3.1 & 2 Flashcards
(127 cards)
1
Q
How is the periodic table arranged
A
By increasing atomic no.
In periods showing repeating trends in physical and chemical properties
In groups having similar chemical properties
2
Q
First ionisation energy
A
Removal of 1 mole of electrons from each atom in 1 mole of gaseous atoms
3
Q
Metallic bonding
A
Strong electrostatic attraction between cations and delocalised electrons
4
Q
Solid giant covalent lattices
A
Networks of atoms bonded by strong covalent bonds
5
Q
Atomic radii across a period
A
Decreases
More attraction to nucleus
Greater no. of protons
6
Q
When is the ionisation energy higher
A
With a smaller radius (higher nuclear charge)
Greater attraction
More energy needed
7
Q
How do inner quantum shells affect the IE
A
Inner quantum shells shield the electrons from the attractive force of the nucleus
The more inner quantum shells
The greater the shielding effect
The lower the IE
8
Q
When a molecule sublimates what bonds are broken
A
The intermolecular forces not the covalent bonds
9
Q
Why is Mg harder than Na
A
No. of delocalised electrons
Charge on cations
Cations are smaller, so the charge density is even greater
10
Q
Why is there an increase in mp and bp of metals from Gp 1 to Gp3
A
Increased strength of metallic bonding: The charge on the positive ion is greater More delocalised ions (1 in 1+ vs 3 in 3+) Smaller ions (greater charge density)
11
Q
Properties of metals
A
Malleable
Ductile
High mp and bp
Good electrical conductors
12
Q
Solubility of giant ionic lattices
A
Ionic lattices dissolve in polar solvents like water
Cations attract delta negative charges on the O molecule in H2O
Anions attract delta positive charges on the H molecules in H2O
13
Q
Solubility of simple molecular lattices
A
London forces can form between molecules and non-polar solvents (hexane), weakening the lattice
14
Q
Group 2 characteristics
A
Fairly high melting and boiling points
Low densities
15
Q
Group 2 reactions
A
2 Mg (s) + O2 (g) —> 2MgO (s) Ca (s) + 2 H20 (l) —> Ca(OH)2 + H2
16
Q
Second IE
A
X (g) —> X^+ (g) and e^-
X^+ (g) —> X^2+ and e^-
17
Q
Why does reactivity increase down Group 2
A
Further from nucleus
More shielding
1st and 2nd IE decreases
18
Q
Why does solubility and pH increase down a group
A
Mg(OH)2 —> Mg2+ and 2 OH-
Dissociation causes more hydroxide ions to be released —> more alkaline solution
19
Q
Down Group 7
A
No. of full shells increase
Outermost electrons further from nucleus
Atomic radius increases
20
Q
Forces in Group 7 molecules
A
London forces
More electrons = stronger London forces
21
Q
What do halogens react to form
A
Halide ions
22
Q
Are halogens reduced or oxidised
A
Reduced
Oxidising agent
23
Q
Why does reactivity decrease down Group 7
A
Nuclear charge increases down a group
Atomic radius increases
Shielding increases (outweighs nuclear atraction)
Attraction between nucleus and electron decreases
Halogen less readily accepts an electron
24
Q
Do halide ions get oxidised or reduced
A
Oxidised
25
Why is iodide the most powerful reducing agent
It wants to lose electrons
26
How does chlorine react with water
To form an acidic solution containing HCl and HOCl (chloric (I) acid)
27
Risks of chlorinating water
Cl2 gas is toxic
| Chlorinated hydrocarbons may form (carcinogens)
28
Benefits of chlorinating water
Effectively kills many harmful microorganisms
| Reduction in cholera and typhoid rates
29
Bleach
Formed when Cl2 is added to cold dilute NaOH solution
Active ingredient is sodium chlorate (NaOCl)
Cl2 (aq) + 2 NaOH (aq) —> NaOCl (aq) + Nacl (aq) + H2O (l)
30
Tests for halide ions
AgNO3 used to test
| Silver halides are insoluble
31
Cl- (aq) and Ag+ (aq) —>
AgCl (s)
| White precipitate
32
Br- (aq) and Ag+ (aq) —>
AgBr
| Cream precipitate
33
I- (aq) and Ag+ (aq) —>
AgI
| Yellow precipitate
34
What does AgCl dissolve in
Dilute ammonia
35
AgBr dissolves in
Concentrated ammonia
36
Does AgI dissolve
No
37
Colours with cyclohexane
Cl2 - pale green
Br2 - orange
I2 - violet
38
Order of identifying ions
Carbonate
Sultate
Halide
39
Carbonate test
Dilute nitric acid
Bubbles of gas, effervescence if positive
If you bubble CO2 through lime water it goes from clear to cloudy
40
Sulfate test
Add Ba^2+ (Ba(NO3)2)
| Forms white precipitate if positive - barium sulfate
41
Halide test
Dissolve halide in H2O and add AgNO3
AgX precipitate colour if positive
Also test if dissolves in NH3
42
Ammonium test
Add warm NaOH and place on damp litmus paper
| Turns blue
43
Why is the order of the ion tests important
BaCO3 (s) and Ag2SO4 (s) are both insoluble white precipitates but only BaCO3 would give off bubbles of gas so you’d automatically know
44
If you have a mixture of ions:
Do tests in same order
Keep adding HNO3 until reaction stops
Add excess of barium nitrate to ensure all SO4^2- form a precipitate - which can be filtered
Then test filtrate for halide ions
45
Why is dilute nitric acid used in the carbonate test
It does not form any precipitates whereas HCl and H2SO4 do
46
Direct measurement of enthalpy change
```
Q = m * c * delta-t
Delta-H = -Q/n
```
Use n of moles that is not in excess
47
Hess’ law
Enthalpy is a state function, depends only on state
If a reaction can take more than one route and the initial and final conditions are the same, the total enthalpy change is the same
48
Catalyst
A substance that increases the rate of a chemical reaction without being used up in the overall reaction.
