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ocr chem > module 3 - periodic table and energy > Flashcards

module 3 - periodic table and energy Flashcards

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1
Q

periodicity definition

A

repeating pattern of physical and chemical properties

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2
Q

how is the periodic table arranged

A
  • by increasing atomic number of elements
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3
Q

how are periods and groups organised

A

periods = similar chemical and physical properties
groups = similar chemical properties

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4
Q

where are the s,d,p and f blocks located

A

s = left
p = right
d = middle
f = bottom

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5
Q

electron shielding definition

A

electron shielding is the decrease in the attraction of the outer shell electrons in the nucleus.
the greater the number of shells, the greater the electron shielding

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6
Q

how does electrons shielding affect ionisation energy

A

when electron shielding increases, ionisation energy decreases as it requires less energy to remove the outer shell electron

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7
Q

how does electron shielding differ across a period

A

electron shielding remains the same across a period because each period has the same number or shells

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8
Q

how does electron shielding differ down a group

A

electron shielding increases down the group because there are more shells as you go down the group

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9
Q

why are there small decreases in first ionisation energy in periods 2 and 3

A

-magnesium - the outer electrons are in a 3s orbital, whereas in aluminium the outer electron is in a 3p orbital. The 3p orbital is further from the nucleus so there is more electron shielding and distance from the nucleus so electrostatic attraction decreases. The second drop is from Phosphorus to Sulphur. This is because it is the first time electrons are paired up in 3p orbitals, so there is added electron-electron repulsion.

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10
Q

first ionisation energy definition

A

the energy required to remove one electron from each atom in one mole of gaseous atoms

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11
Q

how does atomic radius differ across a period

A

atomic radius decreases because:
- nuclear charge increases
- electron shielding remains the same

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12
Q

how does atomic radius differ down a group

A

atomic radius increases because:
- increased number of shells
- increased shielding

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13
Q

trend in first ionisation energy across periods 2 and 3

A
  • first ionisation energies will increase
  • as you go along a period, atomic number increases (number of protons in nucleus) which means that atomic radius decreases and so there is a stronger electrostatic attraction between electrons and the nucleus
  • however there are small decreases in the first ionisation energy
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14
Q

why are there small decreases in first ionisation energy across periods 2 and 3

A
  • highest energy electron occupying a p orbital, which is slightly higher in energy than an s orbital.
  • so there is a dip in first ionisation energy
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15
Q

metallic bonding definition

A

strong electrostatic attraction between cations (positive ions)
and delocalised electrons

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16
Q

what structure do metals form

A

giant metallic lattice structures

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17
Q

components of giant metallic lattice structures

A
  • each metal atom forms a positive ion
  • the positive ions are arranged into a regular lattice structure
  • outer shell electrons are delocalised which can move through the structure, and conduct electricity
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18
Q

in which part of the periodic tables are giant covalent lattices formed

A

right side

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19
Q

mtp and btp of giant covalent lattices

A
  • very high
  • large amount of energy needed to break strong covalent bonds between atoms
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20
Q

what lattices does carbon form

A

diamond, graphite and graphene

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21
Q

structure of diamond and silicon

A
  • form giant 3D structures with atoms bonded in a tetrahedral arrangement by covalent bonds
  • all 4 outer shell electrons involved in covalent bonding so the electrons cannot move and are unable to conduct electricity
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22
Q

structure of graphite

A
  • forms a giant planar structure with many planes weakly held together by covalent bonds
  • 3 outer shell electrons involved in covalent bonding within each layer, which means the outer shell can move and conduct electricity
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23
Q

graphene structure

A

single layer of atoms held together by covalent bonds

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24
Q

trend in melting points for giant metallic lattices

A
  • giant metallic lattice held together by strong metallic bonds between positive ions and delocalised electrons
  • large amount of energy needed to break the metallic bonds and melting points are high
  • melting point increases from lithium-beryllium and from sodium-magnesium-aluminium because the charge on the positive ion and number of delocalised electrons both increase.
  • attraction between particles increase and more energy is needed to break the metallic bonds
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25
Q

trend in melting points for giant covalent lattices

A
  • lattice held together by strong covalent bonds between atoms
  • large amount of energy is needed to break the covalent bonds and the melting points are high
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26
Q

trend in melting points for simple molecular lattices

A
  • weak london forces between molecules hold the lattice together
  • small amount of energy is needed to break the london forces and the melting points are low
  • fluctuations in melting points of multiple elements because london forces increase with the number of electrons in the molecules so it requires more energy to break those bonds
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27
Q

