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Chem Flashcards

1
Q

Definitions of relative isotopic mass
Relative atomic mass
Relative molecular mass

A

Relative istopic mass: is the mass of one ISOTOPE compared to one twelfth of the mass of one carbon atoms of carbon-12
Relative atomic mass:is weighted mean mass of one ATOM compared to 1/12th of the mass of one atoms carbon 12
Relative molecular mass:is the AVERAGE mass of a molecule compared to one twelfth of the mass of one atom of carbon-12

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2
Q

Hydrated salt- practical

A

Weigh and empty clean dry crucible and lid.Add 2g of hydrated salt .Heat strongly with Bunsen burner.Let it cool and weigh again.reheat and re weigh.

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3
Q

Making A Solution

A
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4
Q

Ideal gas equation

A

PV=nRT

Pascal
M3
Mol
8.31 for R
Temperature +273 from C* to K

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5
Q

Avogradros constant

A

6.02 * 10^23
No of particles=moles * avogradore

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6
Q

Other formulas
Density
Percentage yield
Atom economy

A

Density= mass/ volume
Percentage yield= actual yield /theoretical yield *100
Atom economy=mass of useful products/mass of all reactant *100

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7
Q

Acids and Bases

A

Acids release H+ ions in aqueous solutions e.g. HCl,H2SO4,HNO3, CH3COOH
Bases neutral is acids.Alkalis are soluble bases and release OH- ions. NaOH,KOH,NH3 aq

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8
Q

REDOX

A

Oxidation is the process of electron loss
Oxidation number increases

Reduction is electron gain
Oxidation number decreases

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9
Q

Oxidation numbers

A

H 1
F -1
Cl,Br,I -1
O -2

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10
Q

Reactions of acid with metal

A

Acid + metal——> salt + H2

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11
Q

Sub energy levels

A

S holds 2 spherical
P holds 6 dumbbells
D holds 10
F holds 14

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12
Q

What are some exceptions to the rules of sub shells

A

4s fills up before 3d

The elements Cr and Cu are exceptions to the rule.
Each electron occupies one orbital before pairing to prevent repulsion between each other.

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13
Q

Ionic bonding

A

Ionic bonding is the electrostatic force of attraction between oppositely charged ions formed by electron transfer.Its stringer and have higher melting points when the ions are smaller/ have higher charges, e,g,MgO Mg+2 O-2.These form giant ionic lattices

High m.p
Non conductor of electricity when solid- only when molten or in solution
Usually soluble

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14
Q

Covalent Bonding

A

Covalent bond is the strong electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atoms
Dative covalent bond is when the shared pair comes from only one of the bonding atoms.
Structure:simple molecular (only used the words molecules and intermolecular forces when taking about simple molecular substances)

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15
Q

Comparison giant ionic and molecular

A

Giant ionic has high b.p and m.p. Due to strong forces
Molecular has low due to weak i.m. Forces

Giant ionic are usually soluble
Molecular usually poorly soluble

Conductivity is poor for molecular

General description:Gina ionic and crystalline solids
Molecular are mostly gases & liquids

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16
Q

Shapes of molecules

A

Linear : 2 bonding 180*
Trigunal planar:3 bonding 120*
Tetrahedral: 4 bonding 109.5
Trigunal pyramidal 3 bonding 1 lone 107*
Non-linear 2 bonding 2 lone 104.5
octahedral 6 bonding 90*

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17
Q

Electronegativity

A

Electronegativity is the relative tendency of an atom in a covalent bond in a molecule to attract electrons in a covalent bond to itself.