Does this by providing an alternative route with a lower activation energy
49
Homogeneous catalysis
A reaction in which the catalyst and reactants are in the same physical state usually aq or g
50
Heterogeneous catalysis
A reaction in which the catalyst and the reactants are in diff. states. The reactant is usually a gas and the catalyst, a solid
51
Dynamic equilibrium
The equilibrium that exists in a closed system when the rate of the forward reaction is equal to the rate of the reverse reaction and the conc. don’t change
52
Le Chateliers principle
When a system in dynamic equilibrium is subjected to a external change, the system readjusts itself to minimise the effect of the change and to restore equilibrium
53
What is Ca(OH)2 used for
Neutralising acid soils
54
What are Mg(OH)2 and CaCO3 used as
‘Antacids’ in treating indigestion
55
Delta H
Reactants - products
56
-ve delta h
Exothermic
57
+ve delta h
Endothermic
58
Activation energy
Minimum energy required to start a reaction
59
Enthalpy change of formation
Formation of 1 mol of a compound from its constituent elements, in standard conditions
Products and reactants in their standard form
60
Standard enthalpy of neutralisation
Formation of 1 mol of water from neutralisation
61
Standard conditions
298K
| 100 kPa
62
Collision theory
For a reaction to take place, particles must collide in the correct orientation with sufficient energy
63
Why does using a catalyst increase the rate
It provides an alternative route with lower Ea so a greater proportion of collisions have an energy > Ea
64
Increasing the temperature increases the rate because
A greater proportion of molecules have an energy > Ea and there are more frequent successful collisions
65
Changing the pressure/conc. changes the rate because
There are more/less particles per unit volume and therefore there are more/less collisions
66
Enthalpy change of formation
Formation of 1 mol of a substance from its constituent elements under standard conditions and states
67
Enthalpy change of combustion
Complete combustion of 1 mole of a substance with excess O2 under standard conditions and states
68
Enthalpy change of neutralisation
Formation of 1 mol of water from neutralisation under standard conditions and states
69
How to measure rate of reaction from a graph
Drawing a tangent at said point and calculating the gradient
70
Factors affecting rate of reaction
Temperature
Concentration
Surface area
With or without catalysts
71
Effect of increasing surface area
All particles have exposed surfaces to react with each other; more frequent collisions
72
What does the highest point of the Boltzmann curve shows
Most probable energy
73
What is the total area under the Boltzmann distribution equal to
The no. of molecules
74
If the forward reaction is exothermic, how does increased temp affect the equilibrium
Shifts to the LHS
75
If the forward reaction is exothermic, how does decreased temp affect the equilibrium
Shifts it to RHS
76
How does increasing pressure affect the equilibrium
Shifts to side with least moles of gas
77
How does decreasing pressure affect equilibrium
Shifts to side with most moles of gas
78
How does the equilibrium shift if the conc of reactants increases
Shifts to the RHS
79
How does the equilibrium shift if the conc. of products increases
Shifts to the LHS
80
If the forward reaction is endothermic, how does increased temp affect the equilibrium
Shifts to the RHS
81
If the forward reaction is endothermic, how does decreased temp affect the equilibrium
Shifts to the LHS
82
What is Kc a measure of
The conc. of products and reactants at equilibrium
83
How is Kc calculated
Kc = [C]^c [D]^d/ [A]^a [B]^b
84
If Kc >> 1
More products are formed at equilibrium
85
If Kc << 1
More reactants formed at equilibrium
86
If Kc = 1
Equal conc. of reactants and products
87
Why do solid giant ionic structures not conduct electricity
They have ions fixed in position by ionic bonds and so don’t conduct
88
Role of a catalyst
Increasing rate of reaction without being used up by the overall reaction
Providing alternative route with lower Ea
89
Homogenous catalysts
Has same physical state as the reactants. Reacts with reactants to form intermediate. This then breaks down to give the product and regenerates the catalyst
90
Heterogenous catalyst