solubility of giant covalent and metallic structures

A
  • A substance can dissolve. in water if it forms strong enough attractions with water molecules. Giant covalent and metallic substances cannot form these strong attractions with water, so they are insoluble.
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28
Q

why is there decreases in mpts and bpts across periods 2 and 3

A
  • as we go across these periods, we are met with the metals first. for metals they have a large number of delocalised electrons which increases the strength of the metallic bonds that hold the metal together. because the metallic bonds are so strong, the melting point increases as it requires more energy to break these bonds.
  • once we reach the non metals in periods 2 and 3, the mpt decreases significantly
  • non metals arent held together by metallic bonds and in the case of oxygen and fluorine, their molecules are held together by weak dispersion forces. and in neon, the atoms are held together by weak dispersion forces. because these forces are so weak, it requires little energy to break the bonds
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29
Q

successive ionisation energy definition

A
  • the energy that is required to remove the electron one after the other.
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30
Q

what does successive ionisation energy depend on

A
  • the number of electrons present in the outermost shell.
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31
Q

why does successive ionisation energy increase between shells

A
  • because the closer the shell is to the nucleus, the easier it is to remove an electron from that shell
  • for example, in sodium, it has one outer shell electron and there is a big jump in ionisation energy from the first to second. This is because the second electron is being removed from a shell that is much closer to the nucleus, meaning there is stronger attraction from the nucleus.
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32
Q

atomic radius trend as you go down a group

A

increases as you go down a group in the periodic table because extra electron shell is added each time, so the distance between the nucleus and outermost electron increases. This means that the attraction between the outermost electron and nucleus decreases

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33
Q

atomic radius trend as you go across a period

A
  • decreases slightly as you go across a period, as the nuclear charge of the nucleus increases, which means that the nucleus pulls the outer electrons closer to it, which decreases the atomic radius.
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34
Q

trend in successive ionisation energy as you go down a group

A

Ionisation energy decreases down a group because atomic radius increases and the outer electrons are shielded from the nucleus by inner electron shells.

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35
Q

trend in successive ionisation energy as you go across a period

A

Ionisation energy generally increases across a period because nuclear charge increases.

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36
Q

what do group 7 elements exist as

A

diatomic molecules

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37
Q

mpts and bpts of group 7

A

strong covalent bonds between atoms, weak intermolecular forces, so low mpts and bpts

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38
Q

trend in mpts and bpts as you go down group 7

A

bpts and mpts increase as you go down the group - larger atoms = more electrons = stronger london forces = more energy needed to break these bonds

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39
Q

reactivity trend as you go down group 7

A

reactivity decreases as you go down
more shells = more shielding effect = less nuclear attraction

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40
Q

benefits and risks of chlorine water

A
  • can kill bacteria in water
  • chlorine is toxic and respiratory irritant
  • reacts with organic compounds to make haloalkanes - can be carcinogenic
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41
Q

redox reactions in group 7

A
  • most common reaction in group 7
  • each halogen becomes reduced (gains an electron)
  • behaves as an oxidising agent - removes electron from other substances
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42
Q

disproportionation definition

A

the reaction in which the same element becomes simultaneously oxidised and reduced

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43
Q

examples of disproportination reactions

A

chlorine + water
Cl2 + H2O –> HCl + HClO

chlorine with sodium hydroxide
Cl2 + 2NaOH —> NaClO + NaCl + H2O

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44
Q

test for halide ions

A

method:
- take 5cm^3 of unknown solution
- add 3-5 drops of nitric acid followed by 3-5 drops of silver nitrate solution

positive result:
- chloride = white precipitate
- bromide = cream precipitate
- iodide = yellow precipitate

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45
Q

test for ammonium ion

A

method:
- take 5 cm^3 of unknown
- add 1cm^3 of sodium hydroxide solution and warm the mixture
- test any gas evolved with damp red litmus paper

positive result:
- litmus paper turns blue

46
Q

redox reactions in group 2

A
  • become reducing agents because it reduces other elements (gives electrons)
47
Q

reactions of group 2 with water

A

forms hydroxides

48
Q

trend in alkalinity of group 2

A

as you go down the group:
- alkalinity increases
- pH increases
- solubility increases

49
Q

use of group 2 compounds in agriculture

A

calcium hydroxide:
- added to fields to increase the pH of acidic soils
- neutralises acid in soil forming water

50
Q

use of group 2 compounds in medicine

A

used as antacids in treating indigestion
contains magnesium and calcium carbonates
acid is mainly HCl so it neutralises it

