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Chemistry > Module 3 Periodicity > Flashcards

Module 3 Periodicity Flashcards

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1
Q

[‘who created the periodic table

A

mendeleev

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2
Q

how did mendeleev order the elements in the periodic table

A

in order of atomic mass

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3
Q

how else did mendeleev order the elements in the periodic table

A

groups of similar properties

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4
Q

why did mendeleev leave gaps in the periodic table

A

he assumed some elements were yet to be discovered

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5
Q

how did he predicts properties of missing elements

A

from group trends

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6
Q

how are elements arranged in the periodic table now

A

linked to their physical and chemical propeties

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7
Q

what is the period (row) of a modern periodic table measured from

A

from left to right the elements are arranged in order of increasing atomic number

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8
Q

what is the group of the modern periodic table measured in

A

every element in the group has the same amount of electrons in the outer shell

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9
Q

what does the number of the period mean

A

gives the number of the highest energy electron shell in an element’s atom

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10
Q

what is periodicty

A

a repeating trend in properties of the elements through each period`

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11
Q

describe the trend in electron configuration

A

across period 2 - the 2s sub-shells fills with two electrons, followed by the 2p sub-shell with six electrons
across period 3 - same pattern but also filling 3s and 3p sub shells

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12
Q

what are the 4 blocks in the periodic table

A

s p d f

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13
Q

what are the blocks

A

elements divided corresponding to their highest energy sub-shell

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14
Q

what is ionisation energy

A

measures how easily an atom loses electrons to form positive ions

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15
Q

what is the first ionisation energy

A

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

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16
Q

what has the nucleus got to do with ionisation energy?

A

electrons are held in their shells by attraction from the nucleus, so the first electron lost is the highest energy level and will have the least attraction to the nucleus

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17
Q

what three factors affect the attraction between the nucleus and outer electrons of an atom

A

atomic radius
nuclear charge
electron shielding

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18
Q

how does the atomic radius affect ionisation energy

A

the greater the distance between the nucleus and the outer electrons, the less the nuclear attractiom`

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19
Q

how does nuclear charge affect ionisation energy

A

the more protons there are in the nucleus of an atom, the greater the attraction between the nucleus and the outer electrons

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20
Q

how does electron shielding affect ionisation energy

A

electrons are negatively charged and so inner-shell electrons repel outer-shell electrons.w

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21
Q

what is the shielding eefect

A

the effect of when inner shell electrons repel outer shell electrons which reduce the attraction between the nucleus and the outer electrons.

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22
Q

what is the second ionisation energy

A

the energy required to remove one electron from each atom in one mole of gaseous 1+ ions of an element to form one mole of gaseous 2+ ions

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23
Q

what is the trend in successive ionisation energies

A

increase of energy is required

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24
Q

explain the trend of successive ionisation energies

A
  • electron ratio increases and the ion gets smaller, attraction of electron to nucleus increases
  • the big jump suggests it is removing an electron closer to the nucleus
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25
Q

what predictions can you make from successive ionisation energies

A

the number of electrons in the outer shell
the group of the element in the periodic table
the identity of an element

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26
Q

what are the trends of ionisation energies in the periodic table down the group

A

ionisation energy decreases
electron being removed is further away from the nucleus

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27
Q

what are the trends of ionisation energies in the periodic table across a period

A

ionisation energy increases
atoms are smaller so attraction gets bigger

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28
Q

explain anomalies across a period for ionisation energies

A

Be>B - removal of p-electron for Boron which is higher in energy than the s-electron removed from berrillium
N>O - there is an electron pair p-electron in oxygen, removal requires less energy, then removing the p-electron due to the electron repulsion

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29
Q

what is the constant property for all metals

A

able to conduct electricity

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30
Q

what is `metallic bonding

A

the strong electrostatic attraction between cations and delocalised electrons

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31
Q

how do solid metal structures get delocalised electrons

A

each atoms has donated its negative outer-shell electron to a shared pool of electrons (delocalised)

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32
Q

describe the cations and delocalised electrons in metallic bonding

A

cations are in a fixed position (which maintains the structure and shape of the metal)
the delocalised electrons are mobile and are able to move throughout the structure

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33
Q

what properties do metals have

A

strong metallic bonds
high electrical conductivity
high melting and boiling points

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34
Q

explain electrical conductivity in metals

A

when a voltage is applied across a metal, the delocalised electrons can move through the structure, carrying charge

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35
Q

explain melting and boiling points in metals

A

most have a high melting and boiling points
- melting point depends on the strength of the metallic bonds holding together the atoms in the giant metallic lattice
- high temps are needed to break the strong electrostatic attraction

