Brainscape's Knowledge GenomeTM

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


See full index

Chemistry > Atomic structure and the periodic table > Flashcards

Atomic structure and the periodic table Flashcards

You may prefer our related Brainscape-certified flashcards:
AP® Chemistry flashcards Periodic Table flashcards Chemistry 101 flashcards Organic Chemistry 101 flashcards
1
Q

What is the relative isotopic mass?

A

the mass of an isotope of an atom relative to the mass of 1/12 the atom of carbon-12

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

What is the relative atomic mass?

A

The average mass of an atom relative to the mass of 1/12 of an atom of carbon-12

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

what is the principal quantum number (n) ?

A

The number given to each shell e.g. the third shell has n=3

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

What happens when we have 2 electrons in the same orbital?

A

They must have opposite spins

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

What is an atomic orbital?

A

A region around the nucleus that can hold up to 2 electrons with opposite spins

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

What are the different types of atomic orbitals?

A

S, P, D, F

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

Define details of S orbitals

A
  • It has a spherical shape and every electron shell contains a single s orbital
  • The size of the orbitals increase with increasing shell number
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
7
Q

Define details of p orbitals

A
  • There are 3 p orbitals
  • They are shaped like a dumb bell
  • They are orientated perpendicular to one another
  • With increasing shell number, the lobes become longer and larger
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
8
Q

How many d orbitals are there?

A

5

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

How many f orbitals are there?

A

7

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

How are electrons imagined?

A

As small spinning charges which rotate around their own axis in either clockwise or anticlockwise direction. The spin creates a tiny magnetic field with each pole facing up or down

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

What happens to electrons with the same spin?

A

they repel each other, known as spin-spin repulsion

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

What do sub shells contain?

A

one or more atomic orbitals

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

What will electrons do first?

A

They will occupy separate orbitals in the same sub shells to minimise this repulsion and have their spin in the same direction (known as Hund’s rule)

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

What will electrons do after this?

A

They will then pair up, with a second electron being added to the first p orbital with its spin in the opposite direction (known as Hund’s rule)

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

How much space does each electron occupy?

A

Even though there is repulsion between them, they occupy the same region of space in orbitals

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

What is the Pauli Exclusion principle?

A

all electrons in an atom must be in different orbitals or have different spins

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

What is the reason for the Pauli Exclusion principle?

A

The energy required to jump to a higher empty orbital is greater than the inter-electron repulsion so they occupy lower energy levels first

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

What is the aufbau principle?

A

Electrons enter the lowest energy orbital available

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

What is Hund’s rule?

A

Electrons prefer to occupy orbitals on their own, and only pair up when no empty orbitals of the same energy are available

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

What is the order the electrons fill the orbitals in?

A

1s - 2, 2s - 2, 2p - 6, 3s - 2, 3p - 6, 4s - 2, 3d - 10, 4p - 6

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
21
Q

Why does 4s come before 3d in the orbital order?

A

It has a lower energy level than 3d

22
Q

Which elements do not have the expected electron structure and why?

A

Chromium and Copper
To decrease the electron repulsion and get a lower overall energy level, the electrons occupy the 3d orbital first. (for copper its 3d10, 4s1 and not 3d9 and 4s2)

23
Q

What is ionisation energy?

A

A measure of the amount of energy needed to remove electrons from atoms

24
Q

What is the first ionisation energy?

A

The energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous positive ions

25
Q

Give an example of the first ionisation energy?

A

Na(g) = Na+(g) +e-

26
Q

What does the value of the 1st I.E. depend on?

A

The electronic structure

27
Q

Why is the ionisation energy value for helium higher than for hydrogen?

A
  • There are 2 protons in the nucleus instead of 1, meaning that the nuclear charge is greater
    -The distance of the outer electron to the nucleus decreases (decreased atomic radius)
  • The outer electron is held more tightly to the nucleus so more energy is required to pull it out of the atom
28
Q

Why is the Ionisation energy for lithium lower than for helium?

