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CHEM compostella > Inorganic chemistry > Flashcards

Inorganic chemistry Flashcards

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1
Q

all types of quantum numbers

A
  1. principal (n): energy of orbital (2 in 2s1)
  2. Orbital (l): letter (s/p/d/f)
  3. Magnetic (m): direction of orbital (x/y/z)
  4. Spin (s): magnetic properties (the 2 in 3p2)
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2
Q

structure of sigma bond

A

overlapping of s orbitals along the plane of the cell nuclei –> cylindrically symmetrical

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3
Q

bonding vs antiboding

A

bonding: low energy, orbitals overlap in phase and inteact constructively –> both orbitals have the same sign

antibonding: high energy, orbitals overlap out of phase and inteact destructively

!!only if there are more electrons in the bonding than antibonding MOs will there be bonding between two atoms

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4
Q

structure of a pi bond

A

overlap of p orbitals but no longer cylindrically symmetrical along intrernuclear axis (hence not as strong as sigma bond)

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5
Q

hybridisation definition

A

Redistribution of the energy of orbitals of individual atoms to give orbitals of equivalent energy

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6
Q

sp hybridisation

A

-linear molecules with 180 bond angle
-involved mixing of an s and p orbital

!! involved in compounds containing any carbon carbon triple bond

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7
Q

sp2 hybridisation

A

-trigonal molecules with a bond angle 120
-involves mixing one s and two p orbitals

!! involved in compounds containing any carbon carbon double bond

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8
Q

sp3 hybridisation

A

-tetrahedral molecules with a bond angle 109.5
-involves mixing one s and three p orbitals

!! involved in compounds containing any carbon carbon single bond

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9
Q

hybridisation of N atoms

A
  • Undergo either sp/sp2/sp3 hybridisation
  • follows same pattern as carbon atoms in terms of bonds and hybrids

!! lone pair of electron is present in the case of sp3 hybridisation –> shape is trigonal pyramidal but with 107 bond angle

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10
Q

hybridisation of O atoms

A
  • Undergo either sp/sp2/sp3 hybridisation
  • follows same pattern as carbon atoms in terms of bonds and hybrids

!! 2 lone pairs are formed during sp3 hybridisation –> molecule has a bent shape and a bond angle 104

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11
Q

electronegativity def

A

ability of an atom in a molecule to attract electrons to itself

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12
Q

trend of electronegativity

A

increases across period
decreases down a group

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13
Q

What causes the formation of polar bonds

A

a difference in electronegativity between two atoms (larger difference means a more polar bond)

!! electrons are pulled towards electronegative atom

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14
Q

What is an electrostatic potential map

A

shows distribution of charges over a molecule (electron density map)

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15
Q

Coordination number of a transition metal complex definition

A

the number of coordiate bonds between the ligand and the central metal ion

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16
Q

types of intermolecular forces

A
  1. london dispersion
  2. permanent dipole
  3. hydrogen bonding (NOF) - has a characteristic bond length around 0.177nm
  4. Ion-dipole forces (ions and water molecules in solution)
  5. Dipole - induced dipole (permanent dipole induces another temporary dipole)
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17
Q

what is the relationship between ion volume and hydrated ion volume

A

smaller ions have the largest volume when hydrated.

larger ions have the smallest volume when hydrated

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18
Q

3 types of thermodynamic systems

A
  1. OPEN: matter and energy can be exchanged with surroundings
  2. CLOSED: energy can be exchanged but not matter
  3. ISOLATED: neither energy nor matter can be exchanged
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19
Q

First law of thermodynamics

A

energy is conserved (any energy lost by system must be gained by surroundings and vice versa)

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20
Q

What is the internal energy of a system

A

Sum of all kinetic and potential energies of the components of the system

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21
Q

Equations used for change in internal energy

A
  1. ΔE = Efinal - Einitial,
    ΔE > 0 the system has gained energy; ΔE < 0 the system has lost energy
  2. ΔE=q+w , q = heat (q>0 heat is absorbed by the system), w = work (w>0 work is done on the system)
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22
Q

Enthalpy definition and equations

A

ENTHALPY: energy of a system at constant pressure

!! H=E+PV, E is the internal energy (ΔE = q + w), P in the pressure and V is the volume

