in solid state = sodium chloride lattice
  • the formation is very exothermic = solid product is stable

An ionic solid forms because the energy of the oppositely charged ions in the crystal lattice is lower than the energy of the original elements

The formation of an ionic solid will depend on:

  • the energy required to form the cation and anion
  • the strength of interaction between cation and anion in solid state
45
Q

Lattice enthalpy

A

Lattice enthalpy = energy required for the conversion of ions in one mole of ionic soild into its gaseous ions

MX (s) –> M+ (g) + X- (g)

Lattice enthalpy can be calcuated:
Lattice energy = -k(q1q2)/r
k = constant related to arrangement of ions
q1 / q2 = charges on ion
r = interionic centre-to-centre distance

(negative sign is in front to cancel out negative (from negative charge on ion)

Lattice enthalpy is proportional to ionic charges as bigger charges = stronger attraction in ionic solid
Inversely proportional to distance, as smaller ions = more energy to separate ions

46
Q

Born-Haber Cycle

A

We can use the born-haber cycle to determine the lattice enthalpy

The difference in energy between NaF(s) and its ions = lattice enthalpy

Steps in cycle:

  • formation
  • atomization
  • ionisation
  • electron affinity
  • lattice enthalpy

Sublimation (vapourisation) of solid element into gas = always endothermic
Ionisation energy = always endothermic

47
Q

X-ray Crystallography

A

X-ray crystallography allowed us to see the ‘structure’ of molecules
-allowed us to see structure of haemoglobin and DNA

X-ray diffraction

  • when x-rays are shot through a molecule at different angles, they are scattered by the regular array positions of atoms giving a DIFRACTION PATTERN
  • we can measure the positions and intensity of scattered x-rays to determine structural arrangement of molecule

X-rays can be considered as sinosoidal waves

  • when two wavelengths come close together and line up (are in phase) = constructively interference
  • the new wavelength formed will have the same wavelength + greater amplitude
  • when two wavelengths are not aligned with each other (out of phase) = deconstructive interference
  • the new wavelength = no wavelength (cancel each other out)

When considering two simple reflection of x-rays from two different layers of atoms separated by a distance of d (lattice planes):
Incident angle = reflected angle
Bottom wavelength travels further than the top wavelength, but it still in phase with the top

Bragg’s Law:
nλ = 2dsinθ
n = integer related to order of reflection
λ = lambda, wavelength of radiation
d = interplanar distance
θ = angle of incidence of the reflection
(smallest value of d can be λ/2 where visible light would be useless)

Technique can be applied to all crystalline materials:

  • proteins/DNA
  • ionic solids
  • metals
48
Q

Structure of Metals

A

Unit cell = basic repeating unit within a structure
crystal structures are made up of unit cells which are related by pure translations in 3D

In a simple cubic structure, there are 8 UNIT CELLS

49
Q

Sphere Close Packing

A

Structures of solids (metals) may be related to the structures based on the close packing of spheres

Close packing of spheres (3D) occurs when layers of spheres stack on top of each other
(when two layers stack on top of each other, a sphere from one layers lays on top of an indentation formed from three spheres of another layer)
-there are TWICE as many indentations as spheres

Spheres can stack in two ways:
Hexagonal Close Packing (HCP) = ABABA packing
third layer stacks directly on top of first layer, fourth layer stacks directly on top of second
-each sphere is surrounded by 12 SPHERES

Cubic close packing (face centred cubic) = ABCABC packing
-fourth layer is directly stacked on top of first layer (first three layers stack differently)
Each unit cell has 4 SPHERES

Bother are close packing = 72% efficiency (highest)

50
Q

Calculating number of spheres in a unit cell

A

Each sphere in a corner of a cubic unit cell = 1/8 of volume

Each sphere on the face = 1/2 of volume inside a given unit cell

(e.g. FCC, 8 x 1/8 = 1 sphere + 6 x 1/2 = 3 spheres = 4 TOTAL SPHERES

51
Q

Other types of Cubic Packing

A

Primitive cubic packing (PCC)

  • spheres only on the corners of unit cell
  • 1 sphere in a unit cell
  • not CLOSE PACKING = 52% efficiency

Body-centred cubic packing (BCC)

  • each sphere is surrounded by 8 OTHER SPHERES
  • 2 spheres per unit cell
  • sphere in centre MUST BE IDENTICAL to surrounding spheres (difference)
  • NOT closing packing (62% efficiency)
52
Q

