Periodic Table, Elements And Physical Chem Flashcards
1
Q
Define relative atomic mass
A
Mean mass of an atom of an element, compared to 1/12th of the mass of an atom of carbon-12
2
Q
Define relative molecular mass
A
Mean mass of a molecule compared to 1/12th of the mass of an atom of carbon-12
3
Q
Define relative isotopic mass
A
Mean mass of an atom of an isotope compared to 1/12th of the mass of an atom of carbon-12
4
Q
Define theoretical yield
A
Amount of product produced assuming no product was lost and all reactants react fully
5
Q
Why is percentage yield never 100%?
A
- you may lose product when transferring from beaker to beaker
- not all the reactants reacted
- may have lost some product as a gas
- may be impurities
6
Q
What is the importance of high atom economy?
A
- Produce less waste and is better for the environment
- Produce less biproducts so less time and money wasted to separating
- Raw materials are used more efficiently so more sustainable
- Compainies will try to use reactions near 100%
7
Q
What are monoprotic acids?
A
1 mole of an acid will produce 1 mole of H+ ions
8
Q
What are diprotic acids?
A
1 mole of an acid will produce 2 mole of H+ (H2SO4)
9
Q
What is a triprotic acid?
A
An acid that will produce 3 mole of H+ (H3PO4)
10
Q
How does ammonia react with acids?
A
Produce an ammonium salt but no water
11
Q
Give the solubility rules
A
Soluble
- all group 1 compounds
- all nitrate compounds
- all ammonium compounds
insoluble
- Ag+, Pb2+ chlorides
- Ba2+, Pb2+ sulfates
- most hydroxide
- most carbonates
12
Q
Describe how to make standard solutions
A
- Weigh solid precisely w/ balance and weighing boat
- Transfer tp a beaker and wash any solid left into beaker using distilled water
- Dissolve w/ distilled and stir to ensure all dissolves
- Transfer to volumetric flask w/ funnel and rinse beaker w/ distilled
- Use distilled to fill to graduation line and use pipette to fill to line when near
- Invert flask a few times= thoroughly mixed
13
Q
Define a reducing agent
A
- lose electrons and are oxidised themselves
14
Q
Define oxidising agent
A
- gain electrons and are reduced themselves
15
Q
What is disproportion?
A
When the same element is simultaneously oxidised and reduced
16
Q
What are the subshells and how many orbitals do they have?
A
S= 1 orbital
P= 3 orbital
D= 5 orbitals
F= 7 orbitals
17
Q
What is shell no. Also known as?
A
Principle Quantum Number
18
Q
Describe the shape of the s orbital
A
Spherical and 2 electrons can move anywhere within it
19
Q
Describe the shape of the piece orbitals
A
3 p orbitals in the shape of dumbells and cak hold up to 2 electrons within this shape. The orbitals are 90⁰ to eachother
20
Q
What is spin pairing?
A
When 2 electrons occupy 1 orbital, they spin in opposite directions
21
Q
Describe how electrons fill orbitals
A
- fill from the lowest energy level upwards
- fill orbitals singly first then pair up due to electron repulsion
22
Q
How are electrons removed from orbitals when forming ions?
A
- removed from highest energy level first
23
Q
How do permanent dipole-dipoles occur?
A
- exist in molecules with polarity
- stronger than London forces
24
Q
Why does water have high SHC and why is ice less dense than water?
A
- strong H bonds mean that lots of energy is needed to change the temp of water
- ice forms a regular structure held by H bonding keeping water molecules far apart
25
Why does BP of hydrogen halides from HCl- HI increase down the group?
- due to increased mass of molecules and bigger electron cloud so more London forces
26
Across the periodic table, what are the block names?
s-block, d-block, p-block
27
Define 1st ionisation energy
Minimum amount of energy needed to lose 1 mole of electrons from 1 mole of gaseous atoms
28
What are the 3 factors affecting ionisation energy?
1. Atomic charge
2. Shielding
3. Atomic size
29
Describe the change in ionisation energy across a period
- increasing number of protons means there is increased nuclear attraction
- shielding is similar and distance from nucleus marginally decreases
** ionisation energy increases
30
Explain why elements like Phosphorous have electrons removed more easily
- involves taking an electron from an orbital with 2 electrons in it
- electrons repel and so less energy is needed to remove the electron
31
What is successive ionisation energy?
