Unit 1 revision For 3.3 Flashcards
1
Q
What are in the p-block?
A
Non-metals
2
Q
What are in the s-block?
A
Metals
3
Q
What does something being in the “p block” actually mean?
A
Outer electron is in the p orbital
4
Q
What does something being in the “s block” actually mean?
A
Outer electron is in the s orbital
5
Q
Where are the metals and where are the non-metals on the periodic table?
A
Left of zig-zag = metals
6
Q
What are the elements on the zig zag line on the periodic table?
A
Semi-metals (Metaloids)
7
Q
What can Metaloids do?
A
React as both metals and non-metals, depending on conditions and what they react with
8
Q
Trend in atomic radii across the groups of the periodic table + explanation
A
Decreases
Increased nuclear charge = pulls electron shells towards itself
9
Q
Trend in atomic radii down the periods of the periodic table + explanation
A
Increases
More electron shells
10
Q
Electronegativity
A
A measure of the ability of an element to attract a pair of electrons to itself in a covalent bond
11
Q
Electronegativity trend across the groups of the period table + explanation
A
Increases
Increased nuclear charge
12
Q
Electronegativity trend down the periods of the periodic table + explanation
A
Decreases
Increased distance from the nucleus
Increased shielding
(Outweighs the increased nuclear charge)
13
Q
Electronegativity values of metals
A
Low
14
Q
Electronegativity values of non-metals
A
High
15
Q
Describe the energies of electrons
A
Fixed
16
Q
What does electronic configuration determine?
A
The chemical properties of an element
17
Q
What’s the name for the numbers of electorn shells?
A
Principal quantum numbers
18
Q
Which electron shell has the lowest energy?
A
The on closest to the nucleus
19
Q
Orbital
A
A region in the space of a loud of negative charge where you’d likely find an electron
20
Q
What can an orbital hold?
A
Up to 2 electrons with opposite spin
21
Q
Why do electrons have opposite spin?
A
In order to stop them from totally repelling each other
22
Q
What’s different about different orbitals?
A
Different energies
23
Q
S orbital shape
A
Spherical
24
Q
P orbital shape
A
Dumbbell
25
Number of orbitals and electrons of sub-shell s
1, 2
26
Number of orbitals and electrons of sub-shell p
3, 6
27
Number of orbitals and electrons of sub-shell d
5, 10
28
Maximum number of electrons in a shell
2n^2, where n is the principal quantum number
29
Number of sub shells in first shell
One
1s
30
Number of subshells in second shell
2s, 2p
31
Number of subshells in third shell
3s, 3p, 3d
32
What do we put as the core with electronic configurations?
Noble gases
33
Which subshell is filled first - 4s or 3d and why?
4s first as it’s of lower energy
34
What does electrons having opposite spin make them?
More stable
35
Rules for electrons filling shells
1.) electrons are placed in the shell with the lowest energy
2.) only 2 electrons can occupy an orbital and they have opposite spins
3.) the electrons are placed in a sub-shell with parallel spins until the sub-shell is half full
36
What do we do to show the electrons in their boxes once the p orbital is reaches?
Half fill the box’s with downs before filling once full
37
Exceptions of the usual electronic configuration writing + explanation
Chromium ——> needs a half-filled 4s box before filling the 3d box to e stable
Copper ——> needs a half filled 4s box before the filling the 3d box to be stable
38
Why are chromium and copper exceptions for writing electronic configurations?
Because the 3d orbital becomes lower in energy than the 4s orbital when filled, which break the rules of filling electron shells
39
How do we write the electronic configuration of ions?
Write it as normal
+ charge —> remove an electron from he subshell with the highest energy (the last in the list)
40
d-subshells of transition metals
Partially filled
41
What do successive ionisation energies provide information about?
The arrange of electrons around the molecules
42
1st ionisation energy equation for sodium
Na (g) —> Na+ (g) + e-
43
What’s important to remember when writing out ionisation energy equations?
1.) write g - we can only ionise gases
2.) show state symbols
3.) show electrons on the RHS
44
Do successive ionisation energies always increase or decrease?
