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Chemistry > Periodic Table, Elements And Physical Chem > Flashcards

Periodic Table, Elements And Physical Chem Flashcards

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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

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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

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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

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4
Q

Define theoretical yield

A

Amount of product produced assuming no product was lost and all reactants react fully

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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
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6
Q

What is the importance of high atom economy?

A
  1. Produce less waste and is better for the environment
  2. Produce less biproducts so less time and money wasted to separating
  3. Raw materials are used more efficiently so more sustainable
  4. Compainies will try to use reactions near 100%
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7
Q

What are monoprotic acids?

A

1 mole of an acid will produce 1 mole of H+ ions

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8
Q

What are diprotic acids?

A

1 mole of an acid will produce 2 mole of H+ (H2SO4)

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9
Q

What is a triprotic acid?

A

An acid that will produce 3 mole of H+ (H3PO4)

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10
Q

How does ammonia react with acids?

A

Produce an ammonium salt but no water

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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

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12
Q

Describe how to make standard solutions

A
  1. Weigh solid precisely w/ balance and weighing boat
  2. Transfer tp a beaker and wash any solid left into beaker using distilled water
  3. Dissolve w/ distilled and stir to ensure all dissolves
  4. Transfer to volumetric flask w/ funnel and rinse beaker w/ distilled
  5. Use distilled to fill to graduation line and use pipette to fill to line when near
  6. Invert flask a few times= thoroughly mixed
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13
Q

Define a reducing agent

A
  • lose electrons and are oxidised themselves
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14
Q

Define oxidising agent

A
  • gain electrons and are reduced themselves
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15
Q

What is disproportion?

A

When the same element is simultaneously oxidised and reduced

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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

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17
Q

What is shell no. Also known as?

A

Principle Quantum Number

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18
Q

Describe the shape of the s orbital

A

Spherical and 2 electrons can move anywhere within it

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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

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20
Q

What is spin pairing?

A

When 2 electrons occupy 1 orbital, they spin in opposite directions

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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
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22
Q

How are electrons removed from orbitals when forming ions?

A
  • removed from highest energy level first
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23
Q

How do permanent dipole-dipoles occur?

A
  • exist in molecules with polarity
  • stronger than London forces
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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
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25
Q

Why does BP of hydrogen halides from HCl- HI increase down the group?

A
  • due to increased mass of molecules and bigger electron cloud so more London forces
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26
Q

Across the periodic table, what are the block names?

A

s-block, d-block, p-block

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27
Q

Define 1st ionisation energy

A

Minimum amount of energy needed to lose 1 mole of electrons from 1 mole of gaseous atoms

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28
Q

What are the 3 factors affecting ionisation energy?

A
  1. Atomic charge
  2. Shielding
  3. Atomic size
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29
Q

Describe the change in ionisation energy across a period

A
  • increasing number of protons means there is increased nuclear attraction
  • shielding is similar and distance from nucleus marginally decreases
    ** ionisation energy increases
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30
Q

Explain why elements like Phosphorous have electrons removed more easily

A
  • involves taking an electron from an orbital with 2 electrons in it
  • electrons repel and so less energy is needed to remove the electron
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31
Q

What is successive ionisation energy?

A
  • 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
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32
Q

Describe the structure and properties of graphite (6)

A
  • 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
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33
Q

Describe the structure and properties of diamond (5)

A
  • 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
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34
Q

Describe the structure and properties of silicon

A
  • same structure as diamond so same properties
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35
Q

Describe the structure and properties of graphene and its uses

A
  • 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
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36
Q

Why are metals good thermal conductors?

A

Giant metallic structures are good thermal conductirs as delocalised electrons can transfer kinetic energy

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37
Q

What do group 2 oxides look like?

A

White solids

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38
Q

Why do group 2 oxides become more alkaline down the group?

A
  • the hydroxide formed when they react with water are more Soluble so more OH- ions are formed
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39
Q

Define electronegativity

A
  • ability for an atom to attract electrons toward itself in a covalent bond
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40
Q

How does electronegativity change down the halide group?

A
  • decreases as atoms get larger and the distance between positive nucleus and bonding electrons increases + more shielding
  • less strongly attracted
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41
Q

How is bleach made?