51
Q

equation to calculate enthalpy change

A

ΔH = Q / n

52
Q

equation to calculate heat loss/gain

A

Q = mcΔT

53
Q

what does it mean if the enthalpy change is negative

A

exothermic reaction

54
Q

what does it mean if the enthalpy change is positive

A

endothermic reaction

55
Q

what goes on x and y axis on an energy profile diagram

A

x = progress of reaction
y = energy

56
Q

construction of an energy profile for an exothermic reaction

A

exothermic = reactants above products

57
Q

construction of an energy profile for an exothermic reaction

A

exothermic = reactants over products

58
Q

what happens to atoms during a chemical reaction

A

break reactant bonds, shuffle, make product bonds

59
Q

energy change units

A

kilo joules per mole
kJmol-1

60
Q

where is activation energy on an exothermic enthalpy change diagram

A

from reactants line, up to the peak

61
Q

qualitative test for halide ions

A
  • make solution of halide
  • acidify with dilute nitric acid
  • add a few drops of silver nitrate solution
  • precipitate of silver halide will be formed

silver chloride: white precipitate
silver bromide: cream precipitate
silver iodide: yellow precipitate

62
Q

qualitative test for halide ions pt 2

A

add ammonia as the white, cream and yellow precipitates can look very similar in colour
- add dilute NH3
- if precipitate dissolves in dilute ammonia, unknown halide is chloride
- if nothing happens, add concentrated NH3
- if precipitate dissolves in concentrated NH3, unknown halide is bromide
- if precipitate doesn’t dissolve in either NH3, then unknown halide is iodide

63
Q

if precipitate dissolves in dilute NH3, unknown halide is:

A

chloride

64
Q

if precipitate dissolves in concentrated NH3, unknown halide is:

A

bromide

65
Q

if precipitate doesn’t dissolve in either NH3, then unknown halide is:

A

iodide

66
Q

qualitative test for carbonates

A
  • add small amount of dilute hcl into test tube
  • equal amount of sodium carbonate solution added to test tube
  • as soon as sodium carbonate solution is added, bung with delivery tube should be attached to the test tube
  • delivery tube should transfer gas which is formed into a different test tube that contains a small amount of limewater
  • carbonate present = effervescence and white precipitate formed
67
Q

qualitative test for sulfates

A
  • add dilute hcl to sample and add a few drops of barium chloride
  • sulfate present = white precipitate
  • or use barium nitrate solution
68
Q

qualitative test for ammonium

A
  • react sample with warm aqueous sodium hydroxide
  • ammonium gas present = pungent smell/ turns red litmus paper blue
69
Q

standard conditions

A
  • 25°C / 298K
  • 101KPa / 1 atm
  • 1 mol/dm³
70
Q

activation energy:

A
  • the minimum energy required for a chemical reaction to occur.
  • the energy needed to break existing bonds in reactants and form new bonds in products.
  • in order for a reaction to occur, colliding molecules must have enough energy to overcome activation energy barrier.
71
Q

enthalpy change of reaction

A

the overall energy taken in from / given out to the surroundings OR the energy difference from reactants to products.

72
Q

enthalpy change of formation

A

enthalpy change when 1 mole of a compound is formed from its elements under standard conditions

73
Q

enthalpy change of combustion

A

enthalpy change when 1 mole of a substance is reacted completely with oxygen under standard conditions

74
Q

enthalpy change of neutralisation

A

enthalpy change when 1 mole of water is formed in the reaction between an acid and a base under standard conditions

75
Q

2 equations used to calculate enthalpy change

A

Q = mcΔT
ΔH = Q/n

76
Q

average bond enthalpy

A

average enthalpy change when 1 mole of gaseous covalent bonds is broken

77
Q

disadvantage of term ‘average bond enthalpy’

A

bond enthalpies are only averages and in reality, bond enthalpies can vary depending on the molecule

78
Q

exothermic and endothermic reactions in terms of breaking and making chemical bonds

A

chemical bonds broken first (endothermic), atoms shuffle and new bonds are formed (exothermic)

79
Q

overall energy change depends on:

A

chemical bonds which have to be broke, and the bonds which are formed

80
Q

equation to calculate energy change using bond enthalpies

A

break - make

81
Q

what goes at the bottom of the triangle in a enthalpy change of reaction using bond enthalpies

A

gas atoms

82
Q

what goes at the bottom of the triangle in a enthalpy change of reaction using formation data

A

elements

83
Q

what way do the arrows face in an enthalpy change of reaction using bond enthalpies

A

down

84
Q

what way do the arrows face in an enthalpy change of reaction using formation data

A

up

85
Q

what way do the arrows face in an enthalpy change of reaction using combustion data

