Chapter 7: Periodicity Flashcards
(98 cards)
1
Q
How many elements were known when Mendeleev arranged the periodic table?
A
- Arranged in order of increasing atomic mass
2
Q
How did Mendeleev organise the groups?
A
With elements of similar properties + left gaps which did not fit assuming that some atomic mass measurements were incorrect + yet to be discovered
3
Q
What did Mendeleev predict?
A
Properties of unknown elements from group trends
4
Q
When were protons discovered?
A
1900s
5
Q
What is ekasilicon?
A
An element Mendeleevd predicted the properties of in 1871
- element discovered in 1886 and was named germanium
- Mendeleev’s predictions were not far off
6
Q
Mendeleev’s predictions for Eka-si vs Ge:
A
Atomic mass: 72 and 72.61
Density: 5.50 and 5.35
Formula: EO2 and GeO2
Oxide density: 4.70 and 4.70
7
Q
How many elements does the periodic table have?
A
As of 2014, has 114 elements arranged in 7 horizontal periods + 18 vertical groups
8
Q
Why is the periodic table important?
A
The first point of reference for chemists everywhere + most labs have it on the wall
9
Q
What does the arrangement of the periodic table show?
A
Trends among elements of
- position of elements are lined to chemical + physical properties so is used to predict properties
10
Q
How are elements arranged?
A
In order of increasing atomic number
- each successive element has atoms with an extra proton
11
Q
What are groups?
A
Vertical columns where atoms have the same number of outer shells + similar properties
12
Q
What are periods?
A
Horizontal rows
- number gives the number of the highest energy electron shell in an elements atoms
13
Q
What is periodicity?
A
A repeating trend in properties of elements
14
Q
Properties of periodicity:
A
- electron configuration
- ionisation energy
- structure
- melting points
15
Q
What is the chemistry of an elements does determined by?
A
Electron configuration, specifically outer + highest energy electron shell
16
Q
Trend across period 2:
A
- 2s subshell fills with 2 electrons followed by 2p sub shell with 6 electrons
17
Q
Trend accord period 3:
A
3s subshell fills with 2 electrons and 3p fills with 6
18
Q
Trend across period 4:
A
3d subshell is involved but the highest shell number is n = 4 where only 4s and 4p sub shells are occupied
19
Q
Trend down a group:
A
- each group have atoms with the same number of electrons in outer shell + atoms with the same number
- this similarity in electron configuration gives elements in the same group their similar chemistry
20
Q
What are the blocks?
A
- s,p,d,f
21
Q
What are the 2 ways of numbering groups?
A
- groups 1-7 then 0, based on s and p blocks
Advantage: group number matched number of electrons in highest energy electron shell - groups 1-18 numbering each column in s,d,p blocks sequentially.
Used in IUPAC in 1988
22
Q
What are group 3-12?
A
Transition elements
23
Q
Old and new name for group 5:
A
Old: 5
New: group 15
Name: pnictogens N,P,As,Sb,Bi
24
Q
Old and new name for group 6:
A
Old: 6
New: group 16
Name: chalcogens O,S,Se,Te,Po
25
Old and new names for group 7:
Old: group 7
New: group 17
Name: halogens
26
Group 0 old and new names:
Old: group 0
New: group 18
Name: noble gases
27
What is ionisation energy?
Measures now easily atom loses electrons to form positive ions
28
first ionisation energy definition:
The 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
29
Equation for first ionisation energy of Na:
Na (g) → Na+ + e-
First ionisation energy = +496
30
Factors affecting ionisation energy:
- electrons attraction to nucleus. The first electron lost will be the highest energy level + have least attraction from nucleus
31
Where is the first electron lost in 1s2,2s2,2p6,3s1?
3s subshell
32
How does atomic radius affect ionisation energy?
- atomic radius, the greater the distance between the nucleus + outer electrons, the less the attraction. The force of attraction falls sharply with increasing distance
33
How does nuclear charge affect ionisation energy?
The more protons there are in the nucleus, the greater the attraction between nucleus and outer electrons
34
How does electron shielding affect ionisation energy?
- electrons are negatively charged so inner shell electrons repel outer shell electrons, causing shielding effect.
- which reduces attraction between nucleus + outer electrons
35
Ionisation energy =
Number of electrons
- e.g He (g) → He+ (g) + e- first ionisation energy
He+ (g) → He2+ (g) + e- second ionisation energy
36
Why is the second ionisation of helium greater than the first?
