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Mocks > Chemistry mock p1 > Flashcards

Chemistry mock p1 Flashcards

1
Q

What is periodicity?

A

Periodicity is the trends of elements across a period

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2
Q

Does atomic radius increase or decrease across a period and why?

A

Atomic radius decreases across a period because the proton number increases, so the positive charge of the nucleus increases. Electrons are pulled closer to the nucleus, making the atomic radius smaller. Extra electrons are added to outer energy level so add no shielding effect.

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3
Q

Does ionisation energy increase or decrease across a period and why?

A

Ionisation energies increase across a period because the proton number increases so there is a stronger nuclear attraction. Extra electrons add no shielding effect.

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4
Q

Why is there a drop in ionisation energy between group 2 and 3?

A

Due to 3p orbital being at a higher energy level than 3s, so is further away from the nucleus. The 3p orbital has shielding from 3s orbital. Overrides effect of increasing nuclear charge

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5
Q

Why is there a drop between group 5 and 6?

A

Electron repulsion
Easier to remove an electron from a partially filled p orbital than a singly filled orbital due to electron repulsion.

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6
Q

How do metal ions make carbonates and nitrates unstable?

A

The presence of metal cations polarises the anion electron cloud, making it less stable

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7
Q

Does thermal stability of carbonates and nitrates increase or decrease down a group and why?

A

Increases, as larger cations cause less distortion than small cations as they have a lower charge density. Down a group, the cations get larger making carbonates/ nitrates more thermally stable.

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8
Q

Are group 1 or 2 carbonates/nitrates more stable and why?

A

Group 1 compounds are more thermally stable as group 2 has a higher charge, and therefore a higher charge density, meaning there is more distortion of the anion and it is less stable.

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9
Q

Do group 1 carbonates decompose when heated?

A

They are stable when heated with a Bunsen burner and therefore do not decompose.
Apart from Li2c03 which decomposes to Li20 and C02

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10
Q

What do group 1 nitrates decompose into when heated?

A

Group 1 nitrates decompose to form the nitrite and oxygen.
2MNO3 (s) –> 2MN02 (S) + 02(g)
Apart from Li2CO3 that decomposes into Li2O, NO2 + 02

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11
Q

What do group 2 carbonates decompose to form when heated?

A

They decompose to form oxide and carbon dioxide
MCO3 (s)–> MO(s) + CO2 (g)

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12
Q

What do group 2 nitrates decompose to form when heated?

A

Oxide, Nitrogen dioxide and oxygen
2M(NO3)2 (s) –> 2MO(s) + 4NO2 (g) + O2 (g)

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13
Q

How is the thermal stability of nitrates tested?

A

Time taken to produce a certain amount of oxygen
(enough to relight a glowing splint)
Time taken for an amount of NO2 to be produced
(toxic so needs to be done in a fume cupboard)

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14
Q

How is the thermal stability of carbonates tested?

A

Time taken to produce a certain amount of CO2
Test for CO2 using lime water- turns it cloudy

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15
Q

What flame colour does Lithium produce?

A

Red

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16
Q

What flame colour does sodium produce?

A

Orange/ Yellow

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17
Q

What flame colour does potassium produce?

A

Lilac

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18
Q

What flame colour does rubidium produce?

A

Red

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19
Q

What flame colour does caesium produce?

A

Blue

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20
Q

What flame colour does calcium produce?

A

Brick-red

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21
Q

What flame colour does strontium produce?

A

Crimson

22
Q

What flame colour does barium produce?

A

Green

23
Q

How is a flame test carried out?

A

Clean nichrome wire by dipping in HCL and holding in a Bunsen burner flame
Dip nichrome wire into sample and hold on the edge of a blue/ naked flame until colour is observed

24
Q

Why does Magnesium not produce a colour?

A

The light emitted is outside the visible spectrum for humans

25
Q

Why do metal ions produce a colour when passed through a Bunsen burner?