- the removal of more than 1 electron from the same atom
- General increase as you are removing an electron from an increasingly more positive ion
32
Describe the structure and properties of graphite (6)
- each carbon bonded 3 times with 4th electron delocalised
-giant covalent, lots of covalent bonds so high MP
- delocalised electrons= conducts electricity
- insoluble as covalent bonds to strong to break
- layers slide easily as weak forces between them
- low density as layers are far apart
33
Describe the structure and properties of diamond (5)
- each carbon bonded 4 times in tetrahedral
- tightly packed sk good heat conductor
-high MP due to many covalent bonds and very hard
- doesn't conduct electricity as no delocalised electrons
- insoluble as covalent bonds too strong to break
34
Describe the structure and properties of silicon
- same structure as diamond so same properties
35
Describe the structure and properties of graphene and its uses
- 1 layer of graphite and is 1 atom thick made up of hexagonal rings
- delocalised electrons= conducts electricity
- delocalised electrons strengthen covalent bonds= high strength
- lightweight and transparent as 1 atom thick so used in phone screens, aircraft shells
36
Why are metals good thermal conductors?
Giant metallic structures are good thermal conductirs as delocalised electrons can transfer kinetic energy
37
What do group 2 oxides look like?
White solids
38
Why do group 2 oxides become more alkaline down the group?
- the hydroxide formed when they react with water are more Soluble so more OH- ions are formed
39
Define electronegativity
- ability for an atom to attract electrons toward itself in a covalent bond
40
How does electronegativity change down the halide group?
- decreases as atoms get larger and the distance between positive nucleus and bonding electrons increases + more shielding
- less strongly attracted
41
How is bleach made?
2NaOH + Cl2 -> NaClO +NaCl + H2O
a disproportionation reaction
42
What is the equation when chlorine reacts with water and why is this useful?
H2O + Cl2 -> HCl + HClO
- chloric (I)acid ionises in water to make chlorate (I) ions and H3O+
- chlorate (I) ions (ClO-) kill bacteria which is useful for drinking water and pools
43
What are the advanatges and disadvantages of chlorinating water?
:)
- destroys microorganisms that cause disease like cholera epidemic
- long lasting so reduces build up further down the supply
- reduce growth of algae that discolours the water (give bad smell + taste)
:(
- chlorine gas is toxic + irritates respiritory system
- liquid chlorine gives sever chemical burns
- chlorine can react w/ organic compounds present in water to make chloroalkanes (link to cancer)
44
Why may some people object to chlorination of water?
- occurs across the UK and we have no choice
- some say it is forced medication of a whole population
45
Give the 2 alternatives to chlorination of water
1. ozone= powerful oxidising agent that kills microorganisms
- short half lide so not permanent and would be expensive
2. UV light= damages DNA in microorganisms
- ineffective in cloudy water + won't prevent contamination further down the process
46
How do you test for ammonium compounds?
- add sodium hydroxide + gently heat = ammonia gas produced
- damp red litmus will turn blue
47
Define standard enthalpy change of formation
- enthalpy change when 1 mole of a compound is formed from its elemnts in their standard states, under standard conditions
48
Define standard enthalpy change of neutralisation
- enthalpy change when an acid and alkali react to form 1 mole of water, under standard conditions
49
Define enthalpy change of combustion
- enthlapy change when 1 mole of a substance is completely burned in O2 to make CO2 and H2O under standard conditions
50
What are the standard conditions?
- 100kPa
- 298K/ 25ºC
51
How to detrmine if a reaction is endo/exothermic from bonds made/broken
- if more energy released when bonds form than break = exothermic
- if more energy released when bonds break than form = endothermic
- bonds made - bonds broken
52
Define average bond enthalpy
- the energy required to break 1 mole of bonds between two atoms in gaseous state
53
Define actual bond enthalpy
- actual= energy reuqired to break a specific bond in a particular molecule and can be affected by chemical environment and the presence of neighbouring atoms, functional groups
54
What is the difference between actual bond enthalpy and average bond enthalpy?
- actual bond enthalpies are specific to individual molecules and bonds, average are general values representing average strength of a specific type of bond
- actual are typically expreimentally determined for specific molecules, while average are tabulated values based on range of compounds
55
What are the benefits of using catalysts?
- lower temperautrues and pressures required = less money spent and CO2
-less waste is produced allowing scientists to use reeactions w/ better atom economies
56
What are the ways to measure rate of reaction?
1. how long it takes for a precipitate to form (disappearing cross) / however difficult to determine when disappears
2. volume of gas produced (gas syringe)
3. amount of mass lost if gases are produced (use fume cupboard if toxic gas)
57
Define dynamic equilibrium
- exists in a closed system when the rate of the forward reaction is equal to the rate of the reverse reaction
- the conc of reactants and products do not change
58
What kind of equilibrium does Le Chatelier's principle work for?