In helium atom there are two protons attracting two electrons in the 1s sub shell. After the first electron is lost the single electron is pilled closer to the helium nucleus
- the nuclear attraction of the remaining electron increases and more ionisation energy will be needed to remove second electron
37
Second ionisation energy definition:
The energy required to move one electron from each atom in one mole of gaseous 1+ ions of an element to form one mole of gaseous 2+ ions
38
What do successive energies show.
Evidence for different electron energy levels in an atom
39
What does a large increase between 7th and 8th ionisation energies show?
- the either electron must be removed from a different shell closer to the nucleus with less shielding
40
Shells when there is a successive ionisation energy in 7th and 8th ionisation m
- the first shell n =1 is closer to the nucleus contains 2 electrons
- the second shell n = 2 in outer shell contains seven electrons
41
What predications can be made due to successive ionisation energies?
- the number of electrons in outer shell
- the group of the element
- the identity of the element
42
Explanation for a large increase between third and fourth ionisation energies on a table:
- the fourth electron is being removed from an inner shell so there are three electrons in the outer shell + the elements must be in group 13 (aluminium)
43
What do periodic trends in ionisation energies show?
Important evidence for the existence of shells and sub shells
44
Trend in ionisation energy across a period:
A general increase
45
Trend between first ionisation energy between a period:
A sharp decrease in first ionisation energy
46
How can ionisation energy trends be explained?
- atomic radius
- electron shielding
- nuclear charge
47
Trend in first ionisation energy down a group:
Decrease and nuclear charge and shielding increases but effect is outweighed by increased radius
48
Number of shells in He, Ne and Ar:
He: 0
Ne: 1
Ar: 2
49
Trend in ionisation energy down a group:
- atomic radius increases
- more inner shells so shielding increases
- nuclear attraction on outer electrons decreases
- first ionisation energy decreases
50
General trend in ionisation energy across period 2:
- nuclear charge increases
- same shell: similar shielding
- nuclear attraction increases
- atomic radius decreases
- first ionisation energy increases
Li is 3p+ Be is 4p+ and so on
51
Where is the drop of first ionisation energy in group 2 and 3?
General increase but falls in the same two places in each period suggesting there is a periodic cause
52
What causes a drop in first ionisation energy in group 2 and 3?
Existence of subshells, energies, how orbitals fill with electrons
53
What does first ionisation energy of group 2 graph show?
- a rise from lithium to beryllium (filling 2s subshell)
- a fall to boron followed by a rise to carbon and nitrogen (adding one electron to each p orbital)
- a fall to oxygen followed be a rise to fluorine and neon (pairing 2p electrons)
54
Comparing beryllium and boron:
- first fall in ionisation between them is the start of filling 2p subshell
- 2p in boron has a higher energy than 2s subshell in Beryllium.
- first ionisation energy of boron is less than the first ionisation energy of beryllium
55
Comparing nitrogen and oxygen:
- fall in first ionisation between them show start of electron pairing in p orbitals in 2p sub shell
- in O2 the paired electrons in one of the 2p orbitals repel one another so it is easier to remove an electron from oxygen than nitrogen
- so first ionisation energy of oxygen is less than the first ionisation energy of nitrogen
56
Changeover from metal to non metal:
Takes place on a diagonal line from the top of group 13 to bottom of group 17
57
What are elements near division of metals and non metals called?
Semi metals or metalloids
58
What happens as you go down groups of metalloids?
Trend from non metal to metal,
- th divide is clearest in group 14
59
How many metals and non metals are there?
92 metals and 22 non metals
60
State of metals at room remp:
- solids except mercury
61
Properties of some metals:
- Tungsten (W) is hard
- lead (Pb) is soft
- Al is light
- Osmium (Os) is twice as dense as lead
62
Constant property for all metals:
- can conduct electricity
- charge must be able to move within a rigid structure
63
Structure in solid metal:
Each atom has donated its negative outer shell electrons to a shared sea of electrons which are delocalised through the whole structure
- positive ions left behind consist of the nucleus and inner electron shells of metal atoms
64
Metallic bonding definition:
The strong electrostatic attraction between cations and delocalised electrons
65
Position of cations and electrons in solid metals:
- delocalised electrons are mobile and are able to move throughout the structure + only electrons move
66
What is the structure of metals?