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36
Q

describe solubility in metals

A

metals do not dissolve
- any interaction between polar solvents and the charges in a metallic lattic would just cause a reaction, instead of dissolving

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37
Q

what is a giant covalent lattice

A

many billions of atoms are held together by a network of strong covalent bonds

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38
Q

what are the properties of a giant covalent lattice

A
  • high melting and boiling points (strong covalent bonds)
  • insoluble (covalent bonds way too strong to be broken by solvents)
  • non-conductors of electricity (apart from graphite and graphene)
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39
Q

what is graphene

A

a single layer of graphite composed of hexagonally arranged carbon atoms linked by strong covalent bonds

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40
Q

what is graphite

A

composed of parallel layers of hexagonally arranged carbon atoms , like a stack of graphene layers
- layers are bonded by london forces
bonding only uses three of carbon’s four outer-shell electrons
- spare electron is delocalised between the layers, so electricity can be conducted as in metals

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41
Q

describe melting points across period 2 and 3

A

melting point increases from group 1 to 14
there is a sharp decrease in melting point between group group 14 and 15
melting points are rather low from group 15 to geoup 18

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42
Q

how many outer shell electrons does group 2 elements have?

A

2

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43
Q

what block are group 2 outer shell electrons in?

A

the s-block

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44
Q

what is the most common reaction in group?

A

redox reaction

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45
Q

show the redox reaction of calcium

A

Ca(s) → Ca2+(aq) + 2e-.

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46
Q

in a redox reaction is calcium being oxidised or reduced

A

oxidised

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47
Q

in the equation Ca(s) → Ca2+(aq) + 2e- describe what would happen

A

another species will gain these 2 electrons and will be reduced
the group 2 element is called a reducing agent because it has reduced another species

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48
Q

what is a reducing agent

A

when an element reduces another species

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49
Q

what happens when a group 2 element reacts with oxygen

A

it forms a metal oxide

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50
Q

what happens when magnesium and oxygen react

A

magnesium burns with a brilliant white light and forms white magnesium oxide

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51
Q

what happens when a group 2 element reacts with water

A

an alkaline hydroxide, M(OH)2 and hydrogen gas

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52
Q

describe the reactivity of group 2 elements with water

A

increases when you go down the group

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53
Q

what is formed when a group 2 element reacts with dilute acid

A

a salt and hydrogen gas

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54
Q

show the equation of redox reaction of a group 2 element and a dilute acid

A

metal+acid -> salt + hydrogen

55
Q

does the reactivity increase or decrease down the group when a group 2 element reacts with dilute acid

A

increases

56
Q

why do the group 2 element increase in redox reactions when you go down a group

A

the atoms of group 2 elements react by losing electrons to form +2 ions. the formation of 2+ ions from gaseous atoms require the input from gaseous atoms requires the input of two ionisation energies

57
Q

what happens when group 2 element oxides react with water,

A

releases hydroxide ions OH- and forming alkaline solutions of the metal hydroxide

58
Q

are group 2 hydroxides soluble

A

only slightly soluble in water

59
Q

what happens when group 2 hydroxide solution becomes saturated

A

any further metals and hydroxide ions will be a solid precipitate

60
Q

does the solubility of hydroxides increase down the group

A

yes, the resulting solutions contain more OH- ions and more alkaline

61
Q

how do you show that solublity, ph and alkalinity increase down the group 2 group

A
  1. add a spatula of each group 2 oxide to water in a test tube
  2. shake the mixture
62
Q

what can group 2 compounds be useful in

A

agriculture and medicine

63
Q

what commpound is used in agriculture

A

calcium hydroxide

64
Q

what does calcium hydroxide do to fields

A

increases pH of acidic soil

65
Q

what are the halogens

A

group 7 (17)

66
Q

name one thing about group 7

A

they are the most reactive non-metallic group

67
Q

describe characteristics of halogens on earth

A

halogens occur as stable hallide ions (Cl-, Br-, I-) dissolved in sea water or combined with sodium or potassium

68
Q

at room temperature what do halogens exist as

A

diatomic molecules (X2)

69
Q

what does flourine appear as at room temperature

A

pale yellow gas

70
Q

what does chlorine appear as at room temperature

A

pale green gas

71
Q

what does bromine appear as at room temperature

A

red-brown liquid

72
Q

what does iodine appear as at room temperature

A

shiny grey-black solid

73
Q

what does astatine appear as at room temperature

A

never been seen

74
Q

describe the trend when you down the group of the halogens

A
  • more electrons
  • stronger london forces
    -more energy required to break the intermolecular forces
    -boiling point increases
75
Q

how many outer shell electrons do halogens have?

