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CHEM compostella > Inorganic chemistry > Flashcards

Inorganic chemistry Flashcards

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
Q

What is Hess’s law

A

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

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26
Q

What is the criteria for a spontaneous reaction

A
  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)
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27
Q

What has an entropy of 0

A

pure crystalline substances at absolute zero

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28
Q

Gibbs free energy definition

A

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)

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29
Q

Colloids vs Suspensions vs Solutions

A

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

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30
Q

what factors affect whether a substance dissolves in another

A
  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
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31
Q

Dissolution vs Dissociation

A

DISSOLUTION: compound dissolving into solvent anf forming a solution

DISSOCIATION: ionic compound dissociating into its ions when added into solvent

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32
Q

What salts are always soluble

A

nitrate/ammonium and alkali metal salts

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33
Q

Henry’s law definition

A

as the pressure increases the solubility of a gas in a liquid increases

34
Q

how does temperature affect solubility

A

IONIC AND MOLECULAR SOLIDS: generally an increase in temp causes an increase in solubility

GASES: increase in temo causes decrease in solubility

35
Q

Concentration def, and different ways of expressing it (4)

A

CONC: amount of solute per quantity of solvent:

  1. mass/volume
  2. mass/mass
  3. volume/volume
  4. molar conc (mol/vol)
36
Q

Colligative properties

A

-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
Q

Osmotic pressure definition and equation that describes it

A

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
Q

Isotonic vs hypotonic vs hypertonic

A

isotonic: concs in equilibrium - equal movement

hypotonic: water moves in (sweling) -> hemolysis in RBC

hypertonic: water moves out (shrinking) -> crenation in RBC

39
Q

What isotonic solutions are used for the human body?

A
  1. 0.9% NaCl solution
  2. 5.0% glucose solution
40
Q

Reaction rate def

A

change in the conc of reactants/products per unit of time

41
Q

How to write a rate equation

A

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
Q

Factors affecting rate of a reaction

A
  1. physcial state
  2. nature of reaction
  3. conc of reactants
  4. reaction temp
  5. pressure
43
Q

what factors affect the value of the rate constant

A

temperature and the presence of a catalyst

44
Q

Arrhenius equation

A

k = A e ^-Ea /RT

where A is the frequence factor

45
Q

What affects the rate law of an overall reaction?

A

the rate determining step in the mechanism

46
Q

Comparison of features of lab reactions (L) vs biological reactions (B) - (4)

A

-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
Q

Features of an equilibrium constant equation

A

-products over reactants (each to the power of their mole ratio)

constant K depends on temp

48
Q

What can be deduced about an equation based on the magnitude of K

A

K >1 -> forwards reaction
K < 1 -> backward reaction

49
Q

reaction quotient Q definition

A

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<K - reactants formed

50
Q

What is the difference between the reaction quotient and equilibrium constant K for a reaction

A

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
Q

what is the relationship between gibbs energy and the equilibrium constant K

A

lnK = -DG 0 / RT

(because at equilibrium DG is 0)

52
Q

Le Chateliers principle def

A

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
Q

What is the equilibrium constant for a net reaction made of 2+ reactions

A

the product of the equilibrium constants for each individual reaction

54
Q

A substance that conducts electricity in aqueous solution is known as an

A

electrolyte

55
Q

definition of concentration in parts per million

A

mass of solute (g)/ mass of solution (g) x 10^6

OR

volume of solute (mL) / volume of solution (mL) x 10^6

56
Q

heat of fusion def

A

The amount of energy needed to melt 1g of a substance

57
Q

What is the vant hoff coefficient for glucose and urea

A

both have i = 1

58
Q

calories to joules conversion

A

1 kcal = 4186 Joules

59
Q

what is the sign of bond dissociation energies

A

positive and endothermic (bcos breaking of bonds requires energy)

60
Q

bond dissociation energy def

A

heat change of reaction for breaking a covalent bond by equally dividing the electrons within the bond

61
Q

does the K of a reaction alone give us enough info to predict rate of a reaction

A

no, only which reaction is favoured

62
Q

what does a negative pH indicate

A

that the [H+] ions is more than 1moldm-3

63
Q

definition of pH

A

pH = -log10[H+]

64
Q

strong vs weak acid

A

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
Q

Ka values and what they say about the acid

A

The stronger the acid the higher the value of Ka

66
Q

pKa values and what they say about the acid

A

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
Q

bronsted lowry def of acids and bases

A

ACID: proton donor

BASES: proton receiver
(lone pair)

68
Q

Kw definition

A

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
Q

name for compound that can act both as a base and acid

A

amphoteric – water

70
Q

action of a buffer solution upon addition with H+ and OH- ions

A
  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
Q

buffer in the blood made of

A

carbonic acid and bicarbonate ions

72
Q

standard electron potential def

A

the emf produced when a standard half cell is connected to a hydrogen half cell at standard conditions

73
Q

potential diff of a cell equation

A

Ecell of cathode - E cell at anode

74
Q

E cell value purpose

A

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
Q

meaning of a positive overall e cell value

A

reaction is feasible (in that direction)

76
Q

what is the gibbs free energy for a cell reaction in standard conditions

A

DG = -nF(overall Ecell)

n= number of electrons exchanged
F= faraday constant

77
Q

nerst equation

A

E= Eo− 0.0592V/n (log10Q)

E = potential diff
Eo = at standard conditions
Q = reaction quotient

78
Q

relationship between e cell and K equilibrium

A

directly proportional

79
Q

definition of Kb

A

Kw/Ka

80
Q

relationship between pH and pKa in buffers

A

pH = pKa+ log {[CB] / [WA]}

where the CB is conjugate base
and WA is the weak acid

81
Q

if the equilibrium constant Kc is greater than 1 for a given reaction, predict the signs of ΔG° and E° at the same temperature.

A

K>1
DG<0
E>O

CHEM compostella (3 decks)

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