Factors:increases across as proton no. Increases and atomic radius increases as the electrons are pulled in more
It decreases down a group because distance between the nucleus and outer electron increases and the shielding increases

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18
Q

Dipole

A

Polar covalent bonds have unequal distribution of electron and produces a charge separation (dipole) ends. S+ and S-

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19
Q

Intermolecular bonding

A

Induced dipole-dipole interactions:All molecular substances (not in ionic).In any molecule the electrons are moving randomly and constantly this causes electron density to fluctuate adn parts of the molecule becomes more or less negative .These temporary dipoles induces dipoles in neighbouring molecules,called induced dipoles.The more electrons,the stronger.
Permanent dipole-dipole:polar molecules,stronger,higher b.p.,they occur in addition to London forces.

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20
Q

Hydrogen bonding

A

Occurs between hydrogen attached to nitrogen,oxygen or fluorine, with a lone pair.
Ice:when in solid state molecules are held further apart which explains lower density.
Iodine:crystal contains regular arrangement of weak London forces. I2

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21
Q

Periodicity

A

Is a repeating pattern across different period

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22
Q

First Ionisation Energy

A

Is energy needed to remove an electron from each atom in one mole of gaseous atoms
H(g)——> H+(g) + e-

Factors affecting:
Attraction of the nucleus,the more protons the more attraction
The distance of electrons from nucleus:the bigger atom,the further ,the weaker the attraction
Shielding :outer electrons are repelled by inner electrons

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23
Q

Pattern of period 2 elements

A

He has highest f.i.e, shell closest to nucleus,no shielding,has one more proton than H.
F.i.e decrease down a group bcs outer electron around in shells further from nucleus and more shielding
F.i.e increases across a period same shielding effect,electron are added proton increases decreasing atomic radii.

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24
Q

Anomalies in period 2&3 f.i.e

A

Na has lower f.i.e than Neon because Na is more shielded ,electron easier to remove.
Small drop between Mg and Al because Al is starting to fill up 3p whereas Mg has full 3s?Ekectrons in 3p are slightly easier to remove because 3p electron are higher in energy and shielded .
Small drop from P to S.Sulphur has 4 electron in 3p and 4th electron is starting to pair.When the second electron is added to a 3p orbital there is a slight repulsion between which makes the second electron easier to remove.

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25
Q

Metallic bonding

A

Is the electrostatic force of attraction between the positive metal ions and the delocalised electrons
Giant metallic lattice ,Li,Be ,Na,Mg,Al
Factors:
Number of protons,the more protons in the stronger
Number of electron per atom,the more the stronger
The smaller the ion the stronger

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26
Q

Macromolecules

A

Diamond:tetrahedral 4 atoms in covalent bonds
Graphite:planar of C atoms 3 covalent bonds,4th electron delocalised.
B,C ,Si
They have very high melting points

27
Q

Comparison macromolecules & metallic

A

High b.p and m.p
Both insoluble in water
Conductivity: poor for diamond,good for graphite.
Good for metallic
Conductivity when molten:poor for macromolecules
General description:solids for macromolecules
Shiny metal,malleable slide over easily for metallic

28
Q

Group 2

A

Meltin point decreases as you go down,atomic size increases,ion further away from electron.
Reactivity increases down the group.Atomic radii increased,more shielding,nuclear attraction decreases easier to remove electrons.
MgO is a white solid.MgO will react slowly without a flame.

29
Q

Reactions of Group 2

A

Water:mg reacts with steam to MgO or with warm water to Mg(OH)2 .Others reacts with cold water to produce hydroxides+H2
You would observe fizzing,metal dissolving ,heating up,calcium forms white precipitate
Acid:forms salt and hydrogen

30
Q

Action of water on metal 2 oxides

A

CaO+H2O——->Ca(OH)2 (aq7 pH 12 bcs it’s quite soluble
It’s also used in agriculture to neutralise soils.If excess soil too alkaline.Aquoues CaOH can be used to test for CO2
MgO+H2O——>Mg(OH)2 (s) pH9 bcs mgOH is slightly soluble.Used in medicine to neutralised excess acid in stomach and treat constipation.Helps with indigestion.Safe to use as its weakly alkaline

31
Q

Halogens-diatomic

A

F2-pale yellow gas,very reactive
Cl2-greenish reactive gas,poisonous
Br2-red liquid,gives brown poisonous fumes
I2-shiny grey solids sublimes to purple gas
Higher m.p and b.p as you go down,larger molecules,more electrons,stronger London forces.
A halogen more reactive will displace the less reactive.
Reactivity decreases down as atoms get bigger more shielding so they less attract and accept electrons .

