Brainscape's Knowledge GenomeTM

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


See full index

Chemistry A Level > Atomic Structure > Flashcards

Atomic Structure Flashcards

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

Electron relative mass

A

1/1840

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

Nucleus

A
  • carries virtually all of the mass of the atom

* positively charged

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

The chemical nature of an atom is determined by…

A

The number of electron it has, and their arrangement

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

Relative atomic mass

A

The average mass of one atom compared to 1/12 that of the mass of one atom of carbon-12

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

Relative isotopic mass

A

The mass of one atom of an isotope compared to 1/12 of the mass of one atom of carbon-12

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

Relative molecular mass

A

The average mass of a molecule compared to 1/12 of the mass of one atom of carbon-12

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

Relative formula mass

A
  • the same as relative molecular mass, but for ionic compounds (as these are not molecules and giant covalent substances - the size of the molecule varies - it cannot have a precise molecular mass)
  • used for compounds with giant structures
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
8
Q

First ionisation energy

A

The energy required to remove one electron from each atom in one mole of gaseous atoms to produce one mole of gaseous ions with a 1+ charge

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

Equation for first ionisation energy

A

X(g) -> X+(g) -> e-

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

The first ionisation energy size provides information about…

A

… the force of attraction between the nucleus and the outer electrons

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

The attraction of an electron to the nucleus depends upon

A
  • the nuclear charge
  • the distance from the nucleus (greater = less attraction)
  • shielding (more inner electrons = less attraction)
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
12
Q

Nuclear charge

A
  • the amount of protons relative to the amount of electrons

* the higher the charge, the greater the attraction

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

Trend of first ionisation energies across a period

A
  • increase
  • electrons are added to the same electron shell (same distance, shielding)
  • nuclear charge increases; attraction to nucleus increases
  • therefore more energy is needed to remove the electron
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
14
Q

Trend across the periodic table for first ionisation energy

A
  • group 1 has the lowest

* noble gases have the highest

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

Trend for first ionisation energy across a group

A
  • decrease

* distance and shielding increase

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

Second ionisation energy

A

The energy required to remove one electron from each 1+ ion in one mole of gaseous ions to produce one mole of 2+ gaseous ions

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

Ionisation energy

A

The energy required to remove the readily available electron from an atom, by overcoming the attraction to the nucleus

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

General principles of first ionisation energy

A
  • as nuclear charge increases, there is a stronger attraction, so ionisation energy increases
  • as distance increases, attraction falls very rapidly - ionisation energy decreases
  • shielding - inner shells repel outer electrons away from the nucleus
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
19
Q

Spin-pairing

A

Causes repulsion; offsets attraction to the nucleus; paired e-s are removed more easily

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

Successive ionisions énergies

A
  • give us important information about the electronic structure of an element
  • determined which group an element is in
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
21
Q

Trend for successive ionisation energies

A
  • the second ionisation energy of an element is always greater than the first
  • after each ionisation, nuclear charge increases
  • the next electron must be removed from a greater attraction
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
22
Q

E.g. successive ionisation energies for potassium

A
  • first electron is removed from the outer-shell
  • second electron has to be removed from an inner-shell; requires more energy
  • the tenth electron has to be removed from the next inner shell
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
23
Q

If there is a large jump in successive ionisation energy

A

You have moved inwards a cell- all previous electrons were in the outer shell

24
Q

Electronic structure

A

Determines an atom’s chemical and physical properties

25
Q

Aufbau principle

A

The lowest energy sub-levels are occupied first

26
Q

Shell 1

A
  • s subshell

* 2 electrons

27
Q

Shell 2

A
  • s and p subshells

* 8 electrons

28
Q

Shell 3

A
  • s, p and d sub-shells
  • 18 electrons
  • can expand the octet
29
Q

Shell 4

A
  • s, p, d and f subshells

* 32 electrons

30
Q

The higher the energy of the shell…

A

… the more subshells it contains

31
Q

2n^2

A

Gives the number of electrons in any shell

Where n is the shell number

32
Q

Orbitals

A

An area occupied by an electron wave

33
Q

S orbitals are

A

Spherical

34
Q

Electrons in the first shell have the …

A

… same energy value, but opposite spin values

35
Q

Electrons in the p subshell

A
  • are of a higher energy level than those at the s

* found in 3 different regions of the atom: x, y and z

36
Q

The 4s subshell fills up…

A

… before the 3d subshell

Because an exactly half-full/full d sub-shell is particularly stable

37
Q

Hund’s rule

A

Electron orbitals are filled such that once an electron has entered an orbital of a p or d subshell, the orbital will not take on another electron until all other orbitals of the same energy level also have one electron

38
Q

s-block element

A

An element in which the last electron is in the s orbital

39
Q

Similar properties are portrayed by…

A

… elements in the same group; elements with the same number of electrons and electrons in the same type of sub-shell

40
Q

Copper and chromium

A
  • exceptions to the Aufbau principle
  • an electron from the 4s subshell goes to the 3d subshell to give chromium a half-full d subshell and copper a full subshell

Cr: 1s^22s^22p^63s^23p^64s^13d^5
Cu: 1s^22s^22p^63s^23p^64s^13d^10

41
Q

Noble gases

A

Have shells which are completely full

42
Q

[Ar]

A
  • has the same electron configuration as argon

* used because it is a noble gas

43
Q

Group 4 elements

A
  • tend to form giant covalent structures
  • elements towards the bottom of the group are metallic
  • Pb and Sn are metallic because they are large; electrons are further from the nucleus and there is more shielding; e-s are lost more easily
44
Q

Groups 5,6 and 7

A
  • generally have simple covalent structures

* bottom of the groups become more metallic

45
Q

Group 8

A
  • noble gases
  • have complete electron shells
  • monoatomic
46
Q

Moving horizontally…

A

… electron control increases as non-metallic character increases

47
Q

Moving downwards…

A

… electron control decreases due to larger atoms; electrons are further from the nucleus with greater shielding; metallic character increases

48
Q

Transition metals

A

A d-block element that forms one or more stable ions with incompletely filled d orbitals

49
Q

Melting temperatures across a period

A

Rise until group 4, then decrease

50
Q

Metals

A
  • strong forces of attraction between cations and delocalised electrons
  • increases across a period because charge increases (charge density increases)
  • no of delocalised electrons increases; attraction of metallic bonding increases
51
Q

Giant molecules

A
  • strong forces of attraction between nuclei and shared pair of electrons
  • melting temperature is at peak because vast amounts of strong covalent bonds must be broken
52
Q

Simple molecules

A
  • covalent bonds gold atoms strongly in the molecule, but there are weak intermolecular forces
  • melting temperatures are low because intermolecular forces are weak and roughly relative to the site of the molecules
53
Q

Noble gases have the lowest melting temperatures because…

A

… they are monatomic

54
Q

Phosphate

A

(PO4)3-

55
Q

Soluble:

A
  • Na+, K+, NH4+, NO3-
  • Cl- except Ag+ and Pb2+
  • (SO4)2- except Ca2+, Ba2+ and Pb2+
56
Q

Insoluble:

A
  • carbonate
  • oxide
  • hydroxide except Ca2+

Chemistry A Level (12 decks)

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