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23
Q

State function definition

A

value of a state function is dependent only on the present state of a system and not the path the system took to reach that state

EG> internal energy, enthalpy, entropy, free gibbs energy

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24
Q

Standard enthalpy change for a reaction def

A

enthalpy change of reaction when all reactants and products are in their standard states

!! Hreaction = Hformation of products - Hformation of reactants

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25
What is Hess's law
if a reaction is carried out in a series of steps, deltaH for overall reaction is equal to the sum of enthalpy changes for all individual steps
26
What is the criteria for a spontaneous reaction
1. ENTHAPLY CHANGE: highly exothermic can indicate spontaneity but it doesnt have to be the case --> not definitive 2. ENTROPY: spontaneous reactions have a positive total entropy 3. FREE GIBS ENERGY: spontaneous reactions have a negative deltaG (exergonic)
27
What has an entropy of 0
pure crystalline substances at absolute zero
28
Gibbs free energy definition
energy release by reaction that can be used to do work 1. Under standard conditions: DG 0 = enthalpy - (temp)(entropy) 2. Uner non standard conditions: DG = DG 0 + RT ln Q (where Q is the reaction quotient)
29
Colloids vs Suspensions vs Solutions
COLLOID: micro-heterogeneous mixture with larger particles (usually opaque). Particles cannot be filtered from its other components and do not settle out SUSPENSION: heterogeneous mixture with large particles suspended in liquid. Particles do not dissolve in liquid (too large). Can be filtered away from liquid or separated using centrifugation SOLUTION: homogeneous mixture (2+ substances), contain a solvent and solute, can be gaseous/liquid/solid
30
what factors affect whether a substance dissolves in another
1. natural tendence od substances to mix and spread into larger volumes when not restrained 2. Types of intermolecular forces: -solute solute and solvent solvent interactions must absorb less energy that what is released by the formation of solute solvent interactions 3. Temperature and pressure
31
Dissolution vs Dissociation
DISSOLUTION: compound dissolving into solvent anf forming a solution DISSOCIATION: ionic compound dissociating into its ions when added into solvent
32
What salts are always soluble
nitrate/ammonium and alkali metal salts
33
Henry's law definition
as the pressure increases the solubility of a gas in a liquid increases
34
how does temperature affect solubility
IONIC AND MOLECULAR SOLIDS: generally an increase in temp causes an increase in solubility GASES: increase in temo causes decrease in solubility
35
Concentration def, and different ways of expressing it (4)
CONC: amount of solute per quantity of solvent: 1. mass/volume 2. mass/mass 3. volume/volume 4. molar conc (mol/vol)
36
Colligative properties
-properties of ideal solution that depend on concentration of solute but no solute identity 1. boiling point rise 2. melting point decrease 3. change in osmotic pressure
37
Osmotic pressure definition and equation that describes it
Presure that prevents the flow of additional solvent into a solution on one side of a semipermeable membrane OP V = i n R T (where OP is osmotic pressure and i is vant hoff coefficient)
38
Isotonic vs hypotonic vs hypertonic
isotonic: concs in equilibrium - equal movement hypotonic: water moves in (sweling) -> hemolysis in RBC hypertonic: water moves out (shrinking) -> crenation in RBC
39
What isotonic solutions are used for the human body?
1. 0.9% NaCl solution 2. 5.0% glucose solution
40
Reaction rate def
change in the conc of reactants/products per unit of time
41
How to write a rate equation
aA +bB --> cC + dD: the rate equation is: RATE = k [A]^m [B]^n where k is the rate constant m is order of reactant A and n is order of reactant B
42
Factors affecting rate of a reaction
1. physcial state 2. nature of reaction 3. conc of reactants 4. reaction temp 5. pressure
43
what factors affect the value of the rate constant
temperature and the presence of a catalyst
44
Arrhenius equation
k = A e ^-Ea /RT where A is the frequence factor
45
What affects the rate law of an overall reaction?
the rate determining step in the mechanism
46
Comparison of features of lab reactions (L) vs biological reactions (B) - (4)
-L uses organic solvents, B uses aqueous solvents -L uses temp range, B uses organism temp -L uses chemical catalysts, B uses enzymes -L has catalysts with little specificity, B has enzymes with very high specificity
