7.1 - 7.3 Flashcards
(51 cards)
1
Q
Back in the early 1800s how were elements ordered in the periodic table
A
- according to atomic mass
2
Q
How is the periodic table arranged now
A
- in increasing atomic number (proton number)
3
Q
Explain Döbereiner’s triads
A
- he grouped elements in 3s (triads) according to their characteristics
- he realised that element like Cl, Br and I had similar characteristic
- and he said that Br had a mass that fit half ways between Cl and I
- so he grouped them together
4
Q
Explain Newland’s octaves
A
- grouped elements in order of mass
- he noticed that every 8th element had similar properties
- he likened this pattern to octaves on a piano and so he called it the law of octaves
5
Q
How did Newland’s octave theory break down
A
- when some of the transition metals didn’t fit his pattern
6
Q
Who created the current periodic table similar to the one that we have today
A
- Dimitri Mendeleev
7
Q
How did Mendeleev order the elements
A
- by atomic mass
8
Q
Mendeleev took Newland’s work and adapted it, but what was the difference between his work and Newland’s work
A
- Mendeleev LEFT GAPS where elements didn’t fit Newlands theory
9
Q
How are the elements grouped in the periodic table
A
- in terms of similar chemical properties
10
Q
Complete the sentence:
Mendeleev was so confident in his theory that he could…
A
- predict the properties of undiscovered elements where he left the gaps
11
Q
What does all elements in the same group have
A
- the same number of electrons in the outer shell
12
Q
What does the group number relate to
A
- the number of electrons in the outer shell
13
Q
What are periods in the periodic table
A
- the rows going across
14
Q
What do elements in the same period have
A
- the same number of electron shells
15
Q
What are the s block elements
A
- H2, He
- Group 1 and 2 elements
16
Q
What are the p block elements
A
- Groups 3,4,5,6,7 and 8 elements
17
Q
What are the d block elements
A
- all the transition metals
- La and Ac in the bottom two rows
18
Q
What are the f block elements
A
- the bottom 2 rows of the periodic table excluding La and Ac
19
Q
Define First Ionisation energy
A
- the energy required to remove 1 mole of electrons from each atom in 1 mole of gaseous atoms of an element to form 1 mole of gaseous 1+ ions
20
Q
Write the equation for the First Ionisation energy of Na
A
Na (g) —> Na+ (g) + e-
21
Q
Is all ionisation endothermic or exothermic and why
A
- endothermic (positive) as it requires energy
22
Q
What are the 3 factors affecting ionisation
A
1) Shielding
2) atomic size
3) nuclear charge
23
Q
How does shielding affect ionisation
A
- the more electron shells between the nucleus and the electron that is being removed the less energy is required as there is a weaker attraction
24
Q
How does atomic size affect ionisation
A
- the bigger the atom the further away the outer electrons are from the nucleus
- there is then a weaker attraction between the nucleus and the outer electron so less energy is required to remove electrons
25
How does nuclear charge affect ionisation
- the more protons in the nucleus the bigger the attraction between the nucleus and the outer electrons
- this means more energy is required to remove the electron
26
How does first ionisation energy change as you go down a group and why (in terms of atomic size, shielding)
- it DECREASES
ATOMIC SIZE:
- because the atomic radius increase as we go down the group so the outer electron is further away from the nucleus. So the force of attraction is weaker so less energy is required to remove an electron
SHIELDING:
- shielding increases as we go down the group. There are more shells between the nucleus and the outer electron so the attractive force is weaker so less energy is required to remove the outer electron
27
What does ionisation data provide strong evidence for
- evidence for shells and proves Niels Bohr’s model of the atom is correct
28
How does first ionisation energy change as you go across a period and why (in terms of atomic size, shielding and nuclear charge)
- it INCREASES
NUCLEAR CHARGE:
- as we go across the period there is an increase in proton number, which increase the attraction between the outer electrons and the nucleus, meaning more energy is needed to remove the electron
SHIELDING:
- is similar and the distance from the nucleus marginally decrease, so shielding has no effect
29
What are the 2 exceptions when it comes to ionisation energy’s across a period
- Al
- S
30
What does the decrease in ionisation energy at aluminium provide evidence for
- atoms having sub-shells
31
Why does aluminium deviate from the general increase in ionisation across period trend
- because the outer most electron sits in a higher energy level sub- shell (3p) which is slightly further from the nucleus that the outer electron in Mg (3s)
32
What does the decrease in ionisation energy at sulfur provide evidence for
- electron repulsion in an orbital
33
Why does sulphur deviate from the general increase in ionisation across period trend
- phosphorus and sulfur both has outer electrons in the 3p orbital so the shielding is the same. However sulfur has one electron spindle pair and phosphorus doesn’t and removing an electron from sulfur involves taking it for a an orbital with 2 electrons in.
