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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

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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?

A

2

21
Q

How many electrons can the 2nd shell hold?

A

8

22
Q

How many electrons can the 3rd shell hold?

A

18

23
Q

How many electrons can the 4th shell hold?

A

32

24
Q

What is an orbital?

A

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

25
Q

What are the 4 types of orbitals?

A
  • s
  • p
  • d
  • f
26
Q

What is the shape of the s orbital?

A

spherical

27
Q

What is the shape of the p orbital?

A

dumb-bell shape

28
Q

How many orbitals are found in a s sub-shell?

A

1 orbital

29
Q

How many orbitals are found in a p sub-shell?

A

3 orbitals

30
Q

How many electrons can be found in a s sub-shell?

A

2 electrons

31
Q

How many electrons can be found in a p sub-shell?

A

6 electrons

32
Q

How many orbitals are present in a d subshell?

A

5 orbitals

33
Q

How many electrons can be in a d subshell?

A

10 electrons

34
Q

How many orbitals are found in a f subshell?

A

7 orbitals

35
Q

How many electrons can be found in a f subshell?

A

14 electrons

36
Q

What are the rules for electron arrangement in a shell?

A
  • 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
Q

Why does 4s orbital fill before 3d orbital?

A

4s orbital has a lower energy than 3d orbital before it is filled

38
Q

What is meant by periodicity?

A

The repeating trends in chemical and physical properties

39
Q

What change happens across each period?

A

elements change from metals to non-metals

40
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 under standard conditions

41
Q

What is the equation for the first ionisation energy of magnesium?

A

Mg(g) —-> Mg+(g) + e-

42
Q

What are the factors that affect ionisation energy?

A
  • atomic radius
  • nuclear charge
  • electron shielding or screening
43
Q

Why does first ionisation energy decrease between group 2 and 3?

A
  • group 3 outer most electrons are in p orbitals

- in group 2 outer electrons are in s orbitals so are easier to remove

44
Q

Why does first ionisation energy decrease between group 5 and 6?

A
  • 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
Q

Does first ionisation energy increase or decrease between the end of one period and the start of another?

A
  • decrease
  • there is an increase in atomic radius
  • so an increase in shielding
46
Q

Does first ionisation increase or decrease down a group?

A
  • 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
Q

Describe the structure, bonding and forces between every element across period 2

A

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
Q

Describe the structure, bonding and forces between every element across period 3

A
  • 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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