Brainscape's Knowledge GenomeTM

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


See full index

A Level Chemistry > Atomic Structure > Flashcards

Atomic Structure Flashcards

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

Define Nucleus

A

Most of the mass of an atom and very small. Contains protons and neutrons

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

Define electrons

A

Orbit the nucleus in shells and take most of the space of an atom.

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

Proton relative charge

A

+1

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

Proton relative mass

A

1

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

Neutron relative charge

A

0

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

Neutron relative mass

A

1

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

Electron relative charge

A

-1

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

Electron relative mass

A

1/1840

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

Mass number

A

number of protons and neutrons in the nucleus

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

Atomic number

A

Protons in the nucleus

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

Define ion

A

Different number of electrons and protons

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

Negatively charged ions

A

Gained an electron

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

Positively charged ions

A

Lost an electron to form a full shell

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

Define isotopes

A

Elements with the same proton number but a different number of neutrons

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

Define relative atomic mass

A

the weighted mean mass of an atom of an element, compared to 1/12th of the mass of an atom of carbon-12

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

Define relative molecular mass

A

mean mass of a molecule, compared to 1/12th of the mass of an atom of carbon-12

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

Relative isotopic mass

A

the mass of an atom of an isotope, compared to 1/12th of the mass of an atom of carbon-12

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

What axis is the abundance on?

A

y

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

What is abundance shown as

A

%

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

What is m/z

A

the mass of an isotope divided by charge. As most have just a + charge this is the same as the isotopic mass.

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

How to find relative atomic mass (mass spectra)

A

Relative atomic mass = (abundance A X m/z A) + (Abundance B X m/z B)
———————————————————————–
Total abundance

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
22
Q

Stages of predicting mass spectra

A

1) Write the percentages as decimals
2) Create a table showing the isotope combinations in a molecule of whatever your dealing with. Multiply the decimal from abundances of each isotope to get the relative abundance for each molecule.
3) Any molecules which are the same add the abundances up
4) . Divide all the relative abundances worked out before by the smallest value. This will give you a whole number ratio which can be used to predict your spectra.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
23
Q

Mass spectra- molecules
What do peaks show?
What is the significance of the M+

A

Fragments of the original molecules
The M+ peaks is the last peak or the molecular ion peak, which is the same as the relative molecular mass of the molecule

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
24
Q

Mass spectra- molecules

What is the m/z meaning

A

m/z is just the mass of a fragment divided by charge. As most have just a +1 charge this is the same as the fragment mass.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
25
Q

Electron configuration- subshells

A

Electron shells are split into 4 subshells

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
26
Q

S Subshell

Orbital number + how many electrons can they hold?

A

s – has 1 orbital can hold 2 electrons

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
27
Q

P subshell

Orbital number + how many electrons can they hold?

A

p – has 3 orbitals can hold 2 × 3 = 6 electrons

28
Q

d subshell

Orbital number + how many electrons can they hold?

A

d – has 5 orbitals can hold 2 × 5 = 10 electrons

29
Q

f subshell

How many orbitals + how many electrons can they hold?

A

f – has 7 orbitals can hold 2 × 7 = 14 electrons

30
Q

S orbital shape

A

The s orbital is spherical and the 2 electrons can move anywhere within this sphere.

31
Q

P orbital shape

A

There are 3 p orbitals in the shape of dumbbells each can hold up to 2 electrons can move anywhere within this shape.

32
Q

What is Hund’s Rule

A

‘all orbitals of equal energy will be singly occupied before any are doubly occupied’, meaning electrons will only pair up only when other degenerate orbitals aren’t empty

33
Q

What is Pauli’s Exclusion Principle?

A

‘no two electrons can have the same quantum numbers/ be in identical quantum states’. Hence two electrons in the same orbitals must have opposite spins

34
Q

With ions what is the rule for removing/adding electrons?

A

With ions you just add or remove electrons from the highest energy level first.

35
Q

What is different about Chromium and Copper’s electron configuration?

A

An electron from the 4s orbital moves into the 3d orbital to create a more stable half full or full 3d sub-shell respectively.

36
Q

Which groups are the s block elements?

A

1/2

37
Q

Which groups are d block elements?

A

transition elements

38
Q

Which groups are known as the p block elements?

A

3-8

39
Q

Where are the f block elements located?

A

The block at the bottom

40
Q

What does the electromagnetic spectrum show?

A

The electromagnetic spectrum shows types of radiation at different frequencies.

41
Q

What is line spectra ?