A
  • Although it has 3 protons instead of 2, the valence electron is in another shell
  • Filled inner shells exert a shielding effect which lowers the effective nuclear pull
  • It is further away from the nucleus so the electron has a lower nuclear attraction and less energy is required to pull it away
29
Q

Why does the 1st ionisation energy show a general increase across a given period?

A

There is an increased nuclear charge, but no extra electron shielding so the distance from the nucleus stays the same but there is a greater nuclear charge

30
Q

Why does the ionisation energy decrease down a group?

A
  • All though there is an increase in nuclear charge, there is an increase in shells so the distance of the outer electron to the nucleus increases
  • Shielding effect increases so the attraction of valence electrons to the nucleus decreases (increased atomic radius)
  • The outer electron is held more loosely to the nucleus so becomes easier to remove
31
Q

Why is the ionisation energy for oxygen lower than for nitrogen?

A

All though there is no new electron shell here, the extra electron has paired up with one of the electrons already in one of the 2p orbitals. The repulsive force between the 2 paired-up electrons means that less energy is required to remove one of them

32
Q

What are successive ionisation energies of an element?

A

More than one electron can be removed from an atom and each time you remove an electron there is a successive ionisation energy

33
Q

What is the second ionisation energy?

A

The energy required to remove 1 mole of electrons from 1 mole of gaseous 1+ positive ions to form one mole of 2+ positive ions

34
Q

What is an example of the second ionisation energy?

A

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

35
Q

Why are successive ionisation energies always greater than the previous one?

A

The electron is being pulled away from a more positive species

36
Q

When do large increase occur?

A

When there is a change of shell as there is a big decrease in shielding

37
Q

What can large increases be used for?

A

To predict the group of an unknown element

38
Q

What happens wherever there has been a large increase in ionisation energy?

A

There has been a change in energy level from which the electron has been removed

39
Q

Why does the ionisation energy across a period increase?

A
  • The nuclear charge increases
  • The atomic radius of the atoms decrease so outer shell is pulled closer to nucleus
  • Shielding remains constant
  • Becomes harder to remove an electron as you move across so more energy is needed
40
Q

What do elements across a period show?

A

Repeating patterns in physical and chemical properties

41
Q

Why is there a decreased atomic radius across a nucleus?

A

There is an increased nuclear charge, so a higher pull from the nucleus

42
Q

Why is there a small drop in the 1st ionisation energy between Mg and Al?

A

Mg has outer electrons in the 3s sub shell whereas Al starts to fill the 3p sub shell. Al’s electron is slightly easier to move as 3p electrons are higher in energy

43
Q

What is electronegativity?

A

The power of an atom to attract the 2 electrons in a covalent bond

44
Q

What is the trend of electronegativity across a period?

A

Across a period, there are more protons in the nucleus so a smaller atomic radius so there is a stronger attraction between the nucleus and 2 electrons in a covalent bond

45
Q

How is melting and boiling points affected by the size of an atom in metallic bonding?

A
  • The bigger the atom, the more electrons there are, so the greater the forces
  • The greater the number of electrons holding them together, the greater the B.p and M.p
46
Q

What also makes the bonding stronger in metallic bonding?

A

A smaller sized ion with a greater positive charge

47
Q

Do metallic bonds need higher or lower energy in general?

A

Higher

48
Q

What molecular term is used to describe silicone?

A

Macromolecular

49
Q

Why does silicone have a very high boiling point and melting point?

A

It has many strong covalent bonds between atoms so a high energy is needed to break them

50
Q

What are the examples of simple molecular?

A

cl2(g), s8(s), P4(s) - (this is an example for one period but it counts for all the other periods as well)

51
Q

What are the forces between the simple molecular?

A

There are weak London forces between molecules so little energy is needed to break them - it has a low melting and boiling point

52
Q

Why does S8 have a higher melting point than P4?

A

It has more electrons (s8=128, p4=60) so it has stronger London forces between molecules

53
Q

What is Argon classified as?

A

Monoatomic - there are weak London forces between the atoms

Chemistry (7 decks)

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