- homogenous equilibria where reactants and products are in the same state
59
What happens when you decrease/ increase pressure in an equilibrium?
- increase means equ will shift to side with least gaseous moles to reduce pressure
- decrease mean equ will shift to side with most gaseous moles to increase the pressure
60
What is the compromise between when using pressures and temperatures for industrial reactions?
- temperature compromises between yield and rate
- pressure compomises between yield/speed and cost + safety
61
How does temperature affect Kc?
- changing temperature will change equilibrium concentrations
- if temp change cause equ to shift to right = _Kc increases_
- if temp change causes equ to shift to left = _Kc decreases_
62
How can we determine where equilibrium lies from Kc value?
- if Kc is small (less than 1) = equ lies to the L as mixture contains mostly reactants
- if Kc large (more than 1) = equ lies to the R as mixture contains mostly products
- if Kc is close to 1 (0.1-10), mixture contains similar amounts of product and reactants
63
How does water dissociate?
- into hydroxide ions and hydroxonium ions
- weakly dissociates as very few ions so we assume concentration of H2O is constant
- 2H2O ---> H3O+ +OH-
- simplified to H2 O -> H+ + OH-
64
Describe Kw
- the equilibrium constant for the self-ionisation of water
- value of 1.00x10-14 mol2dm-6
- Kw= [H+][OH-]
- pure water has an equal conc. of H+ to OH- so Kw=[H+]2
65
How do we work out pH from mono/diprotic acids?
_monoprotic acids_ dissociate and produce one H+ ion
- therefore [H+]=[HA]
_diprotic acids_ dissociate and produce 2 H+ ions for every acid molecule
- therefore, 2[H+]=[HA]
66
How do we figure out pH of strong bases?
- assume they dissociate fully
- the [OH-] value will be the same as the conc.
67
Describe how you would use a pH probe
- must calibrate to give reliable readings
- first place in distilled water and the meter should give pH 7, recalibrate if not
- repeat w/ standard solutions at pH 4 and pH10 (rinse between each pH w/ distilled)
68
What is the equivalence point/ end point?
- the acid has been neutralised fully by the base
- we assume [H+]=[OH-]
- on the verticle section of the graph in the middle
69
How do you choose a suitable indicator?
- must change colour entirely w/in the vertical part of the titration
70
What titrations can methyl orange be used for?
- strong acid/ strong base
- strong acid/ weak base
71
What titrations can phenolphthalein be used for?
- weak acids/ strong base
72
What are titrations are indicators not used for?
- weak base/ weak acid titrations as no sharp pH change
73
Define buffers
- a solution as a system that minimises pH changes on addition of small amounts of an acid/base
- made up of weak acid and salt (excess weak acid + base reacted)
74
What happens when we add an acid to a buffer?
buffer equ: HA ⇌ A- + H+
- H+ ions in the acid react w/ A- ions in the solutions
- equilibrium shifts to L as more HA produced
75
What happens when we add a base to a buffer?
- the OH- ions react w/ H+ in the solution
- causing a low conc. of H+ which is counteracted by Le Chatelier's principle and equilibrium shifts to R
- replaces H+ ions
76
Describe the use of buffers in maintaining blood pH
- keep blood at pH 7.4
- carbonic acid- hydrogen carbonate system :
- H2CO3 ⇌ HCO2- + H+
AND
- H2CO3⇌H2O + CO2
- carbonic acid is controlled by respiration in cells
- when we breathe out CO2 the level of carbonic acid reduces and equ shifts to R to replace
- hydrogen carbonate is controlled by kidneys
77
Describe the iodine clock experiment
1. H2O2 +2I- + 2H+ -> 2H2O + I2
2. 2S2O3 2- + I2 -> 2 I- + S4O62-
- add sodium thiosulfate and starch to excess hydrogen peroxide
- sodium thiosulfate reacts immediately w/ iodine produced in reaction 1
- when no more sodium thiosulfate the iodine reacts with starch = deep blue/black
** vary conc. of iodine and/or hydrigen peroxide and keep everything else constant. This allows us to see what order the reaction is
78
Define orders of reactions
- an order is the power to which a concentration is raised to in the rate equation
- tells us how conc affects rate
79
Define 0 order, 1st order and 2nd order
0= changes in concentration has no affect on rate
1st= changes in concentration is proportional to the change of rate
2nd= changes in concentration have a squared proportional change on rate
80
What do the rate-concentration graphs for 0 order, 1st order and 2nd look like?
zero= horizontal line as changing conc. doesn't change the rate
1st= straight diagonal line as changes to conc. changes rate equally
2nd= curved line as changing conc. squares the rate
81
What does a conc-time graph look like for first and zero order?
zero= straight negative line
first= curved negative line
82
How do you work out k from half life of a 1st order conc-time graph?
