Chapter 7- Periodicity Flashcards

1
Q

how did Mendeleev order the periodic table

A

by atomic mass - he left gaps where required so that elements could be in groups of similar chemical properties

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2
Q

what happens to atomic number moving left to right

A

it increases

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3
Q

what are groups

A

elements with the same chemical properties because they have the same number of outer electrons

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4
Q

what are periods

A

horizontal rows, period number relates to the highest energy level held by electrons in that atom

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5
Q

what does periodicity include

A
  • electron configuration
  • ionisation energy
  • structure
  • melting/boiling points
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6
Q

what are the blocks of the periodic table

A

+ s-block = groups 1 and 2 –> outermost electrons in the s-subshell
+ d-block = transition metals - outer electrons in d subshell
+ p-block = right side = outermost electrons in the P subshell

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7
Q

what is the definition of ionisation energy

A

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

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8
Q

what are the 3 factors affecting ionisation energy

A

nuclear charge, atomic radius, shielding

MENTION ALL 3 IN ANSWERS

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9
Q

explanation of how nuclear charge affects ionisation energy

A

the greater the atomic number/nuclear charge, the more protons there are; this exerts a greater pull on the outer electrons / more electrostatic forces, so more energy is required to remove the outer electron

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10
Q

explanation of how atomic radius affects ionisation energy

A

a greater atomic radius leaves the outer electrons further from the nucleus, across a period atomic radius decreases due to increased nuclear charge, a greater distance gives less electrostatic force so less ionisation energy

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11
Q

explanation of how shielding affects ionisation energy

A

greater shielding reduces the attraction from the nucleus to the outer electrons, this decreases ionisation energy. so more electron shells gives a lower ionisation energy

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12
Q

what happens to ionisation energy across a period

A
  • nuclear charge increases across a period so electrostatic attraction increases
  • atomic radius decreases across a period due to greater nuclear charge. this means the outer electrons are closer to the nucleus so attraction increases
  • shielding remains constant
  • overall ionisation energy increases
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13
Q

what are the two exceptions to the general ionisation energy rule across a period

A
  • where a higher energy subshell is being filled/electron being removed from; this is further away so ionisation energy decreases
  • paired electrons in P subshell are easier to remove due to the repulsion between electrons (only for first paired p electron)
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14
Q

explanation of ionisation energy down a group

A
  • shielding increases as you move down a group; this decreases attraction to outer electrons
  • atomic radius increases down a group; this decreases electrostatic attraction to outer electrons
  • nuclear charge increases down a group but this is not sufficient to overcome the changes due to the other factors
  • overall ionisation energy decreases
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15
Q

how does energy change with successive ionisations

A
  • with each successive ionisation, more energy is required as you are then removing an electron from a positive ion
  • this increases the effective nuclear charge
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16
Q

what does a ‘big jump’ indicate in successive ionisation energies

A

that a new shell has been entered

17
Q

what is metallic bonding

A

metallic bonding occurs when metal atoms ‘donate’ their outer electrons to form positive ions in a sea of delocalised electrons.
the cations are fixed, the electrons are not

18
Q

what are the properties of metallic substances

A

most are:

  • strong due to large electrostatic forces
  • high electrical conductivity- the electrons are free to move a charge
  • high mpt and bpt; ‘a large amount of energy is required to overcome the strong electrostatic forces in metallic bonding’
  • insoluble
19
Q

what is one factor of a metal’s bpt

A
  • the ion that it forms, 1+, 2+ etc.
20
Q

giant covalent structures examples

A

boron, carbon, silicon

21
Q

what are some properties of giant covalent structures

A
  • very high mpt and bpt because a large amount of energy is required to break covalent bonds
  • almost completely insoluble
  • don’t conduct electricity (other than graphite and graphene)
22
Q

why do group 2 metals have a higher bpt and mpt than group 1 metals

A

they donate 2 electrons not 1 so there are greater electrostatic forces

23
Q

mpt and bpt trends across periods 2 and 3

A
  • increases G1 to G2 and aluminium is higher again
  • bpts max at carbon and silicon because covalently bonded
  • tends to decrease after this because only intermolecular forces need to be broken
24
Q

what is an exception of bpts decreasing after carbon and silcon

A

sulphur has a higher bpt than phosphorus because it is S8 but only P4 so there are greater London forces

25
Q

Why is the hydrogen bonding stronger in water than in ammonia (2)

A
  • water has two lone pairs so it can form twice as many hydrogen bonds
  • O-H is a greater dipole than N-H so it has a greater electrostatic attraction
26
Q

if there is a large jump between the Xth and (X+1)th ionisation energies, which position in the period does the element sit

A

the Xth position

- if there is a large jump at 5–> 6 for example then the element must have 5 in its outer shell

27
Q

ideally how should you write the electron configurations

A

in order of numbers so 3d before 4s even though 4s fills first, doesn’t really matter though

28
Q

which things should always be mentioned when comparing the atomic radius of elements in the same period

A
  • nuclear charge and therefore electrostatic attraction

- SAME NUMBER OF SHELLS