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Georges chem > periodicity chp 7 > Flashcards

periodicity chp 7 Flashcards

1
Q

How many elements are on the peridoic table

A

118

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2
Q

what is the atomic number

A

The number of protons in an atom of an element

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3
Q

What is the atomic mass

A

The number of protons and neutrons in the nucleus of an atom of an element

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4
Q

How are the elements in the peridoic table sorted

A

reading from right to left, the elements are arranged in order of increasing atomic number (each successive element has atoms with one extra proton).

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5
Q

Why are elements in groups put into those groups

A

Each element in a group has atoms with the same number of outer-shell electrons and similar properties, as you go down the group the number of shells increase by 1.

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6
Q

why are the elements in periods put into those periods

A

The number of the period gives the number of the highest energy electron shell in an elements atoms.

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7
Q

In what 2 ways are the elements in the peridoic table group

A

periods- horizontal rows
groups- vertical rows

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8
Q

what is periodicity

A

the repeating trend in properties of the elements across the period

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9
Q

what do the blocks on the periodic table represent

A

the blocks represent the highest energy sub-shell that the elements in the block contain.

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10
Q

How many blocks are there and what are they called

A

there are 4 blocks S, P, D & F

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11
Q

what does ionisation energy measure

A

ionisation energy measures how easily an atom loses electrons to from positive ions

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12
Q

what is the first ionisation energy

A

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

the 2nd ionisation energy is the same but going from 1+ to 2+ not 0 to 2+

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13
Q

why is the first ionisation energy of an element always the lowest value.

A
  • 1st electron lost will be in the highest energy level (furthest away) and will experience the least electrostatic attraction from the nucleus and so will require the least energy to remove
  • it will also experience the most sheilding
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14
Q

what are the 3 factors that affect the ionisation energy

A

atomic radius, nuclear charge, electron shielding.

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14
Q

what is the trend in electron configuration across a period

A
  • each period has electron in new highest energy shell
  • period 2 = 2s, 2p
  • period 3 = 3s, 3p, 3d
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14
Q

How does the atomic radius effect the ionisation energy

A

The greater the distance between the nucleus and the outer electrons, the less the nuclear attraction. The force of attraction falls off sharply with increasing distance, so a greater atomic radius would mean a smaller ionisation energy value

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14
Q

what does ionisation energy measure

A

ionisation energy measures how easily an atom loses electrons to from positive ions

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14
Q

How does the nuclear charge effect the 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 and so the greatert the nuclear charge the greater the ionisation energy.

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14
Q

what are the 3 factors that affect the ionisation energy

A

atomic radius, nuclear charge, electron shielding.

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14
Q

what is the trend in electron configuration across a period

A
  • period starts with electron in new highest energy shell
  • electrons added to valence shell (from left to right)

valence = highest energy shell

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14
Q

How many blocks are there and what are they called

A

there are 4 blocks S, P, D & F

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14
Q

what does ionisation energy measure

A

ionisation energy measures how easily an atom loses electrons to from positive ions

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15
Q

why is the first ionisation energy of an element always the lowest value.

A

the 1st electron lost will be in the highest energy level and will experience the least attraction from the nucleus and so will require the least energy to remove.

15
Q

How many blocks are there and what are they called

A

there are 4 blocks S, P, D & F

15
Q

How does the atomic radius effect the ionisation energy

A
  • greater the distance between the nucleus and the outer electrons, the less the nuclear attraction.
  • force of attraction falls off sharply with increasing distance
  • so a greater atomic radius would mean a smaller ionisation energy value
15
Q

whats the trend in 1st ionisation energies across the 1st 20 elements

A

-There is a general increase in the 1st ionisation energy across each period (H->He, Li->Ne, Na->Ar).
-There is a sharp decrease in 1st ionisation energy between the end of one period and the start of the next period
(He->Li, Ne->Na, Ar->K).

16
Q

what is the trend in the 1st ionisation energy down a group

A

1st ionisation energies decrease down a group.

17
Q

Explain the trend in decreasing 1st ionisation energies down a group

A
18
Q

explain the trend is 1st ionisation energies increasing across a period

A
19
Q

what is unique about the ionisation energies across period 2 and 3

A

there are 2 drops in ionisation energies in each period at the same point

20
Q

explain the dips in ionisation energies in period 2 and 3

A

the fall in 1st ionisation energy from beryllium to boron marks the start of filling the 2p sub-shell.

-The 2p sub-shell in boron has a higher energy than the 2s sub-shell in beryllium. Therefore, in boron the 2p electron is easier to remove than one of the 2s electrons in beryllium. The 1st ionisation energy of boron is less than the 1st ionisation energy of beryllium.

