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Chemistry definitions/formulas/sample answers > Chem (not finished yet) > Flashcards

Chem (not finished yet) Flashcards

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1
Q

Definition of electronegativity

A

The ability of an atom to attract bonding electrons towards itself.

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2
Q

is the C-H bond polar/non-polar?

A

C-H bond is polar for redox purposes
C-H bond is non-polar for organic/IMF purposes (all other purposes basically)

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3
Q

model answer when comparing between non-polar/polar molecules

A

(without endpoint given)
- the one with the larger, more polarisable cloud size [estimated using Mr] has the higher BP due to its intermolecular dispersion forces requiring more energy to overcome etc.

(with endpoint given, e.g. why is the BP of H2O > H2S)
- if both molecules are polar, they want you to compare intermolecular dispersion forces
- H2O has h-bonding etc..
- if both molecules are non-polar, compare electron cloud size and polarisability

only talk about electron cloud size and dipole-dipole interactions in ISOLATION, never together. To be safe you can just use intermolecular dispersion forces if not sure

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4
Q

what is VESPR theory?

A
  • EP-EP repulsion > EP-BP repulsion > BP-BP repulsion
  • EPs and BPs space themselves out to minimise repulsion
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5
Q

What makes a solvent soluble?

A

If the new bonds that can be formed btwn the solute and solvent are stronger than the originial bonds btwn solute molecules and solvent molecules, the solute dissolves in the solvent.
(Basically if the new bonds formed are stronger)

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6
Q

Covalent vs ionic vs metallic character

A

Covalent character:
- Maximum covalent character occurs when both atoms have high E-, and E- difference is LOW
- Greater diff in E- = less covalent character

Ionic character:
- increases as diff btwn E- increases

Metallic character:
- Maximum metallic character occurs when both atoms have low E-, and E- difference is LOW

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7
Q

Definition of volatile/volatility

A

volatile = ability to change into gas easily

NOT gain/lose electron

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8
Q

Ability to conduct electricity model answer (compare Na and NaI)

A
  • Na has a metallic structure
  • NaI has a giant ionic lattice structure
  • in molten state, both possess MOBILE CHARGE CARRIERS in the form of delocalised electrons (Na) and mobile Na+ and I- ions (NaI) = both structure are electrically conductivie
  • in solid state, Na still has delocalised electrons…
  • in solid state, NaI’s ions are locked in a crystal lattice structure and cannot move..
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9
Q

Sample answer: explain why ___ is polar/nonpolar?

A

e.g:
CO2 is non-polar as it is symmetrical about all bond axes, and so dipole movements cancel out

H2s is polar as it is not symmetrical about all bond axes, and so dipole movements do not cancel out

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10
Q

Practice drawing hydrogen bonds (use NH3 as an examples)

A

!!

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11
Q

Can benzene C6H6 conduct electricity due to delocalised electrons?

A

no (delocalised electrons in organic arene rings cannot move outside the ring)= benzene cannot conduct electricity

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12
Q

Definition/conditions needed for resonance

A
  • When there are 3 or more consecutive atoms with pi bonds/lone pairs
  • A pair of electrons can move btwn the atoms to form at least 2 different forms with the same stability (formal charge)

Result: bond lengths are AVERAGED out

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13
Q

Definition of sigma/pi bonds

A

pi bond:
- formed btwn the side-on overlap of p-orbitals

sigma bond:
- formed btwn the head-on overlap of any 2 kinds of orbitals

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14
Q

Definition of hybridisation + table to rmbr

A

Hybridisation is the combination of atomic orbitals from at least 2 different subshells to form new hybrid orbitals of equal energy, which maximises the bonding potential (no. of bonds an atom can form) of an atom

no. of pi bonds. | hybridisation
0 | sp3
1 | sp2
2 | sp

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15
Q

Relationship btwn cm3 and dm3 :(

A

dm3 is LARGER than cm3 omfg

convert cm3 to dm3 = divide by 1000

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16
Q

Draw the 3 tables for stoichiometry

A

Mass
Mol | Mr

Vol of gas (dm3)
Mol | Stp = 22.7 (273K), 1 atm
Rtp = 24.0 (297K), 1 atm

                                 Mol  Vol of solution (dm3)   |   Conc of solution (mol/dm3)

“standard conditions” refer to RTP, not STP

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17
Q

molecular formula

A

molecular mr/emperical mr*empirical formula

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18
Q

Finding empirical formula table

A

% (or g)
Mr
mol
ratio

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19
Q

Avogadros constant

A

6.02*10^23

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20
Q

Oxidising agents examples

A

KMNO4, K2Cr2O7

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21
Q

Reaction kinetics: what is overall order?

A

Sum of coefficients of chemicals involved in the RATE-DETERMINING step

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22
Q

How to measure rate of rxn using an actual set-up if GAS IS PRODUCED?

