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A-Level Chemistry > Topic 1: Atomic structure and the Periodic Table > Flashcards

Topic 1: Atomic structure and the Periodic Table Flashcards

1
Q

What was stated in Dalton’s atomic theory?

A
  • atoms are tiny particles made of elements
  • atoms cannot be divided
  • all the atoms in an element are the same
  • atoms of one element are different to those of other element
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2
Q

What did Thompson discover about electrons?

A
  • they have a negative charge
  • they can be deflected by electromagnetic fields
  • they have very small mass
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3
Q

Explain the current model of the atom

A
  • protons and neutrons are found in the nucleus
  • electrons orbit the nucleus in shells
  • most of the atom’s mass is in the nucleus
  • most of the atom is empty space between the nucleus and the electrons
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4
Q

What is the charge for a proton and electron?

A
  • proton is +1

- electron is -1

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5
Q

Which particle has the same mass as a proton?

A

neutron

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6
Q

Which 2 particles make up the majority of an atom’s mass?

A

protons and neutrons

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7
Q

What does the atomic number show about an element?

A

The number of protons

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8
Q

How is mass number calculated?

A

number of protons+ number of neutrons

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9
Q

How to calculate the number of neutrons?

A

number of neutrons= mass number - atomic number

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10
Q

Define isotope

A

Atoms of the same element with different number of neutrons and the same number of electrons but different mass numbers

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11
Q

Why do different isotopes of the same element react in the same way?

A
  • neutrons have no impact on the chemical reactivity

- reactions involve electrons, isotopes have the same number of electrons in the same arrangement

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12
Q

Define relative atomic mass

A

The weighted mean mass of an atom of an element compared with 1/12 of the mass of an atom of carbon-12

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13
Q

Define relative isotopic mass

A

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

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14
Q

The relative isotopic mass is the same as which number?

A

mass number

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15
Q

What 2 assumptions are made when calculating mass number?

A
  • contribution of the electron is neglected

- mass of both the proton and neutron are taken as 1.0 u

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16
Q

How do you calculate the relative formula mass and relative molecular mass?

A

add the relative atomic masses of each of the atom making up the molecule or formula

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17
Q

What are the uses of mass spectrometry?

A
  • identify unknown compounds
  • find relative abundance of each isotope of an element
  • determine structural information
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18
Q

What does the principal quantum number indicate?

A

The shell occupied by the electrons

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19
Q

What is a shell?

A

A group of orbitals with the same principle quantum number

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20
Q

How many electrons can the 1st shell hold?

21
Q

How many electrons can the 2nd shell hold?

22
Q

How many electrons can the 3rd shell hold?

23
Q

How many electrons can the 4th shell hold?

24
Q

What is an orbital?

A

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

25
What are the 4 types of orbitals?
- s - p - d - f
26
What is the shape of the s orbital?
spherical
27
What is the shape of the p orbital?
dumb-bell shape
28
How many orbitals are found in a s sub-shell?
1 orbital
29
How many orbitals are found in a p sub-shell?
3 orbitals
30
How many electrons can be found in a s sub-shell?
2 electrons
31
How many electrons can be found in a p sub-shell?
6 electrons
32
How many orbitals are present in a d subshell?
5 orbitals
33
How many electrons can be in a d subshell?
10 electrons
34
How many orbitals are found in a f subshell?
7 orbitals
35
How many electrons can be found in a f subshell?
14 electrons
36
What are the rules for electron arrangement in a shell?
- electrons are added one at a time - lowest available energy level is filled first - each energy level must be filled before the next one can fill - each orbital is filled singly before pairing - 4s is filled before 3d
37
Why does 4s orbital fill before 3d orbital?
4s orbital has a lower energy than 3d orbital before it is filled
38
What is meant by periodicity?
The repeating trends in chemical and physical properties
39
What change happens across each period?
elements change from metals to non-metals
40
Define first ionisation energy
The energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous 1+ ions under standard conditions
41
What is the equation for the first ionisation energy of magnesium?
Mg(g) ----> Mg+(g) + e-
42
What are the factors that affect ionisation energy?
- atomic radius - nuclear charge - electron shielding or screening
43
Why does first ionisation energy decrease between group 2 and 3?
- group 3 outer most electrons are in p orbitals | - in group 2 outer electrons are in s orbitals so are easier to remove
44
Why does first ionisation energy decrease between group 5 and 6?
- group 5 electrons are in the p orbital as single electrons - group 6 outermost electrons are spin paired, with some repulsion so they are slightly easier to remove
45
Does first ionisation energy increase or decrease between the end of one period and the start of another?
- decrease - there is an increase in atomic radius - so an increase in shielding
46
Does first ionisation increase or decrease down a group?
- decrease - shielding increases ---> weaker attraction - atomic radius increases ---> distance between the outer electrons and nucleus increases ---> weaker attraction - increase in the number of protons is outweighed by increase in distance and shielding
47
Describe the structure, bonding and forces between every element across period 2
Li + Be ---> giant metallic, strong attraction between positive ions and delocalised electrons, metallic bonding - B + C ---> giant covalent, strong forces between atoms, covalent - N2,O2,F2,Ne ---> simple molecular, weak intermolecular forces between molecules, covalent bonding within molecules and intermolecular forces between molecules
48
Describe the structure, bonding and forces between every element across period 3
- Na, Mg, Al ---> giant metallic, strong attraction between positive ions and delocalised electrons, metallic bonding - Si ---> giant covalent, strong forces between atoms, covalent - P4, S8, Cl2, Ar ---> simple molecular, weak intermolecular forces between molecules, covalent bonding within molecules and intermolecular forces between molecules

A-Level Chemistry (20 decks)

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