Increase
45
Why do successive ionisation energies always increase?
Distance form the nucleus decreases, so the nuclear attraction increases
Each electron removed = less electron-electron repulsion = the shells are drawn closer to the nucleus
Greater effective nuclear charge (same number of protons holding fewer electrons)
46
Ionisation energy
The energy required to remove one or more electrons from an atom
47
What all have their own ionisation energy?
Every electron in an atom
48
Molar first ionisation energy
The energy required to remove one mole of electrons from one mole of is gaseous atoms to form one mole of gaseous ions
49
Standard conditions
298k
1 atmosphere pressure
50
Whats the name for ionisation energy if it occurs under standard conditions?
Standard molar first ionisation energy
51
Equation for the ionisation of hydrogen
H (g) —> H+ (g) + e-
52
Is ionisation endo or exothermic? Why?
Endothermic (requires energy to remove electrons from an atom)
53
Unit for energy changes
kJmol-1
54
Why do outer electrons not feel the full force of the positive charge of the nucleus?
Inner electrons partially screen them from the effects
55
Why are there considerable variations in ionisation energies?
The inner electron shells screen the outer shells form the full effects of the positive nucleus
56
How do we determine a group from successive ionisation energies?
- look for a significant jump in ionisation energy = from an inner principal quantum shell
- whichever one it’s jumped *from* is the group
57
Valence electrons
Electrons in the outer shell
58
How do we identify a specific element using successive ionisation energies?
The electrons before the jump are the valence electrons. Find the element with that valence.
59
What do we fill and empty first with transition metals when working out successive ionisation energies?
4s orbital
60
What does having the highest last ionisation energy mean for an element?
Greatest nuclear charge
61
Why do we have to put (g) in ionisation equations?
Can only ionise gases
62
What factors does the value of a 1st ionisation energy depend upon?
1.) nuclear charge (i.e - number of protons)
2.) outer electron distance from the nucleus
3.) screening produced by filled inner energy levels
63
What is screening caused by?
Full shells
64
What happens to ionisation energy values across the periods? Why?
Increase in ionisation energy as…
Nuclear charge increases
Electrons are added to the same shell = no effect on screening
More energy required to remove them from the atom as the outer electrons are drawn closer to the nucleus
65
On an ionisation energy graph, what are the peaks?
Group 0 (noble gases)
66
On an ionisation energy graph, what are the troughs?
Group 1
67
2 exceptions to increasing ionisation energies across the period + explanation
Be to B and Mg to Al
Boron does have an extra proton, but there’s additional screening by a full 2s level and less energy needed to remove the 2p electron
=dip on graph + decrease in the first ionisation energy
N to O and P to S
With oxygen, the new electron goes into an orbital which already has its own electron, causing increased repulsion between 2 electrons of the same orbital, therefore its easier to remove the electron
(This outweighs the effect of the extra proton)
68
Change in ionisation energy from hydrogen to helium + explanation
Hydrogen has a high ionisation energy, as it consists of a single electron close to the nucleus, which it’s strongly attached to with no screening. With helium, the electron being removed is once again close to the nucleus and is unscreened, but the value of the ionisation energy is now much higher than hydrogen as the nucleus now has 2 protons attracting the electrons instead of 1
69
Ionisation energy when going from helium to lithium + explanation
For lithium, the outer electron is in the second energy level - the additional proton in the nucleus is outweighed by the screening given by the 1s^2 electrons = ionisation energy decrease
70
What happens to ionisation energy down a group and why?
Decreases, as shielding outweighs th effects of nuclear charge
71
What outweighs what - shielding or nuclear charge when it comes to ionisation energies?
Shielding
72
What might you have to use when working with ionisation energies and why?
“Log”, as they can get high
73
Successive ionisation energy
To successively remove all of the electrons until all have been removed
74
How many successive ionisation energies does an element have?
As many as it has electrons
75
What do we need to remember when writing successive ionisation energy equations?
Whatever labelled as = target
76
Why do the first row of d-block elements have similar 1st ionisation energies?