- k= ln(2/time of half life)
- units are s-1
83
What is the rate determining step?
- slowest step in a multi-step reaction
the whole reaction rate depends on how quick the rate determing step is
84
What happens to the rate constant when Ea is decreased?
- rate constant gets bigger as rate of reaction increases due to more particles having enough energy to collide and react
85
What happens to the rate constant when temperature increases?
- k increases because particles have more kinetic energy and more likely to collide with at least the Ea
- therefore rate of reaction increases
86
When using the Arrhenius equation, how can we rearrange it to form y=mx+c?
lnk= -Ea/R x 1/T +lnA
- lnk=y
- -Ea/R=gradient
- 1/T= x
- lnA=c
87
How does temperature affect Kp?
**eg if exo forward and endo backward**
- increase temperature= equilibrium shifts to endo thermic to counteract increase in temp. and Kp decreases (denominator increases)
- decrease in temperature= equ shifts to exothermic to counteract decrease in temp. and Kp increases
88
How does pressure affect Kp?
- it doesn't because the partial pressure ratio of reactant to product stays the same
89
Define lattice enthalpy
- formation of 1 mole of an ionic lattice from gaseous ions
- measure of the strength of ionic bonding in a giant ionic lattice
90
Define enthalpy change of 1st ionisation
- enthaply change when 1 mol of gaseous 1+ ions are made from 1 mol of gaseous atoms
91
What affects the strength of ionic bonding?
1. size of the charge = bigger charge means stronger electrostatic attraction between ions/ more energy needed to overcome these forces
2. size of the ion (ionic radii)= smaller the ion, the stronger the electrostatic attraction because smaller ions can pack together more closely/ more energy needed to overcome the forces
92
What is charge density?
- high charge density means ions have a high charge and small ionic radii
93
Define enthalpy change of solution
- enthalpy change when 1 mol of an ionic substance is dissolved in the minimum amount of solvent
- can be worked out by lattice dissociation enthalpy (when you break the lattice into its gasous ions) and enthalpy of hydration
94
What must happen for a substance to dissolve?
1. substance bonds must break (endothermic)
2. new bonds form between the solvent and the substance (exothermic)
95
Define enthalpy change of hydration
- when 1 mol of aqueous ion is made from 1 mole of gasous ions
96
What affects enthalpy change of hydration?
1. higher charge attracts water molecules more strongly due to stronger electrostatic attraction= _more exothermic enthalpy of hydration_
2. smaller ions have a higher charge density so can attract water molecules more strongly= _more xothermic enthalpy change of hydration_
97
Define entropy
- the measure of disorder in a system (more disorder=higher entropy)
- way in which energy can be shared out between particles
- JK-1mol-1
98
How will a reaction tend towards?
- a reaction will tend towards more disorder and hence an increase in entropy
99
What are the conditions for standard entropy?
- 1 mol of substance
- 100kPa
- 298K
100
Describe Gibbs Free Energy
- ΔG= ΔH-TΔS
- all must be in Jmol-1, K and JK-1mol-1
- ΔG is Jmol-1
101
Describe feasibility of a reaction
- a reaction is feasible in theory if ΔG is _-ve or 0_
- even if a reaction is feasible, may not observe a reaction as Ea may be too high/ rate of reaction is too slow
102
How can we work out the lowest temperature at which a reaction becomes feasible?
- when ΔG=0 a reaction is just feasible
- therefore T=ΔH/ΔS
103
Where do electrons flow from in electrochemical cells?
- flow from more reactive metal to a less reactive one
104
What are electrode potentials?
- each half cell has an electrode potential measured in volts
- tells us how easily the half cell gives up electrons
- most negative half cell undergoes oxidation
-
105
What is the standard Hydrogen electrode?