21
Q

explain the dropping in ionisation energies between nitrogen and oxygen

A

the fall in 1st ionisation energy from nitrogen to oxygen marks the start of electron pairing in the p-orbital of the 2p sub-shell
-in nitrogen and oxygen the highest energy electron are in a 2p sub-shell
-in oxygen, the paired electrons in one of the 2p orbitals repel one another, making it easier to remove an electron from an oxygen atom than a nitrogen atom
-therefore the 1st ionisation energy of oxygen is less than the 1st ionisation energy of nitrogen

22
Q

what is the one propertie constant across all metals

A

all metals conduct electricity

23
Q

explain the atomic structure of metals

A

-each atom has donated its negative outer-shell electrons to a shared pool of electrons, which are delocalised throughout the structure.
-the cations left behind consist of the nucleus and the inner electron shells of the metal atoms.
-the cations are fixed in position, maintaining the structure and shape of the metal
-the delocalised electrons are mobile and are able to move throughout the structure, only the electrons move.

24
Q

what is metallic bonding

A

metallic bonding is the strong electrostatic attraction between cations and delocalised electrons.

25
Q

what is the structure that results from metallic bonding

A

giant metallic lattice

26
Q

what are the most common properites of metals

A

-strong metallic bonds - attraction between positive ions and delocalised electrons
-high electrical conductivity
-high melting and boiling point

27
Q

why are metals electrically conductive

A

metals conduct electricity in solid and liquid states. When a voltage is applied across a metal, the delocalised electrons can move through the structure, carrying charge.

28
Q

why do most metals have high melting and boiling points

A

the melting point depends on the strength of the metallic bonds holding together the atoms in the giant metallic lattice.
-for most metals, high temperatures are necessary to provide the large amount of energy needed to overcome the strong electrostatic attraction between the cations and electrons.
-this strong attraction results in most metals having high melting and boiling points.

29
Q

do metals dissolve

A

metals do not dissolve and instead react

30
Q

what atoms can from giant covalent lattices

A

boron, carbon and sillicon all form giant covalent lattices

31
Q

How does nuclear charge effect the 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 and so more energy is needed to overcome this greater attraction.

31
Q

How does electron sheilding effect ionisation energy

A
  • reduces ionisation energy
  • electrons on lower energy levels repel electrons on higher energy levels
31
Q

what is the definition of second ionisation energy

A

The second ionisation energy is the energy required to remove 1 electron from each ion in 1 mole of gaseous 1+ ions of an element to form 1 mole of gaseous 2+ ions.

the 3rd ionisation energy and so on all follow the same definition

32
Q

what does successive ionisation energy provide evidence for

A

successive ionisation energies provide evidence for the different electron energy levels in an atom

32
Q

How does successive ionisation energies provide evidence for shells

A

the large increase between the 7th and 8th ionisation energies suggests that the 8th electron must be removed from a different shell, closer to the nucleus and with less shielding

32
Q

what predictions could be made using 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

33
Q

explain the dropping in ionisation energies between S block and P block

A
  • lower because boron has a single electron unpaired electron in its 2p sub-shell
    ^as subshell in incomplete and at higher energy level easier to remove
33
Q

whats the difference between simple molecular lattices and giant covalent lattices

A
  • simple molecular lattice held together by weak intermolecular forces, relatively small, low melting/boiling point
  • giant covalent lattices made of billions of atoms held together by a network of strong covalent bonds, high meliting/boiling point

both covalent bonds

34
Q

what are the properties of giant covalent lattices

A

high melting and boiling point
insoluble in almost all solvents
not electrically conductive (depends on structure)

35
Q

explain the high melting and boiling point of giant covalent lattices

A

covalent bonds are strong, high temperatures are necessary to provide the large quantity of energy needed to break the string covalent bonds

36
Q

explain the insolubility of giant covalent lattices

A

the covalent bonds holding together the atoms in the lattice are far to strong to be broken by interactions with solvents

37
Q

what are the exceptions to the non-conductive giant covalent lattices

A
  • graphene and graphite are both conductive
    in both graphene and graphite only 3 out of the 4 outer shell electrons are used in bonding, the remaining electron is delocalised and so is mobile and able to carry charge

graphene is just 1 layer of graphite

38
Q

why are most giant covalent lattice not electrically conductive

A

in carbon(diamond) and silicon, all 4 outer shell electrons are involved in covalent bonding, so none are available for conducting electricity

39
Q

whats the trend in melting point in period 2 and 3

A

across period 2 and 3
-the melting point increases from group 1 to group 14
-there is a sharp decrease in melting point between group 14 and group 15
-the melting points are comparatively low from group 15 to group 18

the trend in melting points across period 2 is repeated across period 3, and continues across the s- and p- blocks from period 4 downwards.

40
Q

why is there a sharp decrease in melting point from group 14 to group 15

A

the sharp decrease in melting point marks a change from giant to simple molecular structure.

41
Q

why are most giant covalent lattice not electrically conductive

A
  • in carbon (diamond) and silicon all 4 electrons are involved in covalent bonding
  • ^no mobile charge carriers

Georges chem (29 decks)

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