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a-level chemistry > Mod 2 Chapter 5 > Flashcards

Mod 2 Chapter 5 Flashcards

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AP® Chemistry flashcards Periodic Table flashcards
1
Q

Maximum number of electrons in each shell formula?

A

2n^2

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2
Q

As shell number increases

A

Energy increases (as shells are fixed energy levels)

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3
Q

What is the principal quantum number?

A

Shell number/energy level number

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4
Q

What is an atomic orbital?

A

Region around nucleus that can hold up to two electrons with opposite spin (high probability of finding an electron)

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5
Q

How many electrons can orbitals hold?

A

S, p, d and f can hold up to 2 electrons with each orbital having a different shape

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6
Q

Why do electrons have opposite spins?

A

To prevent repulsion

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7
Q

S-orbitals

A

Electron cloud in shape of a sphere and each shell from n=1 contains 1 s-orbital ; greater the shell number, greater the radius of its s-orbital

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8
Q

P-orbitals

A

Shape of a dumbell (x, y and a) ; one orbital can contain max 2 electrons BUT there are 3 p orbitals at right angles to one another therefore each shell from n=2 contains 3 p-orbitals

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9
Q

The greater the shell number n

A

Further the p-orbital is from the nucleus

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10
Q

D orbitals

A

Each shell from n=3 contains 5 d-orbitals

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11
Q

F orbitals

A

Each shell from n=4 contains 7 f-orbitals

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12
Q

What are sub shells?

A

Orbitals of the same type grouped together

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13
Q

How many electrons can fit in the p subshell?

A

3*2 = 6 electrons

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14
Q

How many electrons can fit in the d subshell?

A

5*2 = 10

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15
Q

How many electrons can fit in the f subshell?

A

7*2 = 14

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16
Q

What is important to note with the filling of the 3rd and 4th shell?

A

Highest energy level of 3rd shell (3d subshell) overlaps the lowest energy level in the 4th shell (4s subshell) - THEREFORE 4S SUBSHELL FILLS BEFORE 3D AND LOSES ELECTRONS BEFORE 3D

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17
Q

Why do electrons have opposite spins?

A

Electrons are negatively charged and repel one another - therefore there spin is needed to help counteract the repulsion between the negative charges of the two electrons

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18
Q

How do electrons occur subshells?

A

If there is more that one orbital in subshell, then one electron occupies each orbital and only then can pairing take place with opposite spin

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19
Q

Electron configuration of Krypton?

A

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

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20
Q

How can shorthand electron configuration be expressed?

A

Previous noble gas in the table + the outer electron sub shells (useful for emphasising the similarities in electron configuration of outer shell - same last number of each element in group)

21
Q

Cations

A

Positive ions - lose electrons

22
Q

Anions

A

Negative ions - gain electrons

23
Q

S-block

A

Highest energy electrons in the s-sub-shell (block of two groups)

24
Q

P-block

A

Highest energy electrons in the p-sub-shell (right block of six groups)

25
Q

D-block

A

Highest energy electrons in the d-sub-shell (centre block of 10 groups)

26
Q

What happens to electron configuration with ions?

A

Highest energy sun-shells lose or gain electrons (4s loses before 3d)

27
Q

How to form electron configuration of Ni2+ ; Ni has 28 electrons

A

4s FILLS FIRST
Once full, 3d energy falls below that of 4s therefore 4s loses first as well
Nickel atom : 1s2 2s2 2p6 3s2 3p6 3d8 4s2
Ni2+ atom : 1s2 2s2 2p6 3s2 3p6 3d8

28
Q

2 exceptions to electron configuration?

A

Chromium and Copper ; 3d does want to be half or fully full therefore 1 electron is taken from 4s2

29
Q

What is ionic bonding?

A

Electrostatic attraction between positive and negative ions (holding together cations and anions in ionic compounds)

30
Q

Structure of ionic bonding?

A

Each ion attracts oppositely charged ions in all directions - resulting in a giant (repeating) ionic lattice

31
Q

Features of giant ionic lattice?

A

Each metal ion is surrounded by 6 non metal ions and vice versa - this attraction between oppositely charged ions forms a giant ionic lattice

32
Q

What state are ionic compounds at room temperature?

A

Solids (high MP and BP) - at room temp there is insufficient energy to overcome strong electrostatic forces and therefore high temperatures are needed to provide the large quantity of energy needed to overcome these forces (therefore high MP and BP)

33
Q

What else affects MP and BP?

A

They are higher temperatures for lattices containing ions with greater ionic charged as there is a stronger attraction between the ions - this depends on the size of the ions

34
Q

Solubility of ions?

A

Many ions dissolve in polar solvents like water because the H2O molecules break down the lattice and surround each ion in solution

35
Q

What affects solubility of ions?

A

Large charges, which mean there is a larger ionic attraction, means that water is unable to break down the lattice structure - therefore the compound is not that soluble

36
Q

Why can ionic compounds not conduct electricity in the solid state?

A

They are in a fixed position in the giant ionic lattice and there are no mobile charge carriers

37
Q

Why can they conduct electricity when dissolved/liquid?

A

Ionic lattice breaks down with the ions now free to move - now it is a conductor of electricity in liquid and aqueous states

38
Q

What is covalent bonding?

A

Strong electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atoms

39
Q

How are they bonded covalently?

A

A small molecule, giant covalent structure or a charged poly atomic ion

40
Q

What is the orbital overlap in covalent bonding?

A

Overlap of atomic orbitals - each contains 1 electron to give a shared pair of electrons
The shared pair of electrons is attracted to the nuclei of both the bonding atoms

41
Q

What does it mean by the attraction being localised?

A

Unlike giant ionic lattices where ions attract 6 other oppositely charged ions, here the attraction is localised within the atom itself, creating a molecule - which is the smallest part of a covalent compound that can exist whilst retaining the chemical properties

42
Q

Displayed formula?

A

Relative positioning of atoms and bonds between them shown as lines - pair electrons that are not shared are lone pairs

43
Q

Number of covalent bonds usually formed by carbon, nitrogen, oxygen and hydrogen?

A 
C = 4
N = 3
O = 2
H = 1
44
Q

Expansion of the octet?

A

The outer shell of sulfur in SF6 now has 12 electrons - which is on,y possible from the n=3 shell when a d-sub shell becomes available for expansion

45
Q

What is a dative bond?

A

Covalent bond where the shared pair of electrons has been supplied by only one of the bonding atoms - this was originally a lone pair of electrons

46
Q

Draw NH4+

A

H+ means there are no electrons in hydrogen therefore nitrogen (5 in outer shell) must donate 2 of its own to hydrogen (arrow going towards hydrogen) ; they are all equivalent in magnitude and you cannot tell which is the dative bond

47
Q

Average bond enthalpy?

A

Larger enthalpy = stronger covalent bond (depends on difference in electronegativity)

48
Q

What does electron configuration show you?

A

Number of last subshell (BEFORE the letter) shows period (starts with 1 at helium/hydrogen)
Last number tells you which group IN that block
Last letter shows which block it is in

a-level chemistry (28 decks)

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