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Chemistry Module 1 > Periodicity > Flashcards

Periodicity Flashcards

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1
Q

Define “Ionisation Energy” and its use

A

the minimum energy required to remove the most loosely bound or outermost electron from a gaseous atom in its ground state
- useful for determining the properties of the element
- indicator for the strength of the force operating between the nucleus and the electrons

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2
Q

What affects the magnitude of ionisation energy?

A
  • the distance of the electrons from the nucleus
    the size of the nuclear charge
  • the presence of inner shells (shielding)
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3
Q

What is the equipment used to remove an electron?

A

a discharge tube where the electrons always move away from the cathode.
- the electron from the cathode strikes a gaseous atom X and ionises it, and an electron is removed from X and a gaseous ion is formed

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4
Q

How many ionisation energies would an electron with n electrons have?

A

n amount

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5
Q

Why does it take more energy to remove an electron in each successive ionisation energy?

A

it is harder to remove an electron from a positive ion compared to a neutral atom because the electron has more electrostatic attraction and less electron repulsion.

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6
Q

What happens if there is a massive jump in ionisation energy?

A
  1. a cation is removing an electron from a different and stable shell
  2. successive elements where the noble electron configuration with a full shell requires a lot of ionisation energy
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7
Q

Which atoms would be located at the peaks (for ionisation energy)

A

noble or group 7 gases

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8
Q

Which atoms would be located at the troughs (for ionisation energy)

A

group 1 and group 2 metals

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9
Q

How does ionisation energy increase?

A

increase across a period
decrease down a group

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10
Q

Why would oxygen have a lower ionisation than nitrogen and sulfur less than phosphorous?

A

both nitrogen and phosphorous have half-filled p orbitals, which is a stable configuration.

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11
Q

Are the reactivities or metal consitent with the ionisation values>

A

No, because reactivities depend on how metals react with water and acids, as well as how the metals are packed in their lattices, and the energy required to disrupt the lattice.

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12
Q

What are the representative elements?

A

The elements of the main block, Group I - VIII

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13
Q

What are the two types of radioactive elements?

A

synthetic and natural

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14
Q

What is the significance of elements in the same group?

A
  • they have similar chemical properties as they will lose or gain the same number of electrons
  • this is so that they can attain a noble gas configuration
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15
Q

Define “electronegativity”?

A
  • a measure of the electron attracting power of an atom in a molecule
  • determined using the Pauling scale
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16
Q

Define “metallic character”

A

having the physical properties which are usually attributed to metals:
- ductility
- malleability
- hardness
- ability to conduct electricity and heat well

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17
Q

Define “non-metallic character”

A

having the physical properties which are attributed to non-metals:
- being brittle
- poor conductor of heat and electricity

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18
Q

Define “chemical reactivity”

A

the tendency to form compounds and the rate at which this occurs

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19
Q

Define “electron affinity”

A

a measure of the energy released when one electron is added to an atom of an element

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20
Q

Why is there a change from non-metallic to metallic properties as a group descends?

A

increase in shell number means that the valence electrons are located further from the nucleus and therefore more easily lost by atoms

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21
Q

Why is there a change in atomic radius across a period?

A

as the nuclear charge increases, so do the number of protons and therefore pulls the electron shells inward and decreases the atoms size

22
Q

Descending a group, how does ionisation energy change and why?

A

decreases, as the outer electrons are further from the nucleus, therefore less electrostatic attention, less energy needed to lose an electron

23
Q

Descending a group, how does atomic radius and why?

A

increases, new shell of electrons

24
Q

Descending a group, how does electronegativity change and why?

A

decreases, because the atomic radius increases there is less attraction for electrons

25
Q

Descending a group, how does reactivity change and why?

A

increases for metals
decreases for non-metals
- the most reactive metals are the least electronegative
- the most reactive non-metals are the most electronegative

26
Q

Descending a group, how does valency change and why?

A

stays the same for Group I, II, VI, VII, because valence electrons are well shielded from the nucleus, they are lost more easily

27
Q

Across a period, how does ionisation change change and why?

A

increases, as the electrons are smaller

28
Q

Across a period, how does atomic radius change and why?

A

decreases, nuclear charge decreases therefore an the atom’s outer electrons more strongly

29
Q

Across a period, how does electronegativity
change and why?

A

increases, smaller electrons have the highest electronegativity

30
Q

Across a period, how does reactivity change and why?

A

decreases for metals
increases for non-metals

31
Q

Across a period, how does density change and why?

A

increases, because the atomic radius decreases, whilst the atomic mass increases

32
Q

What is electronegativity used for?

A

determining the bonding type of a compound

33
Q

What type of bonding occurs between large differences in electronegativity?

A

ionic, metals tend to have low electronegativity and non-metals will have high electronegativity; the strongest ionic bonds occur between alkali metals and halogens

34
Q

What type of bonding occurs between small differences in electronegativity?

A

covalent, which can be separated into different levels of polarity. molecules can range from very polar to non-polar.

35
Q

What results in a non-polar substance

A

an electronegativity difference of zero

36
Q

Define “polarity”

A

a measure of how evenly the bonding electrons is shared
non-polar: bonding electrons shared evenly
polar: bonding electrons shared very unevenly.

37
Q

What is the diagonal relationship

A

one step down one explains the similar between one element and another. –> this can be used with electronegativity as an explanation

38
Q

Define effective nuclear charge

A

the net nuclear charge after considering the lessening of nuclear charge from the closed inner shells

39
Q

Cations are smaller than their parent atoms

A
40
Q

Anions are larger than their parent atoms

A
41
Q

Describe the conductivity of non-metals

A

they are negligible

42
Q

How does an element conduct electricity?

A

must contain electrons that are free to move, otherwise known as mobile charge carriers

43
Q

Why are metals good conductors?

A

they have metallic bonding which means positive metals ions are attracted to delocalised and mobile electrons
- This means they are free to move and carry charge

44
Q

Why are metalloids not good conductors?

A

they exist in giant lattice structures
- which means few electrons have enough energy at room temperature to enter high energy levels
- this means there are only a few delocalised electrons
- but at higher temperatures electrons are promoted to the higher energy shells
- therefore there are more delocalised electrons to carry charge

45
Q

Why are non-metals not good conductors?

A

they are held in strong covalent bonds

46
Q

What are melting and boiling points depenent on?

A

strength of forces between particles
- the stronger the force the more energy needed to overcome
- therefore higher temperature

47
Q

What happens when a substance melts?

A

attractive forces are broken or loosened, such that the particles can move freely around each other but are still close

48
Q

What happens when a substance boils?

A

the remaining attractive forces are broken so the particles can move freely and far apart

49
Q

What is the trend of melting points across a period?

A

metals - high
metalloids - very high
non metals - low

50
Q

Why do metals have high boiling points?

A

they have metallic bonding where the positive metal ions are attracted to delocalised mobile electrons

51
Q

Why do metalloids have high boiling points?

A

some exist in network covalent bonds, which is a giant lattice structures in different arrangements like tetrahedral.
- they for a giant molecule or macromolecule
- they have high melting points because the covalent bonds are strong.

52
Q
A

Chemistry Module 1 (4 decks)

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