- used as a reference to measure standard electrode potentials
- the SHE has E which is equal to 0V
- H2 in at 298K, 100KPa and solution of 1moldm-3 H+ ions
- connects to a half cell to find the standard electrode potential of the element
- the solution of 1moldm-3 of H+ is made from 1moldm-3 HCl or 0.5moldm-3 of H2SO4
106
How do we simplify how we set up a cell
- most negative half cell potential on the LEFT side of double line(salt bridge)
- reduced form|oxidised form||oxidised form| reduced form
- Zn(s)|Zn2+(aq)||Cu2+(aq)|Cu(s)
107
How do we predict feasibility of a reaction using electrode potentials?
1. identify which is oxidised and reverse it
2. combine the 2 equations to get the feasible reaction and compare the equation to the one stated in the question
3. the electrode cel potential will always be positive if it is feasible
108
What is a transition element? What elements behave differently when electrons fill orbitals?
- an element that can form at least one stable ion w/ a partially filled d-subshell
- chromium and copper
- chromium has one of its 4s electrons move into the 3d orbital to create a more stable half full 3d subshell
- copper has one of its 4s electrons move into the 3d orbital to create a more stable full 3d sub-shell
109
What elements are in the d block but not transition elements and why?
- scandium isn't because its stable Sc3+ ion has an empty d-subshell so it doesn't have a partially filled d-subshell
- zinc isn't because its stable Zn2+ ion has a full d subshell so is not partially filled
110
How are electrons removed from transition metals?
- lost from 4s before 3d
111
Why are transition metals good catalysts?
- they have variable oxidation states so they can receive and lose electrons in the d-orbitals to speed up reactions
- have surfaces that allow substances to adsorb to the surface= lower Ea
- less energy needed for high temp and pressure
112
What are the risks of using transition elements as catalysts?
1. Copper
- long term exposure to copper can cause damage to liver
- a ring of copper in the eye shows someone is suffering from copper poisoning
2. Manganese
- long term exposure can cause psychiatric issues and physical tremors
113
What is a complex ion?
- where a central transition metal ion is surrounded by ligands bonded by dative covalent (coordinate bonds)
-
114
What are ligands?
- ligands have at least 1 lone pair of electrons where they are used to form coordinate bonds w/ the metal
- can be monodentate (only have 1 lone pair), bidentate (2 lone pairs) or multidentate (can form more than 1 coordinate bond)
115
What is the coordination number and what shapes do they form?
- the number of coordinate bonds in a complex NOT the number of ligands
1. coordination no. 4:
- tetrahedral (109.5) [CuCl4]2-
- square planar (90) Pt[(NH3)2 Cl2]
2. coordination no. 6:
- octrahedral
116
Describe haemoglobin
- haemoglobin is a proteinused to transport O2 around the body in blood
- octahedral structure with 4 nitrogens in the middle section (haem)
- one is from a large protein called globin
- the final coordinate bond is formed with O2 or H2O
117
Describe how haemoglobin works and carries O2
- oxygen substitutes the water ligand in the lungs where O2 conc is high = oxyhaemoglobin which is transported around the body
- oxyhaemobglobin gives up O2 to where it is needed and water takes the place and haemoglobin goes back to lungs to start again
118
Why does CO affect haemoglobin's ability to carry O2?
- if CO is inhaled, the water ligand is substituted w/ CO ligand
- the CO bonds strongly so is not readily replaced w/ O2 or water so O2 can't be transported
- causes O2 starvation so CO is poisonous
119
Describe how cis-platin acts as an anti cancer drug
- the 2 chloride ligands are very easy to displace so they displace and bond to 2 N atoms on DNA molecules inside the cancerous cell
- blocks and prevents the cancerous cell from reproducing by division and the cell will die as it can't repair its damage
:( also prevents normal cells from reproducing, including blood cells = can supress the immune system and increase risk of infection
120
How would you work out the overall order of the reaction?
- use the equation k=rate/Reactant order to try out if the units match
121
What is meant byu the term weighted mean mass?
- the mean mass taking the relative abundancies of isotopes into account
122
What are the signs of ΔH and ΔS to make a reaction feasible?
- ΔH should be negative and ΔS positive so that ΔG is below 0 at all temperatures
123
Explain why rubidium chloride is reasonably soluble at room temperature
- dissolving the solid compound increases ΔS so that it can overcome the effects of ΔH and therefore, ΔG is negative and the reaction is feasible
- the activation energy must also be low enough for the reaction to happen spontaneously
124
What is the purpose of the salt bridge?
- completes the circuit by allowing ions to move between the half cells
- prevents the two solutions mixing together
125
Describe a fuel cell
- uses the enrgy from the reaction of a fuel with oxygen to create a voltage
- oxygen at positive electrode and fuel at negative where oxidation takes p[lace
