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Chemistry > 2.2 > Flashcards

2.2 Flashcards

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1
Q

What is an atomic orbital?

A

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

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2
Q

What is the shape of s-orbitals?

A

spherical

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3
Q

How many s-orbitals does each shell contain?

A

1

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4
Q

What is the shape of p-orbitals?

A

Dumb-bell

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5
Q

How many p-orbitals does each shell contain from n=2?

A

3

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6
Q

How many d-orbitals does each shell contain from n=3?

A

5

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7
Q

How many f-orbitals does each shell contain from n=4?

A

7

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8
Q

What is a sub-shell?

A

A group of the same type of orbitals within a shell (e.g. 2p)

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9
Q

What is the order of orbital filling?

A

1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 …

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10
Q

How do orbitals fill?

A

In order of increasing energy

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11
Q

Why does the 4s sub-shell fill before the 3d sub-shell?

A

The 4s sub-shell has a lower energy level

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12
Q

Why do the 2 electrons in an orbital have to have opposite spins?

A

Opposite spins help to counteract the repulsion between the negative charges.

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13
Q

How do electrons fill an orbital?

A

Orbitals with the same energy are occupied singly until there are no further orbitals available at the same energy level

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14
Q

How can electron configurations be written in shorthand?

A

[Previous noble gas] plus additional outer electron sub-shells.

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15
Q

Where is the s-block on the periodic table?

A

left block of two groups

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16
Q

Where is the p-block on the periodic table?

A

right block of six groups

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17
Q

Where is the d-block on the periodic table?

A

centre block of ten groups

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18
Q

Why does the 4s sub-shell empty before the 3d sub-shell?

A

Once filled, 3d energy level is lower than 4s .

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19
Q

What is different about the electron configuration of Cr?

A

4s1 3d5
it is more stable to have a half full d-subshell than a full s-subshell.

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20
Q

What is different about the electron configuration of Cu?

A

4s1 3d10
it is more stable to have a full d-subshell than a full s-subshell.

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21
Q

What is ionic bonding?

A

The electrostatic attraction between positive and negative ions.

22
Q

What is the structure of ionic compunds?

A

Giant ionic lattice

23
Q

Properties of ionic compounds:

A
  • Solid at room temp
  • High melting and boiling points
  • Dissolve in polar solvents
  • Conducts electricity when molten or dissolved.
24
Q

What is covalent bonding?

A

The strong electrostatic attraction between a shared pair of electrons and the nuclei of bonded atoms.

25
Q

What are the 3 ways atoms can be bonded covalently?

A

small molecule, giant covalent structure, charged polyatomic ion.

26
Q

What does a localised attraction mean?

A

Attraction acts solely between the shared pair of electrons and the nuclei of the two bonded atoms.

27
Q

What is a molecule?

A

The smallest part of a covalent compound that can exist whilst retaining the chemical properties of the compound.

28
Q

When does expansion of the ocet occur?

A

When an atom is able to have more than 8 valence electrons.
e.g. in SF6, all 6 S atoms pair to form 6 covalent bonds.

29
Q

What are multiple covalent bonds?

A

Two atoms share more than one pair of electrons

30
Q

What are dative covalent bonds?

A

The shared pair of electrons has been supplied by only one of the bonding atoms. The shared pair was originally a lone pair.

31
Q

What is average bond enthalpy?

A

A measure of covalent bond strength.

32
Q

What is electron-pair repulsion theory?

A

Electron pairs surrounding a central atom determine the shape of the molecule or ion. Arrangement of electron pairs minimises repulsion, holding the bonded atoms in a definite shape.

33
Q

Why do lone pairs reduce the bond angle?

A

They repel more strongly.

34
Q

3 BP, 0LP name and angle

A

Trigonal planar, 120

35
Q

4 BP, 0LP name and angle

A

Tetrahedral, 109.5

36
Q

6 BP, 0LP name and angle

A

Octahedral, 90

37
Q

2 BP, 0LP name and angle

A

Linear, 180

38
Q

3 BP, 1LP name and angle

A

Pyramidal, 107

39
Q

2 BP, 2LP name and angle

A

Non-linear, 104.5

40
Q

What is electronegativity?

A

The ability of an atom to attract the bonding electrons in a covalent bond.

41
Q

What is the range of electronegativity difference to be classified as a polar covalent bond?

A

0 - 1.8

42
Q

When do non-polar covalent bonds occur?

A

Between two identical atoms as the electrons are shared equally.
Or if the atoms have similar electro-negativity (C-H)

43
Q

When do polar covalent bonds occur?

A

Atoms bonded together are different, one of the atoms attracts the electrons more, causing a permanent dipole.

44
Q

When is a molecule symmetrical?

A
  • No lone pairs, only bonding pairs
  • All the bonds are the same
  • Dipoles cancel out so the molecule is non-polar.
45
Q

What are inter molecular forces?

A

Weak interactions between dipoles of different molecules.

46
Q

What are the 3 main categories of inter-molecular forces?

A
  • Induced dipole - dipole interactions / London forces.
  • Permanent dipole - dipole interactions.
  • Hydrogen bonding.
47
Q

How do London forces form?

A
  1. Movement of electrons produces a changing dipole.
  2. Instantaneous dipole forms
    and induces a dipole on a neighbouring molecule.
  3. Induced dipoles attract one another.
48
Q

How can you increase the strength of London forces?

A

More electrons in each molecule

49
Q

When do hydrogen bonds form?

A

A NOF atom (electro-negative) bonded to a hydrogen atom (electro-positive)

50
Q

What are 3 anomalous properties of water?

A
  • Solid is less dense than liquid
  • Relatively high melting and boiling points
  • Relatively high surface tension and viscosity.

Chemistry (4 decks)

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