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Chemistry A-level > Module 3 content > Flashcards

Module 3 content Flashcards

(61 cards)

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1
Q

How is the periodic table arranged?

A

By increasing atomic/proton number
In periods showing trends in chemical/physical properties

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2
Q

Trends across the periods in table

A

Period 2,3= s and p fill
4 onwards= d shell fills,

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3
Q

Trends down the group

A

Same number of atoms in outer shell
Atoms with the same number of electrons in each sub shell

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4
Q

Blocks in the periodic table

A

1-2= s
3-12=d
13-18= p

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5
Q

What is first ionisation energy?

A

The energy required to remove one electron from each atom in one molecule of gaseous atoms of an element to form one mole of gaseous 1+ ions

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6
Q

Factors affecting ionisation energy

A

First electron lost= highest energy level/least attraction from nucleus
Atomic radius: Higher=less nuclear attraction=decrease ionisation energy
Nuclear charge: more protons in nucleus= higher nuclear attraction= higher ionisation energy
Electron shielding: inner shell electrons repel outer shell electrons, more electrons in middle=higher shielding effect= lower ionisation energy

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7
Q

What affects the values of successive ionisation energies?

A

Gets much larger when changes from a further away shell to a closer one
When 1 e- is lost the positive nucleus attracts fewer electrons so are pulled in more tightly

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8
Q

Trends of first ionisation down a group

A

Decrease down a group:
Atomic radius increases
More inner shells= more shielding
Nuclear attraction decreases

Nuclear charge increase is outweighed by radius

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9
Q

Trends in first ionisation energy across a period

A

Increase across 1st 3 periods:
Nuclear charge increases
Same shell= similar shielding
Nuclear attraction increases
Atomic radius decreases
Each new period down= another shell= lower energy

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10
Q

Sub shell trends in ionisation energy

A

Falls when has to start filling a new sub shell
as had higher energy level
2p sub shell has higher energy level so less energy is required to remove than 2s so ionisation energy is less
Easier to remove electrons when some are paired and repel= lower ionisation energy

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11
Q

Metallic bonding definition

A

Strong electrostatic attraction between cations and delocalised electrons

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12
Q

Structure of a giant metallic lattice

A

Cations are in a fixed position
Delocalised electrons are mobile and can move throughout structure

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13
Q

Properties of metals

A

High conductivity= delocalised electrons carry charge and move through substance
Bp= High as need large amounts of energy to overcome electrostatic attraction between cations and delocalised electrons
Not soluble

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14
Q

Giant covalent structures

A

A network of loads of atoms held together by string covalent bonds

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15
Q

Examples of giant covalent lattices

A

Graphene= 3/4 e- used in bonding, can conduct, planar hexagonal layers, 120 degrees
Graphite= parallel graphite layers bonded w/ London forces
Diamond/silicon= 4/4 e- used in bonds, tetrahedral, 109.5 degrees

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16
Q

Giant covalent lattice properties

A

High bp= strong covalent bonds
Insoluble= bonds too strong
Non conductors apart from graphene/graphite because delocalised electron

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17
Q

Periodic trends in melting points and explanation

A

Across period 2&3:
Increase from Group 1-14 (giant)
Decrease 14-15 (sharp)
Comparatively low from 15-18 (simple)(weak IMFs, London)

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18
Q

Group 2 elements characteristics

A

Each has 2 electrons in outer shell
In outer s sub shell
Redox–>oxidised for a 2+ ion
Group 2 element is a reducing agent (reduces other species)

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19
Q

Trends in reactivity of Group 2 down the group

A

+ water= alkaline hydroxide & H2
Reactivity increases down the group and become stronger reducing agents
Metal+acid–> salt+H2
- React by losing electrons
- Ionisation energies decrease down group because of more shielding/radius

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20
Q

Metal oxide reaction w/ water

A

Form alkaline solution of metal hydroxides
Only slight;y soluble (group 2) so when solution is saturated any more metal and OH- form precipitate

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21
Q

Hydroxides solubility trends

A

Increases down the group
So the solution contains more OH- and alkaline
pH then increases= higher alkalinity

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22
Q

Group 2 uses

A

Calcium hydroxide as lime is added to fields to increase pH
Bases are used at antacids (calcium carbonate) to neutralise acids and Mg(OH)2

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23
Q

Trends in bp of halogens

A

Exist as diatomic molecules at RTP
Boiling point increases down the group
due to more electrons, more London forces, and more energy required to break the IMFs

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24
Q

Halogen definition

A

Outer shell s2p5 and gains 1 electron in redox to form 1- ion
Is an oxidising agent because oxidises another species

