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A-Level Chemistry > Chapter 7 IMPORANT > Flashcards

Chapter 7 IMPORANT Flashcards

1
Q

Mendeleev periodic table then

A

Arranged in terms of atomic mass (didnt know about suabtomic particles then)
Lined up elements with similar properties
Swapped elements and lef gaps for undiscovered
Predicted properties of missing elements from group trends

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2
Q

Periodic table now

A

Arranged in periods and groups - first point of reference for chemists everywhere
Arranged in order of increasing atomic number
Groups - atom eith the same number of outer shell electrons and similar properties
Elemenst arranged in periods - number of highest energy electron shell in an elements atom

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3
Q

Periodicity?

A

Repeating trend in properties across each period - trend from metals to non-metals

1) Electron configuration
2) Ionisation energy
3) Sturcture
4) Melting points

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4
Q

Trend across a period - electron config

A

Each period starts with an electeon in highest energy subshell
Across period 2 - 2s sub shell fills with 2 electrons first followed by 2p with 6
Same across period 3 but with 3s and 3p sub-shells
Across period 4 - although 3d sub shell involved, only 4s and 4p sub shells are occupied (n=4 is highest shell number)

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5
Q

Trend down a group electron config

A

Atoms of elements in same group have same number of electrons in each outer sub-shell - this is what gives them their similar chemistry

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6
Q

Blocks?

A

Elements can be divided corresponding to their higest energy sub-shell to give 4 distinct blocks s, p, d and f

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7
Q

S block

A

First 2 groups (include helium)

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8
Q

D block

A

Transition metals (10 groups)

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9
Q

P block

A

6 groups non-metal

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10
Q

F block

A

14 groups below at bottom - DISREGARD LANTHANIDES ETC THEY ARE PART OF D BLOCK AS WITH TRANSITION METALS

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11
Q

Two ways of numbering groups

A

Old numbers - groups 1-7 and then 0 ; based pn s and p blocks - the advantage of the pld numbering is that the group number matches the number of electrons in outer shell
IUPAC - 1-18 groups - s, d and p blocks sequentially ; approved in 1988 but it takes time for old practices to change

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12
Q

Group 1

A

Alkali metals

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13
Q

Group 2

A

Alkaline earth metals

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14
Q

Groups 3-12

A

Transition elements

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15
Q

Group 7

A

Halogens

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16
Q

Group 0

A

Noble gases

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17
Q

3 factors affecting ionisation energy

A

Atomic radius
Nuclear charge
Electron shielding

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18
Q

What is first ionisation energy?

A

Energy required to remove one electron from each atom in one mole of gaseous atoms of an element to form one mole of gaseous 1+ ions

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19
Q

Exacmple of Na first and second ionisation energy

A

Na (g) -> Na+ (g) + e-

Na+ (g) -> Na2+ (g) + e-

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20
Q

Change from metal to non-metal?

A

Staircase from top of group 3 down to the bottom of group 7 ; elements newr rhe divide (boron, silicon, germanium, arsenic, antimony) can show in-between properties and are called metalloids

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21
Q

Change from non-metal to metal?

A

Going down the group - clearest in group 4 with carbon. 92 metals vs 22 non-metals ; nonxmetals are especially importsnt and in particular elements like carbon hydrogen oxygen and nitrogen

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22
Q

At room temperature metals?

A

All metals except mercury are solids - wide range of properties

23
Q

Constant property of all metals

A

ABILITY TO CONDUCT ELECTRICITY - CHARGE MUST BE ABLE TO MOVE WITHIN A RIGID STRUCTURE FOR CONDUCTION TO RAKE PLACE

24
Q

Metallic bonding

A

In solid metal structure, each metal arom has turned into cations by donating it’s outer electrons to a sea of delocalised electrons ; metallic bonding is strong electrostatic attraction between nucleus and these electrons. CATIONS FIXED - MAINTAINING SHAPE
ELECTRONS CAN MOVE - MOBILE CARRIERES OF CHARGE

25
Q

How are they held together?

A

In a giant metallic lattice

26
Q

Properties of most metals (giant structure + bonding)

A

Strong metallic bonds
High electrical conductivity
High melting and boiling points

27
Q

How do metals conduct electricity?

A

In solid and liquid states - potential difference is applied then delocalised electrons can move through the structure carrying a charge (towards positive terminal) ; contrast eith ionic compounds which have no mobile charge carriers in solid state

28
Q

Highest melting point

A

Tungsten - used in filaments of halogen lamps

29
Q

Lowest melting point

A

Mercury + group 1 elements in peridoic table

30
Q

What does melting point depend on?

