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Chemistry Module 3 > The periodic table and energy > Flashcards

The periodic table and energy Flashcards

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1
Q

What is the first ionisation energy

A

The energy taken to remove an electron from the outer shell of 1 mole of gaseous atoms and turn them to one mole of gaseous 1+ ions

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2
Q

What are the factors that influence ionisation energy?

A
  • Atomic radius
    the larger the radius the smaller the nuclear attraction
  • Nuclear charge
    the higher the nuclear charge the larger the attractive force
  • Electronic shielding
    inner shells repel outer shells as they are all negative, the more inner shields the smaller the nuclear attraction
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3
Q

Describe the trends across a period

A

the number of protons in nucleus increase so their is a higher attraction on the electrons
electrons added to the same shell so outer shell is drawn inwards this decreases the atomic radii
the same number of inner shells = electronic shielding not changing
- attraction between nucleus and outer electrons increase so more energy is needed to move an electron this increases the 1st ionisation energy

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4
Q

describe trends down a group

A

number of shells increase so the distance from the outer electrons from the nucleus so weaker attraction force on electrons
more inner shells this increases shielding therefore weaker attraction
atomic radius increases
- number of protons also increased but this is outweighed by the other factors therefore decreases in 1st ionisation energy

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5
Q

Why is there a decrease in ionisation energy between groups 2-13? Be and B

A

group 13 has outermost electron in p-orbital whereas group 2 has outermost electron in s-orbital, p-orbitals have higher energy so are further away from nucleus so easier to move this decreases 1st ionisation energy

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6
Q

Why is there a decrease between groups 15-16? N and O

A

group 16 has paired electrons in outermost shell therefore experiences repulsion and is easier to get rid of as it has a lower 1st ionisation energy

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7
Q

What is the quantum number used to describe?

A

they are used to describe the electrons in atoms

The number representing the relative overall energy of each orbital which increases with distance from the nucleus

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8
Q

What are the different types of orbitals?

A

s
p
d
f

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9
Q

What is the shape of an s-orbital?

A

sphere

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10
Q

How many electrons can a p-orbital hold?

A

3

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11
Q

Why are the arrows going the opposite way when drawing an electron?

A

they have opposite spins

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12
Q

How many electrons can a s orbital hold?

A

2

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13
Q

What did aristole believe

A

He believed that the world was made out of 4 elements, earth, air, fire and water. this is similar to the states that we have today - solid liquid air and fire (plasma)

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14
Q

What did scientists begin to understand about the elements

A

noticed that some elements could be more easily extracted such as gold as they were found in their native state

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15
Q

what year and what did Antoine-Laurent de Lavoisier produce

A

1789 he produced the first chemical textbook

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16
Q

What was in Antoine-Laurent de Lavoisier’s book

A

complied list of elements that he believed could not be broken down any further

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17
Q

What was the strengths and Weaknesses of Antoine-Laurent de Lavoisier’s periodic table

A

+ showed a difference between metals and non-metals

  • included some compounds and mixtures
  • included terms such as light and heat
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18
Q

What did Jöns Jakob Berzelius produce and when

A

in 1828 he published a table of atomic weights and determined the mass of many compounds

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19
Q

What did Jöns Jakob Berzelius introduce

A

letter base symbols

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20
Q

What did Johann Wolfgang Döbereiner come up with

A

he noticed that certain groups of three elements called triads ordered by atomic weight would have a middle element with a weight and two properties that were roughly an average of the other two elements for example

  • calcium, strontium and barium
  • chlorine, bromine and iodine
  • lithium, sodium and potassium
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21
Q

What did John Newlands do

A

arranged the table in order of relative atomic weights and in 1865 he suggested the law of octaves

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22
Q

What were the limitations of John Newlands table

A
  • many elements were not known about
  • left no gaps for undiscovered elements
  • mixed metals and non-metals
  • order periods down the table and groups across
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23
Q

What did Dimitri Mendeleev come up with

A

published his table in 1869 this is what the modern periodic table is based of off

24
Q

What were the strengths of Dimitri Mendeleev table

A
  • elements with similar properties were arranged in vertical columns
  • gaps left where no element fitted repeated pattern this predicted the property of the undiscovered element
  • order rearranged where properties did not fit for example iodine and tellurium
  • arrangement of elements of groups of elements in order of atomic weights corresponded to their valencies as well as their distinct chemical properties
25
Q

what did Henry Mosley do?

A

in 1913 Henry Mosely determined the atomic number for all the known elements

  • modified Mendeleevs periodic law to read that properties of the elements vary periodically with their atomic number rather than their weight
  • this put elements such as tellurium and iodine in the correct order as well as argon and potassium and cobalt and nickel
26
Q

What did Glenn Seaborg do ?