- and as these 2 electrons in the orbital repel each other less energy is needed to remove an electron from an orbital with 2 than an orbital with 1 electron like phosphorus
34
Define successive ionisation
- the removal of more than 1 electron form the same atom
35
What does the jump in the graph for successive ionisation of an element show
- the jump in energy shows us moving electrons from the shell closer the the nucleus
36
Why is there a general increase in successive ionisation
- as you are removing electrons from an increasingly more positive ion
37
How can we identify an element from a successive ionisation graph
- by looking at the total amount of electrons that were removed
38
What are two examples of giant covalent structures
1) graphite/ graphene
2) diamond
39
Describe the 6 characteristics of graphite
- each carbon is bonded 3 times and the 4th electron is delocalised
- layers slide easily as there are weak forces between the layers
- lots of strong covalent bonds so it has a HIGH MELTING POINT
- delocalised electrons between the layers allow for graphite to CONDUCT ELECTRICITY as they carry a charge
- low density as the layers are far apart
- insoluble (doesn’t dissolve in water) as the covalent bonds r too strong to break
40
Describe the 7 characteristics of diamond/ silicon
- tightly packed, rigid arrangements which allows heat to conduct well
- all 4 carbons bonded 4 times
- tetrahedral shape (109.5)
- can be cut to make gemstones
- VERY HIGH MELTING POINT due to many strong covalent bonds
- DOESN’T CONDUCT ELECTRICITY as there are no delocalised electrons to carry a charge
- INSOLUBLE covalent bonds are too strong to break
41
Describe the 5 characteristics of graphene
- a single layer of graphite
- 1 atom thick so lightweight and transparent
- made up of hexagonal carbon rings
- CONDUCTS ELECTRICITY as there is delocalised electrons free to carry a charge
- high strength dues to covalent bonds
42
What are the uses of graphene
- aircraft shells
- use in super computers and high speed computing
- smart phone screens
43
Define metallic bonding
- the electrostatic force of attraction between the positive metal ions and the delocalised electrons
44
What type of structure do metals have
- a giant metallic lattice structure
- positive metal ions surrounded by sea of delocalised electrons
- electrostatic attraction between positive metal ion and negative delocalised electrons
45
Complete the sentence:
The more electrons at atom can donate to the delocalised system the ….
- higher the melting point
46
Describe the characteristics of metals
- GOOD THERMAL CONDUCTORS—> as the electrons can transfer kinetic energy
- GOOD ELECTRICAL CONDUCTORS—> as the delocalised electrons can carry a charge
- HIGH MELTING PIOINT—> due to the strong electrostatic attractions attractions
- INSOLUBLE—> as the metallic bond is too strong to break
47
Describe the trend in melting point as you go across a period
- general increase as metal ions have an increasing positive charge, increasing the number of delocalised electrons and smaller ionic radius
- so a stronger metallic bond
48
Why is there a big drop in melting point from silicon to phosphorus
- because phosphorus is a simple covalent structure and so the only thing determining its melting point is weak London forces which are easy to overcome
49
Why does sulphur have a higher melting point that phosphorus
- because even though it is also just simple covalent molecule it has 8 sulfur atoms (S8) which means it has a bigger mass so more London forces making them stronger
50
Why does chlorine have one of the lowest melting point out of the period 3 elements
- dues to it being a simple covalent molecule and as it is a small molecule it has a smaller amount of London forces so they are weaker
51
Why does argon have the lowest melting point out for all the period 3 elements
- as it in monoatomic (only exists as individual atoms) so it has even smaller London forces