A
  • carry out the reaction in a chamber attached to a gas syringe and monitor how the vol. of the syringe changes
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23
Q

Factors that affect rate of rxn

A

increase in ____ increase rate of rxn:
- temp (increased avg kinetic energy of rxnt particles = increased avg velocity = increased rate of collisions AND incresed proportion of collisions with energy greater than/equal to EA = increase rate of rxn)

  • particle size (increase rate of collisions = increase rate of rxn)
  • concentration/volume (increase rate of collisions = increase rate of rxn)
  • catalyst (increased proportion of particles w energy greater to/equal to EA = increased in proportion of collisions which are effective = increase rate of rxn)
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24
Q

Arrenhius equation

A

k = Ae^(-Ea/RT)

where:
k = rate constant
A = Arrenhius constant
Ea = activation energy
R = 8.31 J/mol
T = temperature in K

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25
Q

How to find Ea from graph and Arrenhius eqn

A

Graph:
Ea = -(gradient)(R)

Arrenhius eqn:
Ea = -RT( lnk-lnA)

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26
Q

Eqn of lattice energy

A

|LE|∝ (q+q-)/(r+ + r-)

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27
Q

How doe length of carbon chains affect solubility/BP/MP

A

Longer carbon chains = stronger intermolecular dispersion forces = higher MP/BP but LOWER solubility in water

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28
Q

Giant molecular structure examples

A

Diamond
Graphite
SiO2
SiCL
C60 fullerine

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29
Q

Does C=C double bond decolourise bromine water

A

YES!!

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30
Q

Activation energy definition

A

The minimum amt of energy that reactant particles must possess in order to successfully react upon collision

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31
Q

Rate of rxn definition

A

The rate at which reactants are consumed and products are formed in a chemical reaction (typically measured as moldm-3s-1

32
Q

Are all Na salts soluble in water

A

YES!!

33
Q

Formulas to find total energy (heat) produced

A

Q = mc∆T, Q = n∆H

Where:
Q = total energy produced
m = mass (g)
c = specific heat capacity (water is 4.18)
∆T = change in temp in K
n = mol of reactant
∆H = change in temp with respect to reactant

g converts to J
kg converts to kJ

34
Q

Using the graph to find the actual highest temp during a reaction (temperature/time)

A
  • Draw a straight line upwards from when the curve begins
  • Draw another straight line following the downward gradient of the curve
  • the intersection btwn these two lines is the actual highest temp
35
Q

Standard enthalpy change of formation definition

A

Enthalpy change when 1 mol of a compound is formed from its constituent elements in their standard states under standard conditions

+/-, per mol of product

36
Q

Standard enthalpy change of combustion definition

A

Enthalpy change when 1 mol of a substance is completely burned in oxygen under standard conditions

-, per mol of reactant

37
Q

Standard enthalpy change of neutralisation definition

A

Enthalpy change when 1 mol of water is formed in the neutralisation between an acid and alkali under STP

-, per mole of H2O formed

38
Q

Standard enthalpy change of atomisation definition (element)

A

Enthalpy change when an element is converted into 1 mol of gaseous free atoms under STP

+, per mol of product

39
Q

Standard enthalpy change of atomisation definition (compound)

A

Enthalpy change when 1 mol of gaseous molecules are converted to their constituent gaseous atoms under standard conditions

+, per mol of reactant

40
Q

Standard enthalpy change of hydration definition

A

Enthalpy change when 1 mol of gaseous ions is dissolved in large amounts of water to form 1 mol of aqueous ions under STP

-, per mol of reactants

41
Q

Standard enthalpy change of solution definition

A

Enthalpy change when 1 mol of a solid substance dissolves in such a large volume of solvent that addition of more solvent produces no further heat exchange under STP

+/-, per mole of reactant

42
Q

Lattice energy definition

A

Enthalpy absorbed when 1 mol of an ionic crystalline solid is converted to its separate gaseous ions under STP

+, per mole of reactant

43
Q

Electron Affinity definition

A

Enthalpy change measured (in kJ mol-1) when 1 mol of electrons is added to 1 mol of atoms and ions in the gaseous state

First ea= -ve, exothermic
Second ea = +ve, endothermic

44
Q

Ionisation energy definition

A

Energy required to remove one mol of electrons from one mol of gaseous atoms to form one mol of singly positively charged cations under STP

+ve

45
Q

Bond energy definition

A

Energy required to break 1 mol of a covalent bond between 2 atoms in a gaseous state

Bond formed = - bond energy
Bond broken = + bond energy

46
Q

1/2 life properties

A

TIME TAKEN (it is a DURATION, not time from 0) to 1/2 the concentration of a given reactant
- 1/2 life of a reactant is INDEPENDENT of its concentration