A

Seaborg discovered the transuranic elements from 94 plutonium to 102 nobelium
also remodelled the periodic table by placing the actinide series below the lanthanide series
element 106 seaborgium is named after him

27
Q

How are elements arranged

A

arranged in increasing atomic number
groups have similar chemical properties as they have the same number of valence electrons
periods show repeating trends

28
Q

What is periodicity

A

it trend in properties that is repeated across each period

29
Q

What is the period

A

a horizontal row, elements are arranged into groups as they have the same number of valence electrons meaning that they have similar chemical properties

30
Q

What is the group

A

a vertical column

31
Q

What Is the variation in electron structure

A

elements in groups have the same number of electron in their outer shell therefore they have similar chemical behaviour
same orbitals in their group

32
Q

What do group 2 elements make

A

They make hydroxides which are alkaline

they loose two electrons

33
Q

Describe the reactivity trend

A

reactivity increases as you go down the group

34
Q

What are the group 2 physical properties

A 
high melting and boiling points
light metals with low densities
form colourless (white) compounds
35
Q

What are the electronic configuration and ionisation energy

A

they are from the S-block and have the highest energy in in the s-sub shall
- two electrons are lost more easily going down the group therefore reactivity increases going down the group

36
Q

Why does reactivity increase going down group 2 elements

A

as you go down each successive elements has a outer electron in a higher energy level

  • more shielding due to more inner shells
  • larger atomic radii - number of shells increase so the distance from the outer electrons from the nucleus therefore weaker attraction force on electrons
  • number of protons increase but outweighed by the increase in larger atomic radii and shielding, therefore it decreases the ionisation energy and increases the reactivity
  • easier to get rid of 2 electrons in the outer shell
37
Q

Are group 2 elements oxidising agents or reducing agents

A

strong reducing agents with the oxidation states of +2

38
Q

What are the reactions between group 2 elements and oxygen

A

react vigoursly with oxygen to form a metal oxide

39
Q

what are the reactions between group 2 elements and water

A

all except beryllium react with water to form hydroxides as well as a hydrogen gas as you go down the group the elements react more vigrously with water

40
Q

reactions between group 2 elements and dilute acids

A

all group 2 elements except Beryllium react with dilute acids to form a salt and hydrogen, reaction becomes more vigrously going down the group

41
Q

Metal oxide with water reaction equation

A

MO + H2O = M(OH)2

42
Q

solubility in water of group 2

A

soluble and forms an alkaline solution as OH- ions are released

43
Q

what is the group 2 pH

A

between 10-12

44
Q

describe the solubility trend of group 2 elements

A

solubility increases as you go down the group 2 elements this is because it releases more hydroxide ions so there is a more alkaline solution with a higher pH

45
Q

What use does group 2 compounds have

A

mainly use to neutralise acids and are important in the construction industry

  • calcium hydroxide us used by farmers and gardeners as lime is to reduce acidic soil level
  • when suffering from indigestion there is excess acid left in the stomach but having more magnesium hydroxide in the milk which neutralises it and produces a salt and water
  • calcium carbonate building material with limestone but reacts with acid and erodes it
46
Q

What are the properties of the halogens?(group 17)

A

low melting and boiling points

exist as diatomic molecules where X represents the halogen

47
Q

What is the trend in boiling point of the halogens

A

boiling points increase going down the group
- each successive element has an extra shell of electrons therefore more electrons and stronger London forces so more energy is required to break apart the bonds, therefore there is a higher boiling point

48
Q

What is the reactivity trend of the halogens

A

reactivity decreases going down the group

  • atomic radii increases so less nuclear attraction
  • electron shielding increases
  • the ability to gain an electron in the P-sub shell and form a 1- ion decreases
49
Q

What properties do group 17 have

A

strong oxidising agents and form a 1- ion

50
Q

Reaction between magnesium hydroxide and hydrochloric acid

A

Mg(OH)2(s) + 2HCL(aq) = MgCl2(aq) + 2H2O(l)

51
Q

reaction between calcium carbonate and hydrochloric acid

A

CaCO3(s) + 2HCl(l) = CaCl2(aq) + H2O(g) + CO2(g)

52
Q

why is electronic configuration considered a periodic trend

A
  • the pattern repeats itself across each period

- electron sub-shells fill the same order

53
Q

what can be predicted from the successive ionisation energies of an element

A
  • what group the element is in
  • what block the element is in
  • total number of electrons held by an element therefore the actual element itself
  • valance electrons - chemical properties
54
Q

how can successive ionisation energies be used to prove the existence of electron subshells

A
  • large increases in ionisation energy due to number of orbitals decreasing so outer electrons move closer to the nucleus
55
Q

what is the trend in boiling point from left to right across the periodic table

A

metallic to non metallic

56
Q

why can graphite conduct electricity

A
  • not all electrons are involved in bonding
  • each atom has one delocalised electron that can move between the carbon and spread the flow of electricity and conduct electricity

Chemistry Module 3 (6 decks)

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