Brainscape's Knowledge GenomeTM

Browse over 1 million classes created by top students, professors, publishers, and experts.


See full index

Module 3 (7-10) > Chapter 7 > Flashcards

Chapter 7 Flashcards

1
Q

How did Mendeleev order the elements in his periodic table?

A

By atomic mass

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
2
Q

How did Mendeleev put elements into groups?

A

He put elements with similar properties in groups, if they didn’t fit he would swap elements around and leave gaps. (assumed atomic mass was incorrect or some elements were yet to be discovered).

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
3
Q

How are elements organised in the periodic table now?

A

By increasing atomic number

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
4
Q

How were elements assigned to the period/group that they are in?

A

Atoms with the same number of shells in the same period.
Atoms with similar electronic configurations in the outer shell are placed in the same group.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
5
Q

What are outer electrons called?

A

Valence electrons

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
6
Q

What is periodicity?

A

A repeating trend in properties across each period.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
7
Q

What properties do we investigate when looking at periodicity?

A

-electron configuration
-Ionisation energy
-structure
-melting points

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
8
Q

What are the 4 blocks that elements belong to?

A

s-block
p-block
d-block
f-block

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
9
Q

What is the periodic trend of electronic configuration about?

A

The block that an element belongs to as this changes across a period.

The block corresponds to an elements highest energy sub shell.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
10
Q

What is ionisation energy?

A

How easily an atom can lose electrons to form positive ions

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
11
Q

What is the first ionisation energy?

A

The energy required to remove one mole of electrons from one mole of atoms of an element in the gaseous state to form one mole of gaseous ions.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
11
Q

What is the first ionisation energy?

A

The energy required to remove one mole of electrons from one mole of atoms of an element in the gaseous state to form one mole of gaseous ions.

He —-> He+ +e-

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
12
Q

What are the factors that affect ionisation energy?

A

1- atomic radius
2- nuclear charge
3- electron shielding

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
13
Q

How does atomic radius affect ionisation energy?

A

electrons in shells that are further away from the nucleus have less nuclear attraction.

The further the outer electron shell is from the nucleus, the lower the ionisation energy. (down a group the number of shells increases)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
14
Q

How does nuclear charge affect ionisation energy?

A

The more protons there are in the nucleus of an atom, the greater the attraction between the nucleus and the outer electrons so the ionisation energy increases.

There are more electrons in the outer shells of elements across a period.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
15
Q

How does electron shielding affect ionisation energy?

A

The shielding effect is when the electrons in full inner shells repel electrons in outer shells preventing them to feel the full nuclear charge.

This reduces the attraction between the nucleus and outer electrons and results in a decreased ionisation energy (down a group the number of shells increases)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
16
Q

What is the trend in ionisation energy across a period and why?

A

It increases because across a period, the nuclear charge increases.

The distance between the nucleus and outer electron remains reasonably constant (no significant change in atomic radius)
The shielding by inner shell electrons remains the same.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
17
Q

What is the trend in ionisation energy down a group and why?

A

It decreases because the atomic radius increases and the shielding (by inner shell electrons) increases.

Therefore, the attraction between the nucleus and the outer electrons decreases

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
18
Q

What is the second ionisation energy?

A

The 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.

He+ —> He2+ +e-

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
19
Q

Do successive ionisation energies increase or decrease?

A

Increase as the nuclear attraction of the remaining electron increases. After the first electron is lost the second is pulled closer

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
20
Q

Why is there a rapid decrease in ionisation energy between the last element in one period and the first element in the next period?

A

The increased distance between the nucleus and the outer electrons

The increased shielding by inner electrons

These two factors outweigh the increased nuclear charge

21
Q

What is each peak in first ionisation energy?

A

A noble gas. (The end of each period)

22
Q

Why is there a slight decrease in the first ionisation energy between beryllium and boron?

A

Beryllium has an electron configuration of 1s2 2s2

Boron has an electron configuration of 1s2 2s2 2P1.

The ionisation decreases slightly as the electron in the 2p sub-shell is easier to remove than one of the electrons in the 2s sub-shell due to the increase in atomic radius/shielding.

23
Q

Why is there a slight decrease in the first ionisation energy between nitrogen and oxygen?

A

Oxygen has 2 electrons in one orbital, when there are 2 electrons in an orbital they have opposing spin so repel. This means that the electron is easier to remove so the ionisation energy for oxygen is less.

24
Q

When does the changeover from metal to non-metal happen in the periodic table?

A

a diagonal line down from boron.

right is non-metal
left Is metal

25
Q

What state are all metals at room temp?

