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CHEMISTRY UNIT 1 > Electrons, Bonding and Structure > Flashcards

Electrons, Bonding and Structure Flashcards

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1
Q

How many orbitals and (max) electrons are there in the s sub-shell?

A

1 orbital, 2 electrons

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2
Q

How many orbitals and (max) electrons are there in the p sub-shell?

A

3 orbitals, 6 electrons

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3
Q

How many orbitals and (max) electrons are there in the d sub-shell?

A

5 orbitals, 10 electrons

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4
Q

How many orbitals and (max) electrons are there in the f sub-shell?

A

7 orbitals, 14 electrons

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5
Q

What are sub-shells of the 1st shell?

A

1s

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6
Q

What are sub-shells of the 2nd shell?

A

2s 2p

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7
Q

What are sub-shells of the 3rd shell?

A

3s 3p 3d

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8
Q

What are sub-shells of the 4th shell?

A

4s 4p 4d 4f

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9
Q

What is an orbital?

A

95% of the time (or any other percentage you choose), the electron will be found within a fairly easily defined region of space quite close to the nucleus. Such a region of space is called an orbital.

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10
Q

What are the shapes of s and p orbitals?

A

S are spherical and P are dumbbell shaped (3p orbitals are at right angle to each other).

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11
Q

What is the definition of the first ionisation energy?

A

It is the energy needed to remove 1 electron from each atom in 1 mole of gaseous atoms to form 1 mole of gaseous 1+ ions.

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12
Q

What is the equation for the first ionisation?

A

X(g) -> X+(g) + e-

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13
Q

What factors affect ionisation energy?

A

Nuclear charge, the distance from the nucleus and shielding.

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14
Q

What is meant by nuclear charge?

A

The more protons there are in the nucleus, the more positively charged the nucleus is, and the more strongly electrons are attracted to it.

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15
Q

How does distance from the nucleus affect ionisation energies?

A

Attraction falls off very rapidly with distance. An electron close to the nucleus will be much more strongly attracted than one further away.

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16
Q

What is shielding?

A

As the number of electrons between the outer electrons and the nucleus increases, the outer electrons feel less nuclear charge attraction. This lessening of pull of the nucleus by inner shells of electrons is called shielding.

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17
Q

What is the definition of second ionisation energy?

A

It is the energy required to remove 1 electron from each ion in 1 mole of gaseous 1 + ions to form 1 mole of gaseous 2+ ions.

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18
Q

What is the equation for second ionisation energy?

A

X+ (g) -> X2+(g) + e-

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19
Q

Why does successive ionisation energy increase within each new shell?

A

Electrons are being removed from an increasingly positive ions - there’s less repulsion amongst the remaining electrons, so they’re held more strongly by the nucleus.

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20
Q

What is an ionic bond?

A

It’s an electrostatic attraction between two oppositely charged ions.

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21
Q

What are the formulas for the following ionic compounds? Nitrate, Carbonate, Sulfate, Ammonium?

A

NO3-, CO3-, SO4^2-, NH4+

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22
Q

Describe the structure of a giant ionic lattice.

A

Ions are packed together in a regular structure (lattice) that is made up of the same basic unit repeated over and over again.

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23
Q

What are the physical properties of ionic structures?

A

Conduct elect. when molten or dissolved as ions in liquid free to move and carry a charge. Have high melting points - held by strong electrostatic forces. Tend to dissolve as they are polar.

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24
Q

Why can’t ionic structures conduct electricity when solid?

A

The ions are held in fixed positions by the strong ionic bonds and arent free to move and carry a charge.

25
Q

What is a covalent bond?

A

A shared pair of electrons.

26
Q

What are some covalent bonds which are special cases?

A

Boron trifluoride has 6 ekectrons in it’s outer shell. Sulfur hexafluoride, sulfur has 12 electrons in its out shell.

27
Q

What is a dative covalent bond?

A

When both electrons in a pair come from one atom.

28
Q

What are some examples of giant covalent lattices?

A

Diamond and graphite.

29
Q

What are giant covalent lattices?

A

Huge networks of covalently bonded atoms. Carbon atoms can form this type of structure because they can each form four strong, covalent bonds.

30
Q

Diamond is the hardest known substance and is made up of carbon atoms. How many covalent bonds do each carbon atom form? What shape do they form?

A

4 - Each carbon atom is covalently bonded to four other carbon atoms. They arrange themselves in a tetrahedral shape - its crystal lattice structure.

31
Q

What is meant by ‘sublime’?

A

Changes straight from a solid to a gas, skipping out the liquid stage.

32
Q

Diamond has a number of properties due to the strong covalent bonds. List them.

A
  • Very high melting point and extremely hard - Good thermal conductor - Can’t conduct electricity (all outer electrons are held in localised bonds) - Won’t dissolve in any solvent
33
Q

Graphite is an allotrope of carbon. What is an allotrope?