47
Q

Standard state of S and P

A

S8 and P4

48
Q

Explain why theoretical LE differs from experimental LE

A
  • LE values are only accurate if a substance is purely ionic. The more covalent character a substance has in its bonding, the more the experimental and theoretical LE values will defer
  • Hence, as ____ has non-negligible covalent character, its theoretical and experimental LE values differ
49
Q

Lewis acid/bases are:

A

Lewis acid: electron pair acceptor
Lewis bases: electron pair donators

50
Q

Gibbs free energy eqn

A

∆G = ∆H - T∆S

where
∆G and ∆H are in kJ
REMEMBER, ∆S is usually in given in J, so MUST CONVERT TO kJ by dividing by 1000 first!!!!!!!

reaction is spontaneous if ∆G<0, not spontaneous if ∆G>0, and is at equilibrium at ∆G = 0

51
Q

Table of ∆H and ∆S to find out easily find spontaneity of a reaction

A

, ∆H = +ve ∆H = ive

∆S = +ve at high temp. all temp

∆S = -ve at no temp at low temp

52
Q

Explain why ∆Hf of an element (e.g. H2) is 0?

A

∆Hf only applies to compounds, H2 is an element

53
Q

Why is the rxn spontaneous/not spontaneous?

A

Spontaneous:
- ∆G is -ve, meaning that the reaction results in the loss of Gibbs free energy in the surroundings.
- Hence the rxn is spontaneous, meaning that it should occur WITHOUT ADDITIONAL ENERGY BEYOND E.A

vice versa for non-spontaneous

54
Q

When finding/comparing ∆S change,

A

Look at the change in the no. of mols of GASES first!!!!

55
Q

Relative atomic mass definition

A

The ratio of the mass of an atom of an element to 1/12 the mass of an atom of carbon-12

56
Q

Relative molecular mass definition

A

The average mass of a molecule divided by 1/12 the mass of a carbon-12 atom

57
Q

Ideal gas law equation (can be used in qns involving pressure, volume, temp, to find the mol of a reactant)

A

PV = nRT

where:
P = pressure (pa)
V = volume (m3!!!)
n = mol of reactant
R = 8.31
T = temperature in K

58
Q

Primary/secondary/tertiary alcohols

A

Primary: 2 or more H atoms bonded to C
Secondary: 1 H atom bonded to C
Tertiary: 0 H atoms bonded to C

59
Q

Properties of primary alcohols

A

PRIMARY ALCOHOLS OXIDISE TO FORM ALDEHYDES WHICH FURTHER OXIDISE TO GIVE CARBOXYLIC ACIDS

To form aldehydes: heat with immediate distilation
To form carboxylic acids: heat under reflux

60
Q

How does changes in state affect ∆H? e.g. (g) to (l)

A

∆H would be exothermic. Condensation is exothermic as it involves the formation of stronger intermolecular forces. Hence by Hess’s law the formation of (l) is more exothermic than the formation of (g)

61
Q

Bronsted-lowry acid

A

Proton (H+) donator

62
Q

Bronsted-lowry base

A

Proton (H+) acceptors

63
Q

redox equations: approach!

A
  1. write down the individual redox equations, referring to TABLE 24. (e.g. for the redox equation btwn MNO4- and Fe2+, write down the indivual eqns first)
  2. equate the e- coefficients and combine the equations!!!!!
64
Q

Celcius and Kelvin DO NOT have 1:1 ratio!!

A

stupid

65
Q

r/s btwn gdm-3 and moldm-3

A

mr = gdm-3/moldm-3

66
Q

formula for combustion

A

CxHy + (x+y/4) O2 = xCO2 + y/2 H2O

67
Q

Ag assumed charge

A

+1 e.g AgCl

68
Q

specific heat capacity vs heat capacity

A

specific: small c, need to use mass in change in Q = mc change in temp

heat capacity: big C, no need to use mass so change in Q = C change in temp

69
Q

Definition of S

A

S, entropy, relates to the degree of disorder in a system

70
Q

Definition of atomic radius

A

half of the distance between the nuclei of neighbouring atoms in the pure element

71
Q

Explain why there is a large difference between theoretical and experimental lattice enthalpies

A

When there is a large difference btwn theoretical and experimental lattice enthalpies, it implies that there is more covalent character in the bonding of the substance

72
Q

Definition of half life

A

time taken to halve the conc. of a given reactant

consistent half life ONLY for first order reactions

73
Q

Why do gases deviate from ideal gas law at high temp?

A

molecules have finite volume`

74
Q

2 pronged argument draw it out!!

A
75
Q

maximum no. of electrons that can occupy a p orbital?

A

2!!! NOT 6!!

76
Q
A

Chemistry definitions/formulas/sample answers (1 decks)

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