A

solids.

but mercury is a liquid

26
Q

What is the one constant property of all metals?

A

Their ability to conduct electricity.
Charge has to move through a rigid solid structure.

27
Q

What structure are metal atoms in?

A

Metallic lattices- metal atoms are tightly packed together and the electrons in their outer shells are free to move throughout the structure.

When the electrons are delocalised, the metal atoms become positively charged ions.
The positive charges repel each other and keep the neatly arranged lattice in place.
There are very strong forces between the positive metal centres and the ‘sea’ of delocalised electrons.

28
Q

What are the free-moving electrons called?

A

delocalised electrons

29
Q

What is metallic bonding?

A

The strong electrostatic attraction between cations and delocalised electrons.

cations are fixed in position and maintain the structure/ shape of the metal.

delocalised electrons are mobile and are able to move throughout the structure.

30
Q

What are the properties of metals?

A

-Strong metallic bonds (attraction between cations and delocalised electrons)
-high electrical conductivity
-high melting/boiling points

31
Q

Why do metals have high melting/boiling points?

A

A lot of energy is required to overcome the strong electrostatic forces of attraction between positive ions and the ‘sea’ of delocalised electrons.

32
Q

why are metals not soluble?

A

Metals do not dissolve. There is some interaction between polar solvents and charges in the metallic lattice but these lead to reactions, rather than dissolving e.g. sodium and water.

33
Q

Why do metals have electrical conductivity?

A

They conduct electricity in both solid and liquid states. This is due to the delocalised electrons which are free to move/carry charge around the structure.

34
Q

What are covalent bonds?

A

Bonds between non-metals where there Is a shared pair of electrons between the atoms.

35
Q

How do many non-metallic elements exist?

A

As simple covalently bonded molecules.

36
Q

What structure do simple covalently bonded molecules form in solid state?

A

Simple molecular lattice structure

37
Q

What forces hold a simple molecular lattice structure together?

A

Weak intermolecular forces

38
Q

Do elements with a simple molecular lattice structure have low or high melting/boiling points?

A

Low

39
Q

What structure do the non-metals boron, carbon and silicon have?

A

giant covalent lattice

40
Q

What forces do the non-metals boron, carbon and silicon have?

A

strong covalent bonds

41
Q

What is the structure of diamond?

A

Diamond is a giant covalent lattice of carbon atoms. Each carbon is covalently bonded to 4 others in a tetrahedral arrangement with a bond angle of 109.5.

Diamond is the hardest substance known.

42
Q

What is the structure of Graphite?

A

In graphite, each carbon atom is bonded to three others in a layered structure.
The layers are made of hexagons with a bond angle of 120.
The spare electrons are delocalised and occupy the space between the layers.
All atoms in the same layer are held together by strong covalent bonds. However, the layers are held together by weak intermolecular forces (meaning that they can slide over eachother)

43
Q

What is the structure of Graphene?

A

Graphene is made of a single layer of carbon atoms that are bonded together in a repeating pattern of hexagons. (very very thin)

44
Q

What is the structure of Silicon oxide?

A

Same structure as diamond. - giant covalent lattice made of tetrahedral units all bonded by strong covalent bonds.

Each silicon is shared by 4 oxygens and each oxygen is shared by 2 silicons. (SiO2)

45
Q

What properties do substances with a giant covalent lattice structure?

A

-High melting and boiling point
-Insoluble in solvents
-non conductors of electricity (exceptions are graphene and graphite)

46
Q

Why do substances with a giant covalent lattice have a high melting/boiling point?

A

These compounds have a large number of covalent bonds linking the whole structure. These require a lot of energy to break.

47
Q

Why are substances with a giant covalent lattice insoluble in solvents?

A

As the covalent bonds holding together the atoms in the lattice are far too strong to be broken by interaction with solvents.

48
Q

Why are giant covalent lattices non-conductors of electricity?

A

They have no free electrons- all of their electrons are involved in a covalent bond.

49
Q

Why can graphene and graphite conduct electricity?

A

They have delocalised electrons between the carbon layers/ between their atoms.

50
Q

What is the trend in melting points across a period?

A

melting point increases from group 1-4

1-3= strong metallic bonding (increases in strength due to the increased forces of attraction due to more electrons in the outer shell).
group 4= giant covalent structure with many strong covalent bonds (need a lot of energy to be broken)

There is a sharp decrease in melting point from group 4-5.

Group 5= simple molecular structures with weak London forces between molecules (which requires little energy to overcome)

Module 3 (7-10) (4 decks)

Brainscape helps you reach your goals faster, through stronger study habits.
© 2024 Bold Learning Solutions. Terms and Conditions