Metal Structures

A

Most common structures in metals = body-centred cubic, cubic close packing (face centred cubic), and hexagonal close packing (ABABA)

BCC examples = Lithium, sodium, potassium (8 surrounded metal centres)

CCP examples = Aluminium, Iron, Copper, Nickel (surrounded by 12 equivalent metal centres)

HCP examples = magnesium, zinc

53
Q

Alloys

A

Alloys = mixture of elements that have metallic characteristics

Substitutional alloys = some of host atoms are replaced by similar sized metal atoms
E.g. brass = 1/3 of copper atoms are replaced by zinc

Interstitial alloys = interstitial spaces are filled by small atoms
e.g. steel = some of holes in between iron are filled by carbon atoms

54
Q

Structure of ionic solids

A

Structure of ionic solids can be easily related to close-packed or cubic packing of ions
e.g. anions would be in structure where cations are in ‘holes’ or interstitial sites

Tetrahedral site = 3 spheres and 1 sphere on top = tetrahedral hole

Octahedral site = 6 spheres = octahedral hole

There is ONE OCTAHEDRAL SITE for each sphere in close packed structures

There are TWO TETRAHEDRAL SITES for every sphere in close packed structures

If spheres were formed a close packed structure, the radius of largest sphere that can fit is:
Tetrahedral site - 0.225R
Octahedral site - 0.414 R

In stable solids, interactions between cations and anions want to be maximised
This is done by cations fitting in holes that are slightly smaller than actual size rather than larger holes
(cations touch anions; anions do not touch each other)

For a cation, it will fit in a tetrahedral site if radius of anion sphere:
Tetrahedral site - 0.225R- < r+ < 0.414 R-

Octahedral site - R+ > 0.414R (only large cations will go, if it is smaller, it will rattle around)

If cations are VERY large, the structure will switch from close packed into SIMPLE CUBIC ARRANGEMENT
Cubic hole = 0.712 R

55
Q

Structure of NaCl

A

NaCl has a face centred cubic arrangement (cubic close packing)

Octahedral sites are occupied by Na ions and surrounded by 6 chloride ions (fits into space a little too small for them)

6 Coordinate Na ions
6 Coordinate Cl ions
(Chloride ions can be in octahedral sites too as it is a larger ion)

8 x 1/8 = 1 sphere + 6 x 1/2 = 3 spheres
4 yellow spheres per unit cell (sodium)

12 sodium ions on edges = 1/4 x 12 = 3 spheres + 1 x 1 = 1 sphere
4 green spheres per unit (chloride)

Ratio of spheres is 1:1 = 1 sphere there is 1 octahedral site

(in other close packed structures, if cations are smaller, the tetrahedral sites are filled instead = 8 tetrahedral sites in every FCC unit cell
2:1 ratio)

56
Q

Structure of Zinc Sulfide (ZnS)

A

Zinc sulfide has two crystalline forms:

  • ZINC BLENDE
  • WURTZITE

The radius of zinc ion = 0.35R x radius of sulfide ion
-zinc ions occupy tetrahedral holes

ZINC BLENDE:
Cubic close packing (FCC) array of sulfide ions with half of tetrahedral sites filled with zinc ions
-Octahedral sites are empty
-each zinc ion is surrounded by 4 ions (tetrahedral environment)
-an alternative is zinc ions in a FCC arrangement but ions would be too far from each other
-coordination of zinc and sulfide ions = 4

WURTZITE:

  • hexagonal close packing of sulfide ions with half of tetrahedral sites filled with zinc ions
  • sulfide are in HCP array and zinc ions are in every second tetrahedral site
  • octahedral sites are empty
  • coordination of zinc and sulfide ions = 4
57
Q

Structure of Cesium Chloride (CsCl)

A

In CsCl, the ratio of radius between cesium to chloride is 0.93 = too large to fit in octahedral hole

CsCl arranges into PRIMITIVE CUBIC STRUCTURE (simple cubic packing)
(looks like body-centred, but spheres are different)

  • Chloride ions in simple packing and caesium ion in cubic site (or vice versa)
  • Coordination number of Caesium and chloride ions = 8
58
Q

Structure of Fluorite (CaF2)

A

Ratio of calcium ions to fluoride is 0.752 = cubic site

1 structure: Primitive (simple) cubic arrangement of fluoride ions and calcium ions in every second cubic site
(only half of cubic sites are filled as theres twice as many fluorides than calcium)