First electron to be lost from all the elements is from the 4s orbital
As you add an extra proton to the nucleus, you add an extra electron to the 3d orbital,screening the effect of the extra proton
77
Why is zinc’s first ionisation energy significantly higher than coppers?
In both, the electron is coming from the 4s level with a complete 3d^10 level inside = identical screening
Increase = attraction of the extra proton in the nucleus
78
Why is it uncommon that components containing Ba3+ ions can exist?
Ba3+ is in group 2 and has 2 outer electrons
Too much energy is needed to remove the 3rd electron
This necessitates removing an electron from the shell closer to the nucleus to form the Ba3+ ion
79
Least repulsive to most repulsive bond pairs
Bond pair-bond pair
Lone pair-bond pair
Lone pair-lone pair
80
When is the effect of lone pairs on a molecule shape considered?
After the arrangement of the electron pairs has been worked out
81
2 electron pairs shape and bond angle
Linear
180
82
3 electron pairs shape and bond angle
Trigonal planar
120
83
4 electron pairs shape and bond angle
Tetrahedral
109.5
84
5 electron pairs shape and bond angle
Trigonal bipyramidal
90/120
85
6 electron pairs shape and bond angle
Octahedral
90
86
What are lone pairs only for?
The central atom
87
Why are bond angles what they are?
They’re as far apart as the bond pars can get
88
What does the valence shell electron pair repulsion theory (VSEPR) state?
The shape adopted by a simple molecule or ion is that which keeps repulsive forces to a minimum
89
What does the VSEPR theory help us determine?
The shapes of molecules or ions
Determine the number of electron pairs in the outer shell (valence shell) of a central atom
90
What do closer bonds result in?
Greater repulsive forces
91
What do covalent bonds consist of?
A pair of electrons
92
What do bond do to other bonds and why?
Repel each other, as bonds consist of electrons and electrons are negatively charged
93
What do bonds do to each other and why?
Push each other as far apart as possible to reduce repulsive forces
Equal repulsion’s = equally spaced bonds
94
Why do most simple molecules have standard shapes and equal bond angles?
Because of the equal repulsion forces between bond pairs
95
Simple molecules
Ones with a central atom and others bonded to it
96
When do lone pair effect shapes of molecules and angles between bonds?
When on the central atom
97
What do lone pairs on the central atom of a molecule effect?
Shapes
Angles between bonds
98
Bond pairs
These electrons are spread out so that they spend time around both atoms in the bond
99
Lone pairs
They’re attached to 1 atom only and are not involved in bonding
100
Are lone pairs involved in bonding?
No
101
How are lone pairs different to bond pairs?
Occupy a smaller volume of space
Greater power of repulsion
102
Main principles of the VSEPR theory
Electrons in the outer shell around the central atom are found in pairs
Electron pairs repel one another as far away as possible until they form the most stable spatial arrangement
103
Why is a shape in a certain way with the VSEPR theory?
To keep repulsive forces to a minimum
104
How do we work out the shapes of molecules and ions?
1.) number of electrons in the outer shell of the central atom
2.) number of electrons provided by the other atoms (1 per atom)
3.) (for ions) what is the contribution of charge?
Add 1 electron for every negative charge
Subtract 1 electron for every positive charge
4.) number of electron pairs (divide total by 2)
5.) number of bond pairs + lone pairs —> effect on shape
(Formula shows bond pairs, lone pairs are whatever is left over from the bond pairs)
105
Water shape + explanation
Bent
Based on a tetrahedron but has 2 lone pairs
106
How do we work out the shape of molecules with double bonds?
Shape is calculated in the same way
Double bond repels other bonds as if it was single = the same shape
107
What can electronegativity be related to for the elements and why?
The redox behaviour of the elements, as it’s a measure of how strongly an element attracts electrons to itself
108
Which elements on the periodic table are the most powerful oxidising agents and why?
Top right
Strong attraction for electrons (high electronegativity values)
109
Which elements on the periodic table are the strongest reducing agents and why?
Bottom left
Little tendency to attract electrons (low electronegativity values)