A

2NaOH + Cl2 -> NaClO +NaCl + H2O
a disproportionation reaction

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42
Q

What is the equation when chlorine reacts with water and why is this useful?

A

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

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43
Q

What are the advanatges and disadvantages of chlorinating water?

A

:)
- 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)

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44
Q

Why may some people object to chlorination of water?

A
  • occurs across the UK and we have no choice
  • some say it is forced medication of a whole population
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45
Q

Give the 2 alternatives to chlorination of water

A
  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
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46
Q

How do you test for ammonium compounds?

A
  • add sodium hydroxide + gently heat = ammonia gas produced
  • damp red litmus will turn blue
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47
Q

Define standard enthalpy change of formation

A
  • enthalpy change when 1 mole of a compound is formed from its elemnts in their standard states, under standard conditions
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48
Q

Define standard enthalpy change of neutralisation

A
  • enthalpy change when an acid and alkali react to form 1 mole of water, under standard conditions
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49
Q

Define enthalpy change of combustion

A
  • enthlapy change when 1 mole of a substance is completely burned in O2 to make CO2 and H2O under standard conditions
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50
Q

What are the standard conditions?

A
  • 100kPa
  • 298K/ 25ºC
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51
Q

How to detrmine if a reaction is endo/exothermic from bonds made/broken

A
  • 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
Q

Define average bond enthalpy

A
  • the energy required to break 1 mole of bonds between two atoms in gaseous state
53
Q

Define actual bond enthalpy

A
  • 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
Q

What is the difference between actual bond enthalpy and average bond enthalpy?

A
  • 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
Q

What are the benefits of using catalysts?

A
  • lower temperautrues and pressures required = less money spent and CO2
    -less waste is produced allowing scientists to use reeactions w/ better atom economies
56
Q

What are the ways to measure rate of reaction?

A
  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
Q

Define dynamic equilibrium

A
  • 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
Q

What kind of equilibrium does Le Chatelier’s principle work for?

A
  • homogenous equilibria where reactants and products are in the same state
59
Q

What happens when you decrease/ increase pressure in an equilibrium?

A
  • 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
Q

What is the compromise between when using pressures and temperatures for industrial reactions?

A
  • temperature compromises between yield and rate
  • pressure compomises between yield/speed and cost + safety
61
Q

How does temperature affect Kc?

A
  • 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
Q

How can we determine where equilibrium lies from Kc value?

A
  • 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
Q

How does water dissociate?

A
  • 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
Q

Describe Kw

A
  • 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
Q

How do we work out pH from mono/diprotic acids?

A

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
Q

How do we figure out pH of strong bases?

A
  • assume they dissociate fully
  • the [OH-] value will be the same as the conc.
67
Q

Describe how you would use a pH probe

A
  • 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
Q

What is the equivalence point/ end point?

A
  • the acid has been neutralised fully by the base
  • we assume [H+]=[OH-]
  • on the verticle section of the graph in the middle
69
Q

How do you choose a suitable indicator?

A
  • must change colour entirely w/in the vertical part of the titration
70
Q

What titrations can methyl orange be used for?

A
  • strong acid/ strong base
  • strong acid/ weak base
71
Q

What titrations can phenolphthalein be used for?

A
  • weak acids/ strong base
72
Q

What are titrations are indicators not used for?

A
  • weak base/ weak acid titrations as no sharp pH change
73
Q

Define buffers

A
  • 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
Q

What happens when we add an acid to a buffer?

A

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
Q

What happens when we add a base to a buffer?

A
  • 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
Q

Describe the use of buffers in maintaining blood pH

A
  • 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
Q

Describe the iodine clock experiment

A
  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
Q

Define orders of reactions

A
  • an order is the power to which a concentration is raised to in the rate equation
  • tells us how conc affects rate
79
Q

Define 0 order, 1st order and 2nd order

A

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
Q

What do the rate-concentration graphs for 0 order, 1st order and 2nd look like?

A

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
Q

What does a conc-time graph look like for first and zero order?

A

zero= straight negative line
first= curved negative line

82
Q

How do you work out k from half life of a 1st order conc-time graph?

A
  • k= ln(2/time of half life)
  • units are s-1
83
Q

What is the rate determining step?

A
  • slowest step in a multi-step reaction
    the whole reaction rate depends on how quick the rate determing step is
84
Q

What happens to the rate constant when Ea is decreased?