A

down

86
Q

collision theory

A

for a reaction to take place, the reactant particles must collide with energy greater than or equal to the activation energy

87
Q

2 conditions that must be met for an effective collision

A
  • particles must collide with the correct orientation
  • particles must have sufficient energy to overcome activation energy barrier
88
Q

how does concentration affect rate of reaction

A
  • more concentrated = greater number of particles in given volume of solvent
  • increase in concentration causes increased collision frequency and therefore increased rate of reaction
89
Q

how does pressure affect rate of reaction

A
  • increase in pressure = molecules have less space to move as there is so many particles
  • number of effective collisions increase due to increased collision frequency
  • increase in pressure = increased rate of reaction
90
Q

what is a catalyst and what does it do

A
  • a substance that increases the rate of reaction without being consumed itself
  • a catalyst can form an ‘intermediate’ with a reactant or it may simply provide a surface on which the reaction takes place
  • provides an alternate pathway of lower activation energy for reaction
  • catalyst regenerated at end of reaction
91
Q

2 types of catalyst

A
  • homogenous
  • heterogenous
92
Q

homogenous catalyst

A
  • when catalyst is in the same physical state as the reactants
  • the catalyst forms an intermediate with the reactants then the intermediate breaks down to form a product
93
Q

heterogenous catalyst

A
  • when catalyst is in a different physical state as the reactants
  • usually the catalyst is a solid and reactants are gases
  • the gases adsorb to the surface of the solid catalysts where they react
  • the products form then desorb from surface
94
Q

catalysts and sustainability

A
  • hugely important in industry
  • save energy, resources, time, fossil fuels, money
  • lowers temperatures and reduces energy demand from combustion of fossil fuels which results in reduction in CO2 emissions
  • processes with high atom economy can be used
  • also speed up chemical processes for greater efficiency and profitability
  • catalysts regenerated so can be used many times
95
Q

what goes on x and y axis of boltzmann distribution curve

A

x = energy
y = number of molecules with a given energy

96
Q

what does the boltzmann distribution curve show

A
  • shows the energy of particles and the number of particles with a given energy
97
Q

what to remember about the boltzmann distribution

A
  • no molecules have 0 energy so curve always starts a bit forward on x axis
  • area under the curve is the total number of molecules
  • there is no maximum energy of molecules, the curve doesn’t meet x axis
98
Q

how is boltzmann distribution curve affected by increasing temperature

A
  • curve has a lower peak, and has a higher surface area under graph
  • shifted to the right
  • at higher temperatures, greater proportion of particles can overcome Ea
99
Q

how is boltzmann distribution curve affected by a catalyst

A
  • peak is higher
  • shifted a bit more to the right
  • lower Ea
  • curve reaches the ‘bottom’ at a steeper decline compared to graph without catalyst
100
Q

conditions for dynamic equilibrium

A

has to happen in a closed system (nothing can enter or leave reaction)

101
Q

how does a dynamic equilibrium occur

A

when rate of forwards reaction is equal to the rate of the reverse reaction
( reactants making products at the same rate products are making reactants - concentrations don’t change)

102
Q

le chatelier’s principle

A
  • principle states that if a change is applied to a system in equilibrium, the equilibrium will shift to undo that change
  • if we increase pressure, equilibrium will shift to decrease pressure (shifts to side with less molecules)
  • if we decrease temperature, equilibrium will shift to increase temperature ( shifts in endothermic direction)
103
Q

how does a catalyst effect dynamic equilibrium

A
  • increases rate of forwards and backwards reaction by same amount resulting in unchanged position of equilibrium
104
Q

explain importance of compromise between chemical equilibrium and reaction rate in deciding operational conditions to chemical industry e.g haber process

A
  • le chatelier’s principle can be used to predict best conditions of temp and pressure to force position of equilibrium to the right to produce max yield of ammonia
  • ideal conditions for max ammonia yield = low temp and high pressure
  • iron catalyst also used
  • could use low temp but very slow rate of reaction so high temp needed to increase frequency of collisions
  • high pressure = high yield but its a safety risk and very expensive to generate
105
Q

what is Kc

A

equilibrium constant

106
Q

what does [ ] mean

A

concentration of

107
Q

Kc =

A

[products] / [reactants]

108
Q

if Kc = 1, this means that:

A

[products] = [reactants]

109
Q

if Kc is very small:

A

[products] &laquo_space;[reactants]

110
Q

if Kc is very big:

A

[products]&raquo_space; [reactants]

111
Q

Kc units =

A

[products unit] / [reactants unit]

ocr chem (8 decks)

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