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Chemistry > Bonding A1 > Flashcards

Bonding A1 Flashcards

Module 2

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1
Q

What is the definition of a metallic bond?

A

The strong electrostatic attraction between cations and delocalised electrons.

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2
Q

How does metal ion charge and ionic radius affect the strength of metallic bonding?

A

The higher the charge, the stronger the bond.
The larger the ionic radius, the weaker the bond.

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3
Q

Describe the structure of metals.

A

Cations are fixed in position - maintaining structure and shape of the metal.
Each atom donates its outer shell of electrons to a shared pool of electrons which are delocalised.
Delocalised electrons are mobile so move throughout structure (can carry charge).

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4
Q

What is the definition of an ionic bond?

A

The strong electrostatic attraction between positive and negative ions. (between metal and non-metal)

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5
Q

How does ion charge and ionic radius affect the strength of ionic bonding?

How does strength of ionic bonding affect the m.p/b.p of an ionic compound?

A

The higher the charge, the stronger the bond.
The larger the ionic radius, the weaker the bond (because the charge is spread over a large surface area).

The stronger the bond, the higher the mp/b.p of an ionic compound.

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6
Q

What do ionic compounds form?

A

Giant Ionic Lattices
- each ion is surrounded by oppositely charged ions.

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7
Q

What is the definition of a covalent bond?

A

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

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8
Q

What is a lone pair of electrons?

A

An electron pair not involved in bonding.

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9
Q

What is the average bond enthalpy?

A

A measurement of bond strength.

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10
Q

In a dot and cross diagram, does there have to be 8 electrons in the outer shell? (Octet Rule)

A

No there are exceptions - it can be more or less.

Elements in period 3 such as P, S and Cl can hold up to 18 electrons in their outer shell so they can form more bonds.

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11
Q

What happens in a dative covalent compound?

A

One element gives both electrons to the bond.

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12
Q

What does the electron repulsion theory state?

A

Electron pairs repel each other to be as far apart as possible.

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13
Q

Put in order the amount of repulsion in terms of electron pairs from greatest to least.

A

2 lone pairs (repel most)
1 lone pair, one bonded pair
2 bonded pairs (repel least)

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14
Q

What is electronegativity?

A

The ability of an atom to attract a bonding pair of electrons.

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15
Q

What happens to electronegativity as you go down the group? Why?

A

It decreases because:
- more shells, so more shielding
- atomic radius increases (outer shell further from the nucleus)
- so valence electrons less attracted to the nucleus.

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16
Q

What happens to electronegativity as you go across a period? Why?

A

It increases because:
- Nuclear charge increases
- atomic radius decreases
- so valence electrons more attracted to the nucleus.

17
Q

What are the 3 types of bonds which come from electronegativity differences?

A
  • Non-polar Covalent (electrons shared equally)
  • Polar Covalent (electrons shared unequally)
  • Ionic (metal gives an electron to a non-metal)
18
Q

What is a dipole?

A

Separation of charges

19
Q

Why are not all molecules with a dipole polar?

A

Because the molecule is symmetrical so the dipoles cancel out.

20
Q

What are intermolecular forces? Are they stronger or weaker than chemical bonds?

A

Attractions between molecules. Much weaker than chemical bonds.

21
Q

What are the 3 types of intermolecular forces? Describe them.

A

London (induced dipole-dipole) forces
- Weak attractions that form between all molecules due to uneven distribution of electrons.

Permanent dipole-dipole forces
- Stronger then London forces
- Exists for polar molecules
- Attractions between partially positive and partially negative ends of molecules.

Hydrogen bonding
-Strongest type of intermolecular force
- A special type of permanent dipole-dipole force when there is a O-H, N-H or F-H bond.
- You need a lone pair of electrons for Hydrogen bonds to form.

22
Q

Describe how London forces form.

A

1) Uneven distribution of electrons due to chance causes a molecule/atom to have an instantaneous dipole.
2) This induces a dipole in neighbouring molecules/atoms.
3) The positive and negative dipoles from each molecule/atom attract one another, forming a London force.

23
Q

How does the size of an atom affect strength of London forces? How does this affect m.p/b.p?

A

The larger an atom, the more electrons, so stronger induced dipole forces.
The stronger these forces, the higher the m.p/b.p.

24
Q

What do you need to include in your drawing of hydrogen bonding?

A
  • Labelled hydrogen bond (dashed line)
  • Lone pair of electrons
  • Show dipoles on atoms involved (Delta plus and minus)
25
Q

Why is water unusual?

A

It has an unusually high m.p/b.p for a molecule with such a low molar mass.

26
Q

Why is solid water (ice) less dense than liquid water?

A

Liquid water: hydrogen bonds constantly broken and reformed.
Solid water: molecules arranged in hexagonal rings to maximise the number of stable hydrogen bonds - not enough kinetic energy to break bonds.

27
Q

What is the name of the shape of a molecule with 4 bonding pairs and 0 lone pairs of electrons? What is the angle?

A

Tetrahedral - 109.5 degrees

28
Q

What is the name of the shape of a molecule with 3 bonding pairs and 1 lone pairs of electrons? What is the angle?

A

Trigonal Pyramidal - 107 degrees

29
Q

What is the name of the shape of a molecule with 2 bonding pairs and 2 lone pairs of electrons? What is the angle?

A

Non-linear - 104.5 degrees

30
Q

What is the name of the shape of a molecule with 6 bonding pairs and 0 lone pairs of electrons? What is the angle?

A

Octahedral - 90 degrees

31
Q

What is the name of the shape of a molecule with 2 bonding pairs and 0 lone pairs of electrons? What is the angle?

A

Linear - 180 degrees

32
Q

What is the name of the shape of a molecule with 3 bonding pairs and 0 lone pairs of electrons? What is the angle?

A

Trigonal Planar - 120 degrees

33
Q

Which molecule shapes can be symetrical and therefore non-polar?

A

Linear, Trignal Planar, Tetrahedral and Octahedral

34
Q

Which elements form giant metallic lattices?

A

Li, Be, Na, Mg, Al

35
Q

Which elements form giant covalent structures?

A

B, C, Si

36
Q

What elements form simple molecular structures? What are the formulas?

A

N2, P4, O2, S8, F2, Cl2, Ne, Ar

37
Q

What is the difference between Graphite, Graphene and Diamond?

A

ALL made from C atoms and are giant covalent lattices. Allotropes (differernt forms of same element in the same state).
All have v high m.p. and insoluble due to strong covalent bonds.

GRAPHITE
- lots of layers: weak intermolecular forces (London) between layers, strong covalent bonds within layers - 3D
- C atoms arranged in hexagons - 3 covalent bonds
- delocalised electrons (1 electron per C atom)
- lubricant, lightweight, insoluble, conducts electricity…

GRAPHENE
- one sheet of C atoms in hexagons - 2D
- similar structure to graphite with bonding.
- lightweight, transparent, conducts electricity

DIAMOND
- covalently bonded to 4 other C atoms - tetrahedral.
- v hard, good thermal conductor (vibrations travel easily)
- Cannot conduct electricity
(Silicon is also C covalently bonded to 4 other C atoms)

38
Q

When can ionic lattices conduct electricity?

A

When dissolved in liquid as ions are free to move.

39
Q

Are metallic lattices soluble?

A

No - because strong metallic bonds.

Chemistry (22 decks)

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