Billions of metal atoms held together by metallic bonding in giant metallic lattice
67
Properties of metals:
- strong metallic bonds - attraction between positive ions and delocalised electrons
- high electrical conductivity
- high melting and boiling points.
68
How can the physical properties of metals be explained?
in terms of the giant structure of the lattice and metallic bonding.
69
When do metals conduct electricity?
solid and liquid states.
70
How do metals conduct electricity?
When a voltage is applied across a metal, the delocalised electrons can move through the structure, carrying charge
71
Ionic compounds electrical conductivity in a solid:
have no mobile charge carriers in the solid state.
72
Melting and boiling point of metals:
Most have high melting and boiling points.
73
What is tungsten used for?
Tungsten, W, has the highest melting point at 3422°C, which is why it is used in the filaments of halogen lamps; other metals would melt.
74
What temp does mercury melt at?
-39 °C
- Other metals that melt at low temperatures include those in Group 1 of the periodic table, which all have melting points below 200 °C.
75
What does the melting point in a metal depend on?
strength of the metallic bonds holding together the atoms in the giant metallic lattice.
76
Why do metals need high temperatures?
to provide the large amount of energy needed to overcome the strong electrostatic attraction between the cations and electrons.
77
Solubility of metals:
- do not dissolve
- interaction between polar solvents and the charges in a metallic lattice would lead to a reaction rather than dissolving e.g. sodium and water
78
What do non-metallic elements exist as?
Simple covalently bonded molecules
79
What do covalent molecules form in a solid state?
Form a simple molecular lattice structure held together by weak intermolecular forces.
- have low melting and boiling points
80
What do many covalent bonds form?
Giant covalent lattice
81
What do carbon and silicon form?
Use their four outer electrons to form four covalent bonds to other carbon or silicon atoms
- gives tetrahedral structure, bond angle 109.5 by electron pairing repulsing
82
Why are giant covalent bonds difficult to break down?
Strong covalent bonds which make stable structures that are difficult to break down
83
Solubility of giant covalent lattices:
Insoluble
- The covalent bonds holding together the atoms in the lattice are far too strong to be broken by interaction with solvents.
84
Electrical conductivity of giant covalent structures:
- non-conductors of electricity.
- except graphene and graphite, which are forms of carbon.
85
Electric conductivity of carbon and silicon:
• In carbon (diamond) and silicon, all four outer-shell electrons are involved in covalent bonding, so none are available for conducting electricity.
86
How can graphene and graphite conduct electricity?
Carbon can form several structures where one of the electrons is available for conductivity.
87
Dot and cross for graphene and graphite:
- only 3 of 4 outer shell electrons are used in covalent bonding
- remaining electron is released into sea of electrons shared by all atoms
- so carbon structures contains planar hexagonal layers are good electrical conductors
88
Graphene and graphite structures:
giant covalent structures of carbon based on planar hexagonal layers with bond angles of 120° by electron-pair repulsion.
89
What is graphene?
A single layer of graphite composed of hexagonally arranged carbon atoms linked by strong covalent bonds
90
Electrical conductivity of graphene:
- same as copper
- it is the thinnest and strongest material made
91
When was graphene discovered?
Graphene was discovered in 2004 by Andre Geim and Konstantin Novoselov from the University of Manchester.
They were awarded a Nobel prize in 2010. Geim famously made graphene by using sticky tape to pull single layers of carbon atoms from the surface of graphite.
92
What is graphite composed of?
Parallel layers of hexagonally arranged carbon atoms, like a stack of graphene layers
93
Strength of graphite:
Layers are bonded by weak London forces
94
How many carbons are used in graphite?
- bonding in hexagonal layers uses 3 of 4 outer shell electrons. The spare electron is delocalised between layers so electricity can be conducted as in metals
95
Melting points across period 2 and 3 on graph:
- melting point increases from Group 1 to Group 14 (4)
- sharp decrease in melting point between Group 14 (4) and Group 15 (5)
- the melting points are comparatively low from Group 15 (5) to Group 18 (0).
96
Why is there a sharp decrease in melting points in groups 2 and 3?
- change from giant to simple molecular structures
97
Melting points in period 2 and 3:
On melting, giant structures have strong forces to overcome so have high melting points. Simple molecular structures have weak forces to overcome
98
Repeated trends in group 2 and 3:
Trend in melting points across period 2 is repeated accords period 3 and continues across s and p block of period 4