A

7

76
Q

describe the electronic configuration of the halogens

A

2 outer electrons in the s-sub shell, and 5 are in the outer p-sub shell

77
Q

what is the most common reaction in the halogens

A

redox reaction

78
Q

what happens to the halogen in a redox reaction

A

each halogen atom is reduced, gaining one electron to form a 1- halide ion with the electronic configuration of the nearest noble gas

79
Q

show an example of a redox reaction with a halogen

A

Cl2 +2e- -> 2Cl-
chlorine is reduced

80
Q

why is halogen an oxidising agent

A

another species loses electrons to halogen atoms so the halogen has oxidised another species

81
Q

what shows that the reactivity of halogens decrease down the group

A

displacement reaction of halogens with halide ions can be carried out on a test tube scale and the results support this

82
Q

describe the displacement reactions of the halogens

A

a solution of each halogen is added to aqueous solutions of the pther halides.gi

83
Q

give an example of the displacement reactions of halogens and halide

A

a solution of chlorine (Cl2) is added to two aqueous solutions containing bromine (Br-) and iodine (I-) ions.

84
Q

what would happen if the halogen added is more reactive than the halide in the displacement reaction

A

a reaction takes place and the solution will change colour

85
Q

how would you tell 2 similar coloured solutions apart

A

an organic non-polar solvent such as cyclohexane can be added
the non polar halogens dissolve more readily in cyclohexane than in water

86
Q

show displacement reaction with chlorine, bromine and iodine

A
  • chlorine has clearly reacted with Br- and I-
  • bromine has reacted with I- only
    -iodine has not reacted at all
87
Q

what is the full equation of the reaction of chloride and bromide ions

A

Cl2 + 2NaBr -> 2NaCl + Br2

88
Q

what is the ionic equation of the reaction of chlorine and bromide ions

A

Cl2 + 2Br- -> 2Cl- + Br2
chlroine is reduced
bromide is oxidised

89
Q

describe flourine and astatine

A

-flourine is a pale yellow gas reacting with almost any substance that it comes in contact with
-astatine is extremely rare because it is radioactive and decays rapidly and the element has never actually been seen (predicted to be the least reactive halogen)

90
Q

what happens to trend of reactibity when you go down the group

A

gain an electron decreases and halogen becomes less reactive
- atomic radius increases
- more inner shells so shileding increases
- less nuclear attraction to capture an electron from another species
- reactivity deacreases

91
Q

what is the strongest oxidising agent

A

flourine

92
Q

what is disproportionation

A

a redox reaction in which the same element is both oxidised and reduced

93
Q

show the equation of chlorine and water

A

Cl2 + H2O -> HClO + HCl
0 -> -1.
0 -> +1 .

94
Q

show the equation of chlorine and cold dilute aqueous sodium hydroxide

A

Cl2 + 2NaOH -> NaClO + NaCl + H2O
0 -> -1
0 -> +1 .

95
Q

what are the benefits and risks of chlorine use

A

chlorine ensures that our water is fit to drink and thta the bacteria is killed
however it is also a very fatal gas - it is a respiratory irritant in small concentrations and large concentrations can be fatal

96
Q

what are more risks of chlorine uses

A

chlorine drinking water cna react with organic hydrocarbons such as methane, formed from decaying vegetation, chlorinated hydrocarbons are formed, which are suspected of causing cancer

97
Q

show the equation of aqueuous halude ions reacting with aqueous silver ions to form precipitates of silver halides

A

Ag+ + X- -> AgX

98
Q

what is used to show the reducing ability of halide ions

A

sulfuric acid H2SO4

99
Q

what are the relationships of sulfuric acid and chloride and bromide ions

A

chloride ions are not powerful enough to reduce H2SO4
bromide ions are more powerful and can reduce H2SO4 to sulfur dioxide SO2

100
Q

show the equation of halide test of bromide and sulfuric acid

A

2H+ + H2SO4 + 2Br -> SO2 + Br2 +2H2O

101
Q

what does qualitative analysis rely on

A

simple observations rather than measurements and can often be carried out quickly on a test-tube scale

102
Q

what observations could be made from qualitative analysis

A

gas bubbles, precipitates, colour changes, identification of gases

103
Q

what is the carbonate test

A

carbonate react with acids to form carbon dioxide gas

104
Q

what is the equation for the carbonate test between dilute nitirc acid and aqueous sodium cabronate

A

Na2CO3 + 2HNO3 -> 2NaNO3 + CO2 + H2O

105
Q

describe what you would do to do a cabronate test

A
  1. in a test tube, add dilute nitirc acid to the solid or solution to be tested
  2. if you see bubbles, the unknown compound could be a carbonate
  3. to prove that the gas is carbon dioxide, bubble the gas through lime water, carbon dioxide reacts to form a white precipitate of calcium carbonate which turns the lime water cloudy
106
Q

what is the equation between lime water and carbon dioxide gas

A

CO2 + Ca(OH) -> CaCO3 + H2o

107
Q

what is the sulfate test

A

to test for sulfate ions
- barium ions react with sulfate ions SO42- to form insoluble white barium sulfate