A

Different forms of the same element in the same state.

34
Q

Describe the structure of graphite.

A

Carbon atoms arranged in sheets of slat hexagons covalently bonded with 3 bonds each. 4th outer electron of each carbon atom is delocalised between the shets of hexagons. The sheets of hexagon are bond together by weak VDW forces.

35
Q

Describe the properties of graphite.

A

Slippery - weak bonds between layers easily broken (So sheets can slide) - delocalised electrons free to flow and conduct - high melting point - insoluable in any solvent

36
Q

What do metal elements exist as?

A

Giant metallic lattice structures.=

37
Q

Describe the structure of a giant metallic lattice.

A

Electrons in the outermost shell of a metal atom are delocalised and are free to move. This leaves a positive metal ion. The positive metal ions are attracted to the delocalised negative electrons. They form a lattice of closely packed positive ions in a sea of delocalised electrons - this is metallic bonding.

38
Q

Draw a giant metallic lattice structure.

A
39
Q

Draw graphites structure.

A
40
Q

Explain how metallic bonding affects the properites of a metal.

A
  • The more delocalised electrons there are, the higher the melting point.
  • As there are no bonds holding specific ions together, the metal ions can slide past each other when the structure is pulled, so metals are malleable and ductile.
  • The delocalised electrons can pass KE to each other, making them good thermal conductors.
  • Metals are good electrical conductors because the delocalised electrons can carry a current.
  • Metals are insoluable because of the strength of metallic bonding.
41
Q

What is the order of repulsion?

A
  1. Lone-pair/ Lone-pair bond angles are biggest
  2. Lone-pair/Lone-pair bond angles are second biggest
  3. Bonding-pair/Bonding-pair angles are smallest.
42
Q

How many degrees does a lone pair repel?

A

2.5º

43
Q

What are the shapes of molecules with 2 electron pairs on a central atom (e.g. CO2 or H2O)?

A

Linear molecules (180°)

44
Q

What is the shape of a molecule which has 3 electron pairs on a central atom (no lone pairs)?

A

Trigonal planar (120°)

45
Q

What is the shape of a molecule which has 4 electron pairs on a central atom (no lone pairs)?

A

Tetrahedral (109.5°)

46
Q

What is the shape of a molecule which has 4 electron pairs on a central atom (1 lone pair)?

A

Trigonal Pyramidal (107º)

47
Q

What is the shape of a molecule which has 4 electron pairs on a central atom (2 lone pairs)?

A

Non-linear or ‘bent’ 104.5º

48
Q

What is the shape of a molecule which has 6 electron pairs on a central atom?

A

Octahedral (90º)

49
Q

What is electronegativity?

A

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

50
Q

What causes a bond to be polar?

A

In a covalent bond between two atoms of different electronegativies, the bonding electrons are pulled towards the electronegative atom.

51
Q

What causes a bond to be non-polar?

A

The atoms have equal electronegativities and so electrons are equally attracted to both nuclei. Some elements, like carbon and hydrogen, have pretty similar electronegativities, so bonds between them are essentially non-polar.

52
Q
  1. What causes a permanent dipole?
A
  1. In a polar bond, the differences in the atoms’ electronegativities causes a dipole. A dipole is a difference in charge between the two atoms caused by a shift in electron density in the bond. The greater the difference in electronegativity, the more polar the bond.
53
Q

What is a permanent dipole-dipole interaction?

A

The partially positive and negative charges on polar molecules cause weak electrostatic forces of attraction between molecules.

54
Q

Out of hydrogen bonding, temporary dipole-induced dipole (VDW) and permanant dipole-dipole interactions, which is the strongest and weakest?

A
  1. Permanent dipole-dipole interaction
  2. Hydrogen bonding
  3. Temporary dipole-induced dipole (VDW)
55
Q

When can hydrogen bonding ONLY happen?

A

When hydrogen is covalently bonded to fluorine, nitrogen or oxygen. Hydrogen has a high charge density because it’s so small and fluorine, nitrogen and oxygen are very electronegative. The bond is so polarised that the hydrogen of one molecule forms a weak bond with the fluorine, nitrogen or oxygen of another molecule.

56
Q

How do Van der Waals forces arise?

A
  • Electrons in charge clouds are always moving really quickly. At any particular moment, the electrons in an atom are likely to be more one side than the other. At this moment, the atom would have a temporary dipole.
  • This dipole can cause another temporary dipole in the opposite direction on a neighbouring atom. The two dipoles are then attracted to each other.
57
Q

What factors affect Van der Waals forces?

A
  • The atomic/molecular size
  • The number of shells of electrons
58
Q

What types of bond must be overcome in order for a simple molecular substance to boil or melt?

A

VDW forces or hydrogen bonds.

CHEMISTRY UNIT 1 (2 decks)

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