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25
Trends in reactivity of halogens
Decreases down the group Have to gain electrons Down group: Atomic radius increases More shells= more shielding Less nuclear attraction to attract negative electron Weaker oxidising agents
26
How to test for reactivity in halogens
Halogen-halide displacement reactions If halogen added is more reactive: Halogen is displaced from the solution by another The solution changes colour
27
Halogen colours in water and cyclohexane
Chlorine: Water= pale green Cyclohexane= pale green Bromine: Water= orange Cyclohexane= orange Iodine: Water= Brown Cyclohexane= violet Fluorine= pale yellow gas reacts w everything
28
What is disproportionation?
A redox reaction where the same element is both oxidised and reduced
29
Reaction of chlorine w/ water
Disproportion reaction takes place For each Cl2, one is oxidised and one is reduced HClO and HCl products
30
Reaction of chlorine w/ cold, dilute sodium hydroxide
More chlorine dissolves as sodium hydroxide is dissolved Disproportionation reaction NaClO, NaCl, and water products
31
Benefits of chlorine used in water treatments
Kills bacteria
32
Negatives of chlorine used in water treatments
Chlorinated hydrocarbons may form= cancer Chlorine gas
33
Halide ion precipitation reaction
Aq halid ions + Aq silver ions---> silver halid precipitate
34
Test for carbonate
1) Add nitric acid to sample 2) If bubbles produced, bubbled through lime water (Ca(OH)2) 3) If turns cloudy result is positive
35
Sulphate test
Aq barium ions are added to unknown compound solution If dense white precipitate forms= sulphate as barium sulphate= insoluble Use barium nitrate of doing halogen test after not barium chloride
36
Halide tests
1) Add aq silver nitrate to solution of a halide 2) Silver chloride= white precipitate, Silver bromide= cream, Silver iodide= yellow 3) Add Aq ammonia to test solubility as precipitate colours are hard to tell apart 4) Cl,= soluble in dilute Br= soluble in conc, I= insoluble
37
Sequence of tests in qualitative analysis
Carbonates-->sulphate--->halide Because neither sulphate or halide ions produce bubbles and if you do the sulphate test on a carbonate you will also get a white precipitate meaning that you can't tell what it is from so must only to sulphate test once you know there is no carbonate present
38
Cations test
1) Aq NaOH added to solution of ammonium ion 2) Ammonia gas made 3) Mixture warmed and ammonia gas released 4) Moist indicator paper to test. Ammonia alkaline so turns blue
39
Collision theory in rate of reaction
Must collide at correct orientation Must have sufficient energy to overcome activation energy barrier
40
Effects of changing condition on ROR
Increased conc of reactant= higher rate as particles will collide more frequently=more effective collisions Higher pressure of gas= higher ROR, will collide more frequently= more likely effective collisions
41
What is a catalyst?
Increases rate of reaction without being used up Allows reaction to proceed via a different route w/ lower activation energy
42
Types of catalysts
Homogeneous= same form as reactants (forms intermediate and regenerates) Heterogeneous= Different physical state from the reactants (adsorbs reactant molecules onto surface then leave via desorption)
43
Why are catalysts important industrially?
Reduces temperature and energy needs for processes Less electricity/fossil fuels used Cheaper and faster to make
44
Features of the Boltzmann distribution
No molecules have 0 energy Area under the curve= total number of molecules No max energy (does not meet x axis at high energy)
45
How does the Boltzmann distribution change w/ a higher temperature
Peak is lower and shifted to the right Greater proportion of molecules reach activation energy
46
How does a catalyst affect Boltzmann distribution
Greater proportion of molecules over activation energy (more left on graph)
47
La chatelier's principle
When a system in equilibrium is subjected to external change, the system readjust to minimise effect of change
48
Effect on concentration on equilibrium
Increase in concentration of reactants= more products are made so eq shifts to the right
49
Effect of increasing temperature on equilibrium
Shifts to favour the endothermic reaction because they absorb the external heat Endo=+ Exo=-
50
Effect of increase of pressure in equilibrium
Increase in pressure= shift to the side w/ fewer molecules of gaseous atoms
51
Effect of catalyst on equilibrium
Does not change position Speeds up forward and backwards equally Increases rate at which eq is reached
52
Kc equation and significance
Products/reactants mole number as power 1= eq halfway between reactants&products Bigger than 1= towards products Smaller than 1= towards reactants
53
Enthalpy definition
Heat energy in a chemical system
54
Enthalpy change calculation
H(products)-H(reactants)
55
Standard conditions
100kPa 298k 1moldm-3
56
Types of enthalpy changes
Reaction= Molar quantities in chem equation Formation= One mole of a compound is formed from it's elements Combustion= 1 mole of a substance reacts completely w/ O2 Neut= Reaction of an acid by a base-->1moleH2O
57
Energy change calculation
q=mcdeltaT m= mass of surroundings c= 4.18/specific heat capacity Delta T= temp chnage of surroundings
58
What affects measure energy change calculation?
Heat loss to other surroundings Incomplete combustion Evaporation Non-standard conditions = less exothermic
59
Average bond enthalpy
Energy required to break one mole of a specified type of bond in a gaseous molecule Always endothermic Positive enthalpy value
60
What happens to energy levels when bonds form?
Energy is released Exothermic Delta H= negative
61
Calculation of energy change from bond enthalpies
Sum of bond enthalpies in reactants- sum in products

Chemistry A-level (30 decks)

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