A

The energy absorbed from the flame causes the electron to move to a higher energy level (in a excited state)
The electron drops back down to a lower energy level (ground state) due to them being unstable
This releases energy in the form of visible light

26
Q

What determines the colour of the flame produced?

A

The difference in energy between the lower and higher energy levels

27
Q

Draw the shape of the p and s-orbital

A
28
Q

What is the order of subshell notation?

A

1s2 2s2 2p6 3s2 3p6 4s2 3d10

29
Q

Where are s-block elements located?

A

S-block elements are group 1 and 2

30
Q

Where are d-block elements located?

A

D-block elements are the transition metals

31
Q

Do electrons fill orbitals singly or as pairs?

A

Singly due to spin pairing repulsion

32
Q

Where are p-block elements located?

A

P-block elements are non-metals

33
Q

How does emission spectra support the idea of quantum shells?

A

Electrons can only exist in fixed orbits or shells, with a fixed energy. When an electron moves between shells electromagnetic radiation is admitted or absorbed

34
Q

What is the definition of ionisation energy?

A

X(g) —> X+ (g) + e-

35
Q

What is the definition of an ionic bond?

A

An ionic bond is the electrostatic attraction between two oppositely charged ions.

36
Q

How does ionic charge affect the strength of an ionic bond?

A

The greater the charge of the ion the stronger the ionic bond and therefore a higher melting/ boiling point

37
Q

How does ionic radii affect the strength if an ionic bond?

A

Smaller ions have further distance and therefore a stronger electrostatic attraction, making a stronger bond and a higher melting/ boiling point

38
Q

Does ionic radii increase or decrease down a group and why?

A

Ionic radii increases down a group because extra electron shells are added.

39
Q

What is an isoelectronic ion?

A

Ions of different atoms that have the same number of electrons.

40
Q

Does the ionic radius of a set of isoelectronic ions increase or decrease through the set and why?

A

Decreases because the number of electrons stays the same, but the number of protons increases, which increases the nuclear charge and draws the electrons closer

41
Q

What are ionic crystals?

A

Ionic crystals are giant lattices of ions

42
Q

What are the physical properties of ionic bonds?

A

High melting points- ions held together by a strong attraction
Soluble in polar solvents (e.g water) not in non-polar solvents- because particles are charged so pulled apart by polar molecules
Don’t conduct electricity when solid only when molten or dissolved- ions free to move and carry charge in a liquid or dissolved
Can’t be shaped- due to the repulsion between ions making them brittle

43
Q

What is the definition of a covalent bond?

A

Two atoms share electrons so they’ve both got full outer shell electrons. A covalent bond is the strong electrostatic attraction between two positive nuclei and the shared electrons in the bond.

44
Q

What is bond enthalpy?

A

Bond enthalpy is the distance where the attractive and repulsive forces balance each other.

45
Q

How is bond enthalpy related to length of the bond?

A

The higher the electron density between the nuclei, the stronger the attraction between the atoms, the higher the bond enthalpy, and the shorter the bond length.

46
Q

What is a dative covalent bond?

A

A dative covalent bond is a where one atom donates both electrons to a bond. An ammonium ion has a dative covalent bond.

47
Q

What are the properties of giant covalent structures?

A

Very high melting point- very strong bonds
Often extremely hard- due to very strong bonds
good thermal conductors-vibrations travel easily through the stiff latices
insoluble- due to covalent bonds
can’t conduct electricity- no charged ions or free electrons

48
Q

What is a redox reaction?

A

A reaction where reduction and oxidation simultaneously occur.

49
Q

What is the definition of a reduction reaction?

A

gain of electrons/ loss of oxygen/ gain of hydrogen

50
Q

What is the definition of an oxidation reaction?

A

loss of electrons/ gain of oxygen/ loss of hydrogen

51
Q

What is electron affinity?

A

X(g)—> X+ + e-

52
Q
A

Mocks (2 decks)

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