A
  • rate constant gets bigger as rate of reaction increases due to more particles having enough energy to collide and react
85
Q

What happens to the rate constant when temperature increases?

A
  • k increases because particles have more kinetic energy and more likely to collide with at least the Ea
  • therefore rate of reaction increases
86
Q

When using the Arrhenius equation, how can we rearrange it to form y=mx+c?

A

lnk= -Ea/R x 1/T +lnA
- lnk=y
- -Ea/R=gradient
- 1/T= x
- lnA=c

87
Q

How does temperature affect Kp?

A

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
Q

How does pressure affect Kp?

A
  • it doesn’t because the partial pressure ratio of reactant to product stays the same
89
Q

Define lattice enthalpy

A
  • formation of 1 mole of an ionic lattice from gaseous ions
  • measure of the strength of ionic bonding in a giant ionic lattice
90
Q

Define enthalpy change of 1st ionisation

A
  • enthaply change when 1 mol of gaseous 1+ ions are made from 1 mol of gaseous atoms
91
Q

What affects the strength of ionic bonding?

A
  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
Q

What is charge density?

A
  • high charge density means ions have a high charge and small ionic radii
93
Q

Define enthalpy change of solution

A
  • 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
Q

What must happen for a substance to dissolve?

A
  1. substance bonds must break (endothermic)
  2. new bonds form between the solvent and the substance (exothermic)
95
Q

Define enthalpy change of hydration

A
  • when 1 mol of aqueous ion is made from 1 mole of gasous ions
96
Q

What affects enthalpy change of hydration?

A
  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
Q

Define entropy

A
  • 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
Q

How will a reaction tend towards?

A
  • a reaction will tend towards more disorder and hence an increase in entropy
99
Q

What are the conditions for standard entropy?

A
  • 1 mol of substance
  • 100kPa
  • 298K
100
Q

Describe Gibbs Free Energy

A
  • ΔG= ΔH-TΔS
  • all must be in Jmol-1, K and JK-1mol-1
  • ΔG is Jmol-1
101
Q

Describe feasibility of a reaction

A
  • 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
Q

How can we work out the lowest temperature at which a reaction becomes feasible?

A
  • when ΔG=0 a reaction is just feasible
  • therefore T=ΔH/ΔS
103
Q

Where do electrons flow from in electrochemical cells?

A
  • flow from more reactive metal to a less reactive one
104
Q

What are electrode potentials?

A
  • 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
Q

What is the standard Hydrogen electrode?

A
  • 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
Q

How do we simplify how we set up a cell

A
  • 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
Q

How do we predict feasibility of a reaction using electrode potentials?

A
  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
Q

What is a transition element? What elements behave differently when electrons fill orbitals?

A
  • 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
Q

What elements are in the d block but not transition elements and why?

A
  • 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
Q

How are electrons removed from transition metals?

A
  • lost from 4s before 3d
111
Q

Why are transition metals good catalysts?

A
  • 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
Q

What are the risks of using transition elements as catalysts?

A
  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
Q

What is a complex ion?

A
  • ## where a central transition metal ion is surrounded by ligands bonded by dative covalent (coordinate bonds)
114
Q

What are ligands?

A
  • 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
Q

What is the coordination number and what shapes do they form?

A
  • 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
Q

Describe haemoglobin

A
  • 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
Q

Describe how haemoglobin works and carries O2

A
  • 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
Q

Why does CO affect haemoglobin’s ability to carry O2?

A
  • 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
Q

Describe how cis-platin acts as an anti cancer drug

A
  • 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
Q

How would you work out the overall order of the reaction?

A
  • use the equation k=rate/Reactant order to try out if the units match
121
Q

What is meant byu the term weighted mean mass?

A
  • the mean mass taking the relative abundancies of isotopes into account
122
Q

What are the signs of ΔH and ΔS to make a reaction feasible?

A
  • ΔH should be negative and ΔS positive so that ΔG is below 0 at all temperatures
123
Q

Explain why rubidium chloride is reasonably soluble at room temperature

A
  • 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
Q

What is the purpose of the salt bridge?

A
  • completes the circuit by allowing ions to move between the half cells
  • prevents the two solutions mixing together
125
Q

Describe a fuel cell

A
  • 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

Chemistry (4 decks)

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