108
Q

what is the equation for the sulfate test

A

Ba2+ + SO42- -> BaSO4

109
Q

describe the sulfate test

A
  1. add a few drops of dilute hydrochloric to the sample
  2. add a few drops of dilute barium chloride or barium nitrate solution
  3. a white precipitate forms if sulfate ions are present
110
Q

describe halide test

A
  1. add aqueous silver nitrate AgNO3, to an aqueous solution of a halide
  2. the silver halide precipitates are different colours - silver chloride in white, silver bromide is cream-coloured and silver iodide is yellow
  3. add aqueous ammonia to test the solubility of the precipitate. This stage is very useful because the three precipitate colours can be difficult to tell apart
111
Q

what is the sequence of tests when asked to analyse an unknown inorganic compound

A
  1. carbonate
  2. sulfate
  3. halides
112
Q

describe why there is a specific order

A
  • carbonate test you add a dilute acid and are looking for effervescence from carbon dioxide gas, none of other tests produce bubbles with dilute acid so if no bubbles can proceed with the next test
  • sulfate test you are looking for a white precipitate, if you carry out a sulfate test on a carbonate you will alos get a white precipitate
  • halide test, silver carbonate and silver sulfate are both insoluble in water and will also form as precipitates in this test
113
Q

how would you perform a carbonate test when there is a mixture of ions

A
  • if you see bubbles, continue adding dilute nitric acid until the bubbling stops, all carbonate ions will then have been removed and there will be none left to react in the next tests
114
Q

how would you perform a sulfate test when there is a mixture of ions

A
  • to the solution left from the carbonate test, add an excess of Ba(NO3)2
  • filter the solution to remove barium sulfate
115
Q

how would you perform a halide test when there is a mixture of ions

A
  • to the solution left from the sulfate test add AgNO3
  • any carbonate or sulfate ions initially present have already been removed, therefore any precipitate formed must involve halide ions
  • add NH3 to confirm which halide you have
116
Q

describe the test for ammonium ion

A

NH4+ + OH- -> NH3 + H2O
1. aqueuous sodium hydroxide NaOH is added to a solution of an ammonium ion
2. ammonia gas is produced
3. mixture is warmed and ammonia gas is released
4. may smell ammonia, easy to test the gas with moist pH indicator paper, ammonia is alkaline so paper will turn blue

117
Q

what is enthalpy (H)

A

a measure of the heat energy in a chemical system

118
Q

what is the chemical system

A

refers to the atoms, molecules or ions making up the chemicals

119
Q

what can enthalpy be also thought as

A

the enrgy stored within bonds, it cannot be measured but enthalpy changes can

120
Q

what is the enthalpy change

A

H(products) - H(reactants)

121
Q

what is the law of conservation of energy

A

fundamental rule in science that states that energy cannot be created or destroyed

122
Q

what has the law of conservation of energy have to do with enthalpy change

A

heat energy has to be transferred between the system and the surroundings

123
Q

what are the two directions of energy transfer

A

from the system to the surroundings - an exothermic change
-from the surroundings to the system - an endothermic change

124
Q

describe exothermic enthalpy change

A

the chemical system releases heat energy to the surroundings
- any energy loss by the chemical system is balanced by the same energy gain by the surroundings
- enthalpy change is negative
- the temperature of the surroundings increases as they gain energy

125
Q

describe endothermic enthalpy change

A

the chemical system takes in heat energy from th eusrroundings
- any energy gain by the chemical system is balanced by the same energy loss by the surroundings
- enthalpy change is positive
- the temperature of the surroundings decreases as they lose energy

126
Q

what is activation energy

A

the minimum energy required for a reaction to take place

127
Q

if a reaction had a small activation energy would the reaction be fast or slow

A

fast because the energy needed to break bonds is available from the surroundings

128
Q

what is standard value of pressure

A

100kPa

129
Q

what is the standard temperature

A

298K (25oC)

130
Q

what is the standard concentration

A

1 mol dm-3

131
Q

what is the enthalpy change of reaction

A

enthalpy change that accompanies a reaction in the molar quantities shown in a chemical equation under standard conditions

132
Q

what is enthalpy change of formation

A

the enthalpy chnage that takes place when one mole of a compound if formed from its elements
(products - reactants)

133
Q

what is the enthalpy change of combustion

A

the enthalpy change that takes place when one mole of a substance reacts completely with oxygen

134
Q
A

Chemistry (4 decks)

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