Shapes of molecules and intermolecular forces Flashcards

1
Q

What is electron pair repulsion theory?

A

Electron pairs repel each other and take up space around the molecule as far apart as possible.
Lone pairs repel more than bonding pairs.
Bond angles between bonding pairs are often reduced because they are pushed together by lone pair repulsion.

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2
Q

What is linear shape?

A

Molecules with two electron pairs have a bond angle of 180° and is linear.

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3
Q

What is trigonal planar shape?

A

Molecules with 3 electron pairs around the central atom have a 120° bond angle, and is trigonal planar.

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4
Q

What is non-linear?

A

If there are two bonding pairs and one lone pair, the shape is non-linear.
The bond angle will be 117.5°.

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5
Q

What is tetrahedral shape?

A

If there are 4 bonding electrons, all the bond angles are 109.5°, and is tetrahedral.

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6
Q

What is pyramidal?

A

If there are 3 bonding pairs and one lone pair, the bond angle is 107°, and trigonal pyramidal.

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7
Q

What is trigonal bipyramidal?

A

Some central atoms expand the octet, and so have more than 8 bonding electrons.
5 bonding pairs is trigonal bipyramidal, and 3 atoms have 120°, and 2 have 90° between them.

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8
Q

What is octahedral?

A

A molecule with 6 bonding pairs are 90° apart and octahedral.

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9
Q

What is electronegativity?

A

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

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10
Q

What is the Pauling Scale?

A

The measure of electronegativity.
A higher number means an element is better able to attract the bonding electrons.

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11
Q

What is the trend of electronegativity?

A

Increases across periods and decreases down groups.
Fluorine is the most electronegative.
Oxygen, nitrogen and chlorine are also very strongly electronegative.

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12
Q

Why does electronegativity increase across the periods?

A

The number of protons (nuclear charge) increases so attraction between the nucleus and shared electrons increases.
The atomic radius decreases.
Shielding stays the same.
So electronegativity increases.

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13
Q

Why does electronegativity decrease down the groups?

A

Atomic radius increases.
Shielding increases.
The nuclear charge increases but the effect is shadowed.
So electronegativity decreases.

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14
Q

What is shielding of electrons?

A

The nuclear charge is shielded by electrons in the inner shells.
This weakens attraction.

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15
Q

What are polar bonds?

A

The electrons are shared unequally.
An electronegative atom is bonded to a less electronegative atom.

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16
Q

What are non-polar bonds?

A

Electrons are shared equally.
The covalent bonds in diatomic gases are non-polar as the atoms have equal electronegativities so the electrons are equally attracted to both nuclei.
E.g. H, Cl ,O
H - C bonds have very similar electronegativities so are non-polar.

17
Q

What are permanent dipoles?

A

In a polar bond, the difference in electronegativity between the two atoms causes a permanent 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.

18
Q

What are polar molecules?

A

Polar molecules have an overall dipole.
In simple molecules, the one polar bond causes a single permanent dipole, which gives the molecule an overall dipole.

19
Q

What are non-polar molecules?

A

If the polar bonds are arranged symmetrically, so that the dipoles cancel each other out e.g. in CO2, the molecule has no overall dipole and is non-polar.

20
Q

How can bond types be predicted?

A

The greater the difference in electronegativities between the two elements the less covalent and more ionic the bond will be.
Bonds where the electronegativity difference is less than 0.4 will be non-polar covalent bonds.
The electronegativity difference within 0.4-2.0 means the bonds will be polar but mainly covalent.
A difference more than 2 means the bonding is mainly ionic.

21
Q

What are intermolecular forces?

A

Forces between molecules, they’re weaker than covalent, ionic or metallic bonds.
Are induced dipole-dipole interactions, permanent dipole-dipole interactions, and hydrogen bonding.

22
Q

What are induced dipole-dipole interactions?

A

London forces are weak intermolecular forces between all molecules:
Movement of electrons produces a changing dipole in a molecule.
At any time, an instantaneous dipole will exist, but its position is constantly shifting.
This instantaneous dipole induces a dipole on a neighbouring molecule.
The induced dipole induces further dipoles on neighbouring molecules, which then attract one another.

23
Q

How does the strength of induced dipole interactions vary?

A

The greater the number of electrons, the stronger the attraction, and so the greater the energy needed to separate the particles.
Molecules with a greater surface area have more strength, as they have a more exposed electron cloud.

24
Q

What are permanent dipole-dipole interactions?

A

The small S+ and S- charges on polar molecules cause weak electrostatic forces of attraction between molecules.
They happen in addition to induced dipole interactions.

25
Q

What is the strength of permanent dipole interactions?

A

They are stronger than induced interactions, so require more energy to overcome - higher boiling and melting points.

26
Q

What is hydrogen bonding?

A

When hydrogen is covalently bonded to fluorine, nitrogen or oxygen.
Hydrogen has a high charge density due to being small, and F, N and O are very electronegative.
The bond is polarised so a weak bond forms between the hydrogen of one molecule and a lone pair on the other element.

27
Q

What effect does hydrogen bonding have on substances?

A

They are soluble in water and have higher boiling and melting points than molecules of the same size that do not form hydrogen bonds.
Water and Ammonia have very high boiling points when comparing them with other hydrides.

28
Q

How does hydrogen bonding affect ice?

A

In ice, H2O molecules are held together in a lattice by hydrogen bonds.
When ice melts, hydrogen bonds are broken, so ice has more hydrogen bonds than liquid water.
Hydrogen bonds are long, so ice molecules will be further apart than in liquid water.
This makes ice less dense than water.

29
Q

What are the anomalous propertys of ice?

A

Ice has a relatively high melting point - because hydrogen bonds are relatively strong (stronger than London forces).
Ice is less dense than water - it has an open structure / the molecules in ice are held apart by hydrogen bonds.

30
Q

What is metallic bonding?

A

The electrostatic attraction between positive ions and delocalised electrons.

31
Q

What is the solubility of simple covalent compounds?

A

Water is a polar molecule, so only tends to dissolve other polar substances well.
Compounds with hydrogen bonds, e.g. ammonia, can form hydrogen bonds with water molecules, so are soluble.
Non-polar molecules, e.g. methane, are insoluble.

32
Q

What are simple molecular lattices?

A

Simple molecules - small units containing a definite number of atoms with a definite molecular formula - such as Ne, H2, H2O, CO2.
The molecules are held in place by weak intermolecular forces.
The atoms within each molecule are bonded together strongly by covalent bonds.

33
Q

What are the melting and boiling points of simple molecular lattices?

A

The weak intermolecular forces can be broken easily, so have low melting and boiling points.
When it is broken apart during melting, only the weak intermolecular forces break, the covalent bonds are strong and do not break.

34
Q

How do you identify the strongest intermolecular force?

A

Is there a polar bond?
- no - induced dipole interactions
- yes - do the dipoles cancel/ symmetrical?
- yes - induced dipole interactions
- no - is there a H directly bonded to N, O, or F?
- yes - hydrogen bonding
- no - permanent dipole-dipole interactions.