32
Q

Learn the table for chlorine ,bromine and iodine

A
33
Q

Disproportionation

A

Is the name for a reaction where an element simultaneously oxidises and reduces.
Cl2(g)+H2O——>HClO(aq) +HCl (aq)
Chlorine is used in water treatment to kill bacteria.
Cl2 and cold dilute NaOH .Colour of halogen fades to colourless.NaCl and NaClO is used as bleach and disinfect bacteria

34
Q

Halide ions and AgNo3 silver nitrate

A

Add nitric acid (to react with any carbonates present )first and then AgNO3 drop wise .
F does not produce precipitate
Cl white. Br cream. I pale yellow precipitate

Silver chloride dissolves in dilute ammonia
Silver bromide dissolves in con. Ammonia
Silver iodide does not react

35
Q

Qualitative Analysis

A

Presence of CO3 carbonate- add dilute acid .Bubble gas through lime water,will turn cloudy.
Presence of SO4 sulfate-acidified (reacts with CO3 impurities that would give false results)BaCl2 solution used as reagent .If +ve white precipitate form.
Sequence is carbonate,sulfate then halife (false positives due to BaCO3 or Ag2SO4)
Ammonium ion NH4+: react with warm NaOH (aq) forming NH3 gas,this will turn red litmus paper blue

36
Q

Exothermic & endothermic

A

Exothermic:Energy is transferred from system (chemicals) to surrounding.Products have less energy. change in H is negative
Endothermic:energy transferred from surrounding to chemical.Require an input of energy.H is positive

37
Q

Activation energy

A

Minimum energy which particles need to collide to start a reaction

38
Q

Standard conditions

A

100 kP
298 K
Solutions at 1 mol dm3
Normal state at room temp.

39
Q

Enthalpy change of reaction

A

Is the Enthalpy change that accompanies a reaction in the molar quantities shown in a chemical reaction under standard conditions,with all reactants and products in their standard states.
Triangle Hr central line

40
Q

Enthalpy change of formation

A

Enthalpy change that takes place when one mole of a compound is formed from its elements under standard conditions,with all reactants and products in their standard states

41
Q

Enthalpy change of combustion

A

The Enthalpy change that takes place when one mole of a substance is reacted completely with oxygen under standard conditions with all reactants and products in their standards states

42
Q

Enthalpy change of neutralisation

A

Triangle neut H central line
The Enthalpy change that accompanies the reaction of an acid by a base to form one mole of H2P (l) under standards conditions with all reactants and products in their standard states.

43
Q

Calometry formula

A

Energy change Q (J) = mass of solution m (g) * heat capacity c (J per g per K) * temperature change T (K)

44
Q

Calometric method

A
45
Q

Errors in calometric method

A

Energy losses from calorimeter
Incomplete combustion of fuel
Incomplete transfer of energy
Evaporation of fuel after weighing
Heat capacity of calorimeter not included
Measurements not carried out under standard conditions H20 is gas not liquid in here

46
Q

Average bond Enthalpy

A

The mean bond Enthalpy us the Enthalpy change when one mole of bonds of us broken.
Bond breaking absorber energy
Bond making releases energy

47
Q

Collision theory

A

Reactions can only occur when collisions take place between particles having sufficient energy.
At higher concentration ,there are more particles so frequency of collisions increases.

48
Q

Reaction rate definition and how to find it

A

Change in concentration of a substance in unit time. Mole dm3 per seconds.Sodium thiosukftae and HCl see how long it takes for cross to disappear .1/time plot graph.