2 structure: Calcium ions in face centred cubic (cubic close packing) with fluoride ions in all tetrahedral sites (empty octahedral sites)
-fluoride is too big for tetrahedral sites, but it fits as calcium ions are pushed far apart

Coordination number of calcium = 8 (cubic site)
Coordination number of fluoride = 4 (tetrahedral site)

59
Q

Network Solids

A

Network solids are infinite covalent solids that can be 1D (chain), 2D (sheets), 3D

Example of a network solid = SiO2 (quartz in room temperature)
Quartz is based on a tetrahedral SiO4 units linked together

60
Q

Comparison of SiO4 and CO2

A

Both are group 14 oxides but the physical properties are very different

CO2 is gas in normal conditions
Silicon has a very high melting point

The difference is due to structure
CO2 = discrete covalent molecules
SiO2 = 3D covalent network

Silicon does not form good pi bonds with oxygen (p side on overlap is not efficient), it forms 4 bonds to oxygen rather than double bonds

61
Q

Silicates

A

Silicates = ANIONIC Si-O networks
-silicates makes up of earth’s rocks, minerals and soils

Silicates are OXYANIONS of silicon which are usually very large
Oxygen bridging (shared oxygen) = no charge
Oxygen non bridging (terminal oxygen) = negative charge

Silicate units are linked to form single chains = PYROXENES (e.g. jade)
Repeat unit = SiO2 (2-)

Silicate units can be linked up to form rings (e.g. beryl)
Ring formula: Si6O18 (12-)

Double chains of silicates = Mineral tremolite (form of asbestos)
Repeat unit = Si4O11 (6-)
Asbestos are hazardous because it is fibrous material (sharp edges) that is chemically very stable and NOT enzymatically degraded in the lungs (can lead to cancer)

Silicate can be linked to form SHEETS - each tetrahedron shares 3 oxygen atoms
Repeat unit: Si4O10 (4-)
E.g. Talc = bonding is mainly within the sheets, but bonds holding sheets together are relatively weak = talc is slippery

Silicates can be linked to form 3D ARRAYS
Quarts = interwoven, cross linked helical chains of silicates (all oxygens are bridging)

Molten mineral silicates usually crystallize slowly
-when melted silicates cool quickly and solidification occurs WITHOUT CRYSTALLIZATION, a supercooled liquid is formed (e.g. glass)

62
Q

Structure of Oxygen (group 16)

A

Oxygen = strong oxidising agent but most things are stable to oxidation by oxygen in normal conditions as there is a large kinetic barrier

MO description for ground state oxygen = Triplet Dioxygen

  • not very reaction with species that have all their electrons paired
  • Singlet state of oxygen is more reactive (2 unpaired electrons)

Ozone = V shaped molecule
polar, diamagnetic, strong oxidant and a toxic gas
-less stable allotrope of oxygen
-uses : bleach, sterilising water, prevents UV rays from reaching earth surface

63
Q

Structure of Sulfur (group 16)

A

Main form of sulfur is cyclic S8 molecules or chains of sulfur atoms

S2 = produced in gas phase in high temperatures (not stable)

S3 = exists in gas phase in high temperatures

64
Q

Group 16 Elements (Selenium, Tellurium, Polonium)

A

Selenium = 6 distinct forms (like sulfur), 3 structural forms are based on Se8 rings
-Most stable = metallic form = Hexagonal Selenium (helical chains with weak interactions between chains = photoconductive)

Tellurium = only 1 crystalline form with network of spirals with strong interchain interactions

Polonium = alpha polonium (has a simple primitive structure)

65
Q

Metallic Characteristics and Group 13-16 elements

A

As you go down the groups, metallic characteristics increases (directional influences decreases)

66
Q

Group 17 - Halogens

A

All halogens are diatomic with 1 covalent bond
-interactions between molecules increase as you go down the group (increased BP and MP)
-most are gases in room temperature BESIDES IODINE
(iodine = loustrous appearance = metallic character), it has weak conductance in normal conditions but HIGH CONDUCTION in high pressure

General observations:

  • pi bonding is important in influencing structure in period2/1 elements, but not as much for period 3 and higher
  • physical and chemical behaviour is related to structure (not identity of element), like black P and white P or diamond vs graphite
  • further to the right, metallic characteristics are introduced later (e.g. iodine = group 17, last element starts to show metallic characteristics then)

Chemistry 10003 (3 decks)

Brainscape helps you reach your goals faster, through stronger study habits.
© 2024 Bold Learning Solutions. Terms and Conditions