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Chemistry🧪🔫 > 7) PERIODICITY > Flashcards

7) PERIODICITY Flashcards

1
Q

Who arranged the periodic table in the past?

A

Mendeleev

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2
Q

How is the periodic table arranged now?

A

-144 elements
-7 periods
-18 groups

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3
Q

What are the 3 factors in which the elements are arranged in the periodic table?

A

-atomic number
-groups
-periods

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4
Q

What do the positions of elements show?

A

The elements physical + chemical properties

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5
Q

How are the elements arranged according to atomic number?

A

-left —> right
-increasing atomic number

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6
Q

How are the elements arranged according to groups?

A

-elements in the same group have the same number of e- shells
-similar properties

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7
Q

what does a period show?

A
  • the number of highest energy e-
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8
Q

Define periodicity

A

Repeating trend in properties of the elements

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9
Q

What are the 4 types of periodicity?

A

-electron configuration
-ionisation energy
-structure
-melting points

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10
Q

What are the 4 types of sub shells?

A

-s
-p
-d
-f

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11
Q

What is the trend across a period in electronic configuration?

A

1 e- increase between each element

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12
Q

What is the difference in electronic configuration down a group?

A

Elements in the same group have the same number of e- in each sub shell

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13
Q

Define first ionisation energy?

A

Energy required to remove 1 e- from each atom in 1 mole of a gaseous element

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14
Q

what are the factors that affect the ionisation energy?

A

-atomic radius
-nuclear charge
-electron shielding

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15
Q

what are the effects of the atomic radius on the ionisation energy?

A

greater distance = less nuclear attraction = decreased IE

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16
Q

what are the effects nuclear charge on the nuclear charge?

A

more protons = greater attraction = increased IE

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17
Q

what are the effects of electron shielding on the nuclear charge?

A

-more shields = less attraction = lower IE

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18
Q

what does successive ionisation energy mean?

A

as the number of e- removed increases = ion becomes more positive

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19
Q

what is the trend in first IE down a group?

A

decreases

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20
Q

why does the IE decrease down a group?

A

-atomic radius increases
-shielding increases
-nuclear attraction decreases
-IE decreases

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21
Q

what is the trend in first IE across a period?

A

increase

22
Q

why does the first IE increases across a period?

A

-nuclear charge increases
-nuclear attraction increases
-atomic radius decreases
-first IE increases

23
Q

why is there a fall in first IE from Be to B ?

A

-start of the filling of the 2p sub shell
-2p sub shell has a higher energy than 2s
-2p e- easier to remove than 2s in Be

24
Q

why is there a fall in the first IE from N to O ?

A

-in both N + O the highest energy sub shell is 2p
-O has a paired e- = repel each other = easier to remove
-IE in O is lower than N

25
Q

define metallic bonding?

A

strong electrostatic attraction between cations + delocalised e-

26
Q

describe the structure of metallic bonding?

A

-cations fixed in position = maintaining structure
-sea of delocalised e- = mobile

27
Q

at which state are metals able to conduct electricity?

A

solid + liquid

28
Q

how are metals able to conduct electricity?

A

delocalised e- move = carry charge

29
Q

what type of melting + boiling point do metals have?

A

high

30
Q

why do metals have high melting + boiling points ?

A

-strength of the metallic bonds
-high temp needed = provide high energy = overcome strong electrostatic attraction

31
Q

what is the solubility of metals?

A

do not dissolve

32
Q

what does the strength of the metallic bonding depend on?

A

-size of +ve charge ions
-size of metal ion
-number of e- per atom

33
Q

how does the size of the +ve charge ion affect the strength of the metallic bond?

A

greater +ve charge = stronger attraction to delocalised e- =stronger bond

34
Q

how does the size of the metal ion affect the strength of the metallic bond?

A

-small ion = strong bond
-small radius = shorter distance between +ve nucleus + delocalised e- = stronger attraction

35
Q

how does the # of mobile e- per atom affect the strength of the metallic bond?

A

-more mobile e- = more e- holding the lattice together

36
Q

define allotrope

A

different molecular form of the same element = different physical properties

37
Q

what are the different forms of carbon?

A

-diamond
-graphite
-graphene
-buckminsterfullerene

38
Q

what type of bond is present between all types of carbon?

A

covalent bond

39
Q

why is diamond a hard substance?

A

-covalent bonds

40
Q

why does diamond have a high mpt + bpt?

A

strong covalent bonds = a lot of E to overcome

41
Q

why is diamond a good conductor of HEAT ?

A

-covalent bonds
-vibrations of thermal energy = rapidly transferred through the substance

42
Q

why is diamond insoluble?

A

attraction between carbon atoms with covalent bonds in diamond is a lot stronger than the attraction that could occur between the solvent mol.

43
Q

why is diamond a poor conductor of ELECTRICITY ?

A

-all the outer e- in C are fixed between covalent bonds
-no free e- to carry charge

44
Q

describe the structure of graphite

A

-each C atom covalently bonded to 3 other C atoms
-1 outer e- not involved in a covalent bond = delocalised + mobile

45
Q

how are the atoms in graphite arranged?

A

-sheets of hexagons
-layer structure
-joined together by weak intermolecular forces = van der waals forces

46
Q

why can graphite be used as a lubricant?

A

-layers can easily slide over each other
-due to weak intermolecular forces

47
Q

why does graphite have a high mpt + bpt?

A

strong covalent bonds = large E needed to break

48
Q

why is graphite insoluble ?

A

attraction between carbon atoms with covalent bonds in diamond is a lot stronger than the attraction that could occur between the solvent mol.

49
Q

why is graphite a good conductor of electricity?

A

delocalised e- = move between sheets = carry charge

50
Q

why does graphite have a low density?

A

layers are far apart

Chemistry🧪🔫 (8 decks)

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