A

Strength of the metallic bonds holding together the lattice - most elements, a lot of energy is needed to overcome these string electrostatic forces of attraction therefore high melting point

31
Q

Solubility of metals

A

They do not dissolve - instead interactions lead to reactions takinf place

32
Q

Simple molecules?

A

Many non-metals exist as simple covalently bonded molecules ; a simple molecular lattice structure held together by weak intermolecular forces therefore low boiling point

33
Q

Carbon and solicon and boron?

A

Billions of atoms held together by a network of strong covalent bonds to form a giant covalent lattice

34
Q

What is so special about carbon and silicon?

A

4 electrons in outer shell (group 4) - use this to form 4 covalent bonds to other carbon/silicon atoms resulting in a large complex tetrahedral structure with 109.5 degrees electron psir repulsion

35
Q

Properties of giant covalent structures?

A

Insoluble - covalent bonds are too strong to be broken down by any interactions with solvents
High melting and boiling points - high temeprature necesary to beeak the many covalent bonds
Non-conductors of electricity (mostly) - ONLY EXCEPTION IS GRAPHENE AND GRAPHITE ; in carbon and silicon all electrons are used in bonding therefore none available to carry any charge

36
Q

Graphene and graphite

A

Form of carbon where only 3 take part in bonding with each atom releasing 1 into a pool of delocalised elecyrons - bond angles of 120 degrees trigonal planar. Hexagonal layers that slide lver each other

37
Q

Periodic trends in melting points?

A

Across Period 2 and 3 - melting point increases from 1 to 4 (culminating with giant covalent lattices) then there is a sharp decrease in melting point between group 14 and group 15 with the melting points even lower from 15 to 18

38
Q

Where is this trend repeated?

A

Across peripd 4 in s and p block too - sharp decrease in melting point makes a change from giang structures to simple structures ; strong forces to weak forces

39
Q

Atomic radius

A

Greater distance between nucleus and outer electrons = less attraction

40
Q

Nuclear charge

A

More protons = grewter attraction between psootvie nucleus and outer electrons

41
Q

Electron shielding

A

Inner shell electrons repel outer shell ; this repulsion reduces attraction between the nucleus and outer electrons

42
Q

How many ionisation energies?

A

An element has as many ionisation energies as electrons

43
Q

Why is second ionisation energy higher?

A

Fewer electrons pulled towards the same number of protons therefore larger attraction therefore more ionisation energy will be needed to remove the electron

44
Q

Second ionisation energy

A

Energy required to remove one electron from each ion in one mole of gaseous 1+ ions of an element to form one mole of gaseous 2+ ions

45
Q

Number of the ionisation energy?

A

Same as the charge on the ion produced ; second ionisation energy PRODUCES 2+ ion from 1+ ion

46
Q

What do successive ionisation energies allow us to deduct?

A

Number of electrons in the outer shell and the group of the element in the peridoic table and thus the identity of an element

47
Q

Trend in first ionisation energy down a group?

A

Decreases as nuclesr charge increases but this is outweighed by atomic radius and shielding increasing therefore less attraction thus essier to remove outer electron

48
Q

Trend in ionisation energy across a period

A

General increase in first ionisation energy - nuclear charge increased with similar shielding snd a decrease in atomic readius therefore nuclear attraction increases and thus first ionisation energy increases

49
Q

Sub-shell trends in ionisation energy ; periodicity

A

Across both period 2 and period 3 the first ionisation energy does fall in two places ; drops occuring in the same place across each period suggesting that there might be a periodic cause - LINKED WITH EXISTENCE OF SUB-SHELLS, ENERGIES AND HOW ORBITALS FILL

50
Q

Where are two falls across period 2

A

1) Beryllium to boron

2) Nitrogen to oxygen

51
Q

Comparing beryllium and boron

A

Fall in first ionisation energy marks the start of filling the 2p sub-shell. 2p sub+shell has a higher energy level than 2s sub-shell therefore in boron the 2p electron is easier to remove that one of the 2s electrons in beryllium.

52
Q

Comparing nitrogen and oxygen

A

Fall is because in oxygen, the elctrons beginnpairing in the 2p sub-shell ; they repel one another making ti easier to remove an electron from an oxygen than a nitrogen atom. Thus first ionisation energy of oxygen is less than the first ionisation energy of nitrogen

53
Q

Electrons in 2p nitrogen

A

Equally spaced apart and spins are at right angles with equal repulsions

54
Q

Electrons in 2p in oxygen

A

2p electrons start to pair together so the paired electrons repel

A-Level Chemistry (28 decks)

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