47
Features of an equilibrium constant equation
-products over reactants (each to the power of their mole ratio) constant K depends on temp
48
What can be deduced about an equation based on the magnitude of K
K >1 -> forwards reaction K < 1 -> backward reaction
49
reaction quotient Q definition
a number obtained by substituting reactant and products concs (or partial pressures) at any point in a reacton into an equilibrium constant expression Q>K - forms products Q=K - equilibrium Q
50
What is the difference between the reaction quotient and equilibrium constant K for a reaction
K uses the concs at equilibrium whereas Q can use concs at any time in the reaction K is the same for a reaction (only affected by temp) whereas Q changes throughout the time the reaction takes place
51
what is the relationship between gibbs energy and the equilibrium constant K
lnK = -DG 0 / RT (because at equilibrium DG is 0)
52
Le Chateliers principle def
if a system at equilibrium is disturbed by a change in conc/pressure/temp, the system will shift the position of equilibrium to counter the effect of the disturbance
53
What is the equilibrium constant for a net reaction made of 2+ reactions
the product of the equilibrium constants for each individual reaction
54
A substance that conducts electricity in aqueous solution is known as an
electrolyte
55
definition of concentration in parts per million
mass of solute (g)/ mass of solution (g) x 10^6 OR volume of solute (mL) / volume of solution (mL) x 10^6
56
heat of fusion def
The amount of energy needed to melt 1g of a substance
57
What is the vant hoff coefficient for glucose and urea
both have i = 1
58
calories to joules conversion
1 kcal = 4186 Joules
59
what is the sign of bond dissociation energies
positive and endothermic (bcos breaking of bonds requires energy)
60
bond dissociation energy def
heat change of reaction for breaking a covalent bond by equally dividing the electrons within the bond
61
does the K of a reaction alone give us enough info to predict rate of a reaction
no, only which reaction is favoured
62
what does a negative pH indicate
that the [H+] ions is more than 1moldm-3
63
definition of pH
pH = -log10[H+]
64
strong vs weak acid
strong acid fully dissociates into ions and weak acid does not show full dissociation (most of the solution is made up of its molecules )
65
Ka values and what they say about the acid
The stronger the acid the higher the value of Ka
66
pKa values and what they say about the acid
pKa = -log10(Ka) lower pKa = stronger acid pKa = pH at half equivalence (best buffering capacity, where the conc of salt and conjugate base are equal_
67
bronsted lowry def of acids and bases
ACID: proton donor BASES: proton receiver (lone pair)
68
Kw definition
The product of the concentrations of H3O+ and OH− in water or an aqueous solution—symbolized by Kw and equal to 1 × 10−14
69
name for compound that can act both as a base and acid
amphoteric -- water
70
action of a buffer solution upon addition with H+ and OH- ions
1. WHEN H+ IONS ARE ADDED: react with the weak base to form the weak acid 2. WHEN OH- IONS ARE ADDED: react with the weak acid and form the weak base
71
buffer in the blood made of
carbonic acid and bicarbonate ions
72
standard electron potential def
the emf produced when a standard half cell is connected to a hydrogen half cell at standard conditions
73
potential diff of a cell equation
Ecell of cathode - E cell at anode
74
E cell value purpose
the more positive the e cell value the better reduction occurs (hence better oxidising agent) the cell with the higher E cell acts as the positive (cathode where reduction happens)
75
meaning of a positive overall e cell value
reaction is feasible (in that direction)
76
what is the gibbs free energy for a cell reaction in standard conditions
DG = -nF(overall Ecell) n= number of electrons exchanged F= faraday constant
77
nerst equation
E= Eo− 0.0592V/n (log10Q) E = potential diff Eo = at standard conditions Q = reaction quotient
78
relationship between e cell and K equilibrium
directly proportional
79
definition of Kb
Kw/Ka
80
relationship between pH and pKa in buffers
pH = pKa+ log {[CB] / [WA]} where the CB is conjugate base and WA is the weak acid
81
if the equilibrium constant Kc is greater than 1 for a given reaction, predict the signs of ΔG° and E° at the same temperature.
K>1 DG<0 E>O

CHEM compostella (3 decks)

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