A

a way of identifying elements and is evidence for quantum shells

42
Q

Stages of mass spectrometry

A 
Vaporisation
Ionisation 
Acceleration
Deflection 
Detection
43
Q

Briefly describe vaporisation

A

The sample is converted to gaseous state so it can move through the machine

44
Q

Briefly describe ionisation

A

An electron gun fires fast-moving electrons at the particles to knock an electron away and form 1+ ions

45
Q

Briefly describe acceleration (mass spectrometry)

A

The ions are all accelerated so they have the same kinetic energy

46
Q

Briefly describe deflection

A

An electromagnet is used to deflect the ions by creating a magnetic field: the more massive the ions are, the less they are deflected, and the higher the charge on the ions, the more they are deflected

47
Q

Briefly describe detection

A

The ions that have made it through the whole machine are detected at the end and a mass spectrum is made electronically

48
Q

Why is the vacuum pump an important feature?

A

It is important that the ions produced in the ionisation chamber are able to move through the machine without bumping into air molecules; hence a vacuum is created

49
Q

Why is the electron gun an important feature?

A

The particles have to be ionised or else they would not be able to be repelled by the field created by the electromagnet as they would be uncharged

50
Q

What is changed during mass spectrometry?

A

The intensity of the magnetic field- to allow for the masses of ions to be calculated

51
Q

Define 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 1+ ions (measured in kJmol-1)

52
Q

Define successive ionisation energies

A

The energies required to remove each mole of electrons in turn from one mole of gaseous atoms/ions (measured in kJmol-1)

53
Q

How does the proton number affect ionisation energy?

A

The more protons there are in the nucleus, the more positively charged the nucleus is and the stronger the attraction for the electrons. Having more protons increases the ionisation energy

54
Q

What is the importance of the number of electrons between the outer electrons and the nucleus increasing?

A

The number of electrons increase so therefore the outer electrons feel less attraction towards the nuclear charge. This is caused by repulsion from electrons in inner shells which lessens this attraction; shielding is the name of this effect.
More greatly shielded electrons decreases the ionisation energy

55
Q

What is the importance of the distance between outer electrons and the nucleus?

A

Outer electrons in shells close to the nucleus are much more strongly attracted to it those that reside in much further away shells, and so therefore a larger distance between outer electrons and the nucleus decreases ionisation energy

56
Q

Explain why there is an increase in first ionisation energies of period 2 elements

A

Proton number increases, so there is a stronger attraction between the nucleus and outer electrons and so more energy is required to remove it

57
Q

Trends in first ionisation energy down a group

A

general decrease

58
Q

Increase nuclear charge …

A

Increased nuclear charge + Higher energy quantum shells + More shielding = General decrease

59
Q

What do big jumps in ionisation energy show?

A

Big jumps in ionisation energy show where an electron is being removed from a new shell and hence looking at an element’s successive ionisation energies can show which group it is in.

60
Q

Why is there a drop in between groups 2 and 3

A

In group 2- the outer electrons are in s orbital, while group 3 outer electrons are in the p orbital

P orbital is higher in energy- further distance, shielded meaning an decrease in ionisation energy

61
Q

Why is there a drop between group 5 and 6

A

group 5- singly occupied
group 6- 4th electron pairs with the 1st. Electron-electron repulsion which makes the p orbital easier to remove (unstable configuration)
Identical nucleus distance, identical shielding, decrease in ionisation energy

62
Q

Define orbital

A

Orbitals are 3D regions within an atom where electrons have a high probability of being found. Electrons must have opposite spins in order to exist within the same orbital.

63
Q

Explain how electron configuration determines the chemical properties of an element

A

Electronic configuration determines the likelihood that an element would donate/accept an electron / multiple electrons.

This therefore determines how an element may choose to bond with other elements. The nature of this bonding determines an element’s physical characteristics as well as its chemical characteristics.

64
Q

Define periodicity

A

Periodicity is a regularly repeating pattern of atomic, chemical, and physical properties as atomic number increases.

65
Q

Trends in electronegativity

A

Electronegativity increases across a period due to an increase in effective nuclear charge.

This decreases the atomic radius, and the combination of the greater nuclear charge and smaller distance between electrons from other atoms, increases the atoms tendency to attract bonding electrons, hence increasing its electronegativity.

66
Q

What is atomic radius

A

the farthest point form the nucleus where an electron can be found.

67
Q

Trends in atomic radius across a period

A

Atomic radius decreases across a period due to an increase in effective nuclear charge.

This decreases the atomic radius because it attracts the electrons more strongly to the nucleus, hence pulling them closer towards it, decreasing the distance between them.

A Level Chemistry (12 decks)

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