Has diff. physical state to reactants. Reactant molecules are absorbed onto the surface of the catalyst where the reaction take place, activation energy is lowered and the product molecules leave by desorption
91
Catalyst for Haber process
Solid Iron
92
Catalyst for reforming
Pt (s) or Rh (s)
93
Catalyst in hydrogenation of alkenes
Solid Ni
94
Industrial preparation of ethanol
C6H12O6 —> 2 C5H5OH + 2 CO2
95
Benefits of catalysts
Increased sustainability:
| Lowering temps and reducing energy demand from combustion of fossil fuels with resulting reduction in CO2 emissions
96
Thermal decomposition of Group 2 carbonate
CaCO3 —> CaO + CO2
97
Periodicity
Repeating patterns of properties across diff. periods
98
Why do solid ionic lattices not conduct electricity but molten ones do
The ions are fixed in place by ionic bonds when solid but when molten the ions are mobile
99
Which Group 2 carbonate decomposes at the highest temp
BaCO3
100
Are giant metallic structures soluble in water
No
101
Bonding in gaseous hydrogen halides
Mainly covalent with an increasing tendency towards ionic as you go up the group
102
Which particles are attracted in metallic bonding
Cations and delocalised electrons
103
Which halogen most readily forms 1- ions
Fluorine
104
Percentage uncertainty
No. of readings * uncertainty/ (quantity measured) * 100
105
Limitation in enthalpy pag
Fuel evaporates
Heat loss to surroundings
Incomplete combustion.
106
Sources of error in titrations
Not accurately weighing mass of solid - use more precise balance
Not all acid gets transferred to volumetric flask - rinse out beaker and add rinsings to volumetric flask
Insufficient mixing of solution - invert several times
Burette not rinsed - rinse with acid solution before use
107
Which halogen most readily forms 1- ions
Fluorine
108
NO and CO in a catalytic converters
2NO + 2CO —> 2CO2 + N2
CO and NO adsorbed onto surface
CO2 and N2 made, lowered activation energy
Products desorb from the surface
109
General trend in mp and bp from Gp 1 to Gp 4
Increases
110
Why is there an increase in mp and bp between Gp 3 - Gp 4
From Al to Si; giant metallic lattice to giant covalent lattice
Covalent bonds are stronger than metallic bonds in the same period
111
Why is there a decrease in mp and bp between Gp 4 to Gp 7
Go from giant covalent lattice to simple molecular lattice w/ weak IMF (LF) and molecules get smaller e.g. P4 and Cl2
112
Simple molecular lattice
Weak IMF (LF) between molecules
113
The bigger the molecules ...
The more electrons
Stronger induced dipole-dipole forces
Higher mp
114
How does repulsion between electrons in orbitals affect the IE
Less repulsion leads to larger IE
115
Why doesn't a catalyst change the position of eqm
Increases the forward and backward reaction by the same amount
116
Enthalpy change of reaction
No. of moles of reactants specified in the eqn react together
117
Avg. bond enthalpy
One mole of bonds breaking
118
Why may the actual bond enthalpy be diff to the calculated value
Bonds have diff strengths in diff environments
119
Why do industrial manufacturing processes use catalysts
```
Higher atom economy
Reduce CO2 emission
Enable reactions to occur w/ more specificity (correct stereoisomer)
Reactions can occur at lower temp
Saves energy costs
```
120
What makes a reaction exothermic
Bond breaking absorbs energy
Bond forming releases energy
More energy released than absorbed
121
Endothermic
Energy enters system from surrounding
122
Energy profile diagram for endothermic reaction
Products above reactants
| Ea and delta H going up to products
123
How is a dynamic eqm reached
Rate of forward reaction slows down and rate of backwards reaction speeds up
124
Ionic eqn of carbonate test
CO3 2- + 2H+ ---> CO2 + H2O
125
Directly measuring enthalpy change of solution/ neutralisation
Measure out sol using a gradated pippette into a polystyrene cup
Weigh out excess of solid
Place thermometer into cup; stir and record temp every 30s for X mins
At X+1 mins add solid to cup
Gently stir and record temp every 30s until 10mins
Draw cooling curve and extrapolate
126
Why must we extrapolate cooling curves to get an accurate value for delta T
Reaction isn't instantaneous
| Heat is lost as mixture heats up
127
Why may enthalpy change of neut/sol be diff from accepted value
No lid is used - heat loss
Spp heat capacity may be inaccurate
Energy absorbed by polystyrene