49
Q

Catalysts

A

Increases reaction rates without getting used up.They provide an alternative route or mechanism with a lower a.e.
A heterogeneous is in a different phase from reactants.usually solids
A homogeneous catalyst is in the same phase as reactants.reactions proceeds through intermediate species

50
Q

Benefits of Catalysts

A

Lower temp. And pressures ,save energy costs ,fewer CO2 emissions.
Better atom economy with reduced waste,fewer undesired products .
They are often enzymes

51
Q

Investigate rates of reaction

A

Change in volume of a gas using gas syringe
Change in mass if gas is involved
Colorimetry if there is a change in colour

52
Q

Dynamic equilibrium

A

This occurs when forward and backward reactions are occurring at equal rates in a closed system and the concentration of reactants and products do not change.

53
Q

Effect of increasing temperature in equilibrium

A

If temperature increased (decreased) the equilibrium will shift to oppose this and move in the endothermic (exotgermic) direction to reduce (increases)temperature by absorbing heat.
The position of the equilibrium will shift to the …
Giving a lower yield of …..

54
Q

Effect of pressure in equilibrium

A

Increasing (decreasing) pressure will cause the equilibrium to shift towards the side either fewer (more)moles of gas to oppose the change and thereby reduce (increase) the pressure .
Position will shift to …
Giving higher/less yield of…

55
Q

Effect of concentration on equilibrium

A

I2 + 2OH- <———->I- + IO- + H2O (colourless)
(Brown)

Increasing concentration of OH- ions causes the equilibrium to shift to oppose this and move in the forward direction to remove OH- ions.The position of equilibrium will shift toward the right giving higher yields of I- and OH-.
Adding H+ ions reacts with the OH- ions and reduces their concentration so the equilibrium shifts back to the left giving brown colour.

56
Q

Uses of equilibrium : Haber process

A

N2 +3H2 ————> 2NH3

T=450 P=200-1000 atm catalyst=iron

Low temp. Gives good yield but slow rate.high pressure give good yield and good rate.

57
Q

Uses of equilibrium: contact process

A

1- S (s) + O2 (g) ————>SO2 (g)
2- SO2 (g) +1/2 O2 (g) <————>SO3
T=450. P=1 to 2 atm. Catalyst =V2O5

Low temp. Compromise moderate temperature used.

58
Q

Equilibrium: Productions of methanol from CO

A

CO (g) +2H2 (g) <————>CH3OH (g)
T=400 P=50 atm catalyst=chromium and zinc oxide

Compromise temp. Used

59
Q

Hydration of ethane to produce ethanol

A

CH2=CH2 +H2. <—————> CH3CH2OH
T=300 P=70. Catalyst=concentrated H3PO4
Too high pressure would lead to unwanted polymerisation of ethene to poly (ethene)

60
Q

Equilibrium constant Kc

A

Kc= C^c * D^d / A^a * B^b
Use concentration
No units
Kc only changes with temperature.

61
Q

The greenhouse effect-IR

A

CO2 ,methane,water vapour are greenhouse gases that trap earths radiated infra red energy.infrared radiation is absorbed by c=o,c-h and o-h bonds which contributes to global warming.
IR can be used to monitor gases causing air pollution.Modern breathalysers measure ethanol using infrared spectroscopy l

62
Q

CFC

A

Used as aerosols,refrigerants and air conditioning.Has been replaced with HFCs

63
Q

Ozone layer O3

A

Filters suns harmful UV radiation.
Ozone formation O2 + UV light > O+O
O+O2>O3
Ozone depletion O3 +Uv>O2 +O

64
Q

Destruction of Ozone layer

A

Chlorine free radical catalyse decomposition
Cl+O3>ClO +O2
ClO+O>O2 +Cl*
Basically O3+O*>2O2

Nitrogen oxide
NO+O3>NO2+O2
NO2+O*>O2+NO

overall equation
O3+O*>2O2

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