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Chemistry > Bonding > Flashcards

Bonding Flashcards

1
Q

What forces hold ionic bonds together?

A

Electrostatic between oppositely charged ions

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2
Q

What is the term for the structure of ionic compounds?

A

Lattice

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3
Q

What are the properties of ionic compounds?

A

🔹High mpt and solid at room temp
(Due to very strong bonding)

🔹conduct electricity when molten or
dissolved (free e-)

🔹Brittle (lattice structure)

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4
Q

Why are ionic compounds brittle?

A

When struck hard, ions in the lattice are knocked out of regular arrangement.

This small displacement may cause contact between ions of the same charge.

These ions repel each other and the structure shatters

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5
Q

How are atoms in covalent bonds held together?

A

Electrostatic forces of attraction between the positive nuclei and the negative shared electrons

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6
Q

What are the properties of covalently bonded structures?

A 
🔹Gases or volatile liquids at room 
 temp 
🔹Low mpt
🔹Poor conductors of electricity 
 (Molecules are a neutral charge overall so there are no charged particles to carry current)
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7
Q

What ions are formed in an ionic bond?

A 
Positive metal (lost e-) 
Negative n-m (gained e-)
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8
Q

What is the difference between a covalent bond and a dative covalent bond?

A

In a dative covalent bond, both electrons come from one atom

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9
Q

What are the similarities between covalent and co-ordinate bonds?

A

They have exactly the same length and strength once formed

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10
Q

🔹How is a covalent bond between nitrogen and hydrogen atoms represented?

🔹How is a co-ordinate bond between a nitrogen atom and a hydrogen ion represented?

A

🔹N —— H

🔹N —–> H

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11
Q

When representing a dative covalent bond, which direction is the arrow drawn?

A

Points towards the atom accepting the lone pair of e- (the electron deficient atom)

N ——> H(+)

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12
Q

Describe the structure of a metal

A

Lattice of positive metal ions in a sea of delocalised e- held together by strong electrostatic forces of attraction

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13
Q

Why are the e- in a metal delocalised?

A

The outer main levels of the atoms merge so the outer e- are no longer associated with any one particular atom

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14
Q

Why does the number of delocalised ions in different types of metals vary?

A

Metals can lose up to 3 e- in their outer shell. G3 metals have more delocalised e- than G1.

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15
Q

Why are metals good conductors of electricity?

A

Due to delocalised e-
An e- from the negative terminal of the supply joins e- sea at one end of a metal wire while at the same time a different e- leaves the wire at the positive terminal

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16
Q

Why are metals good conductors of heat? (High thermal conductivity)

A

🔹Closely packed ions = increasingly vigorous vibrations

🔹Sea of delocalised e-

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17
Q

What factors effect the strength of a metal?

A

🔹CHARGE ON THE ION
Greater charge = more free e- = stronger electrostatic attraction

🔹SIZE OF ION
Smaller ion = e- closer to nucleus = stronger bond

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18
Q

Why are metals malleable and ductile?

A

Layers slide over each other

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19
Q

What are the properties of metals?

A 
🔹Conductors of heat & electricity 
🔹Strong (usually) 
🔹Malleable (can be beaten into shape)
🔹Ductile (can be pulled into thin wires)
🔹High mpt/bpt -----> solid at room temp (except Hg)
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20
Q

Define electronegativity

A

The power of an atom to attract the electron density in a covalent bond towards itself

21
Q

What scale is used to measure electronegativity?

What range is used?

A

The Pauling scale

0-4

22
Q

Electronegativity depends on what 3 factors?

A

🔹Nuclear charge
🔹Distance between the nucleus and the outer shell electrons
🔹Shielding of the nuclear charge

23
Q

Does electronegativity increase or decrease as you move along a period?

A

Increases

24
Q

Does electronegativity increase or decrease as you move down a group?

A

Decreases

25
Q

How does nuclear charge relate to electronegativity?

A

More protons = more positive charge = stronger attraction between the nucleus and negative charge of e-

26
Q

How does the distance between the nucleus and the outer shell relate to electronegativity?

A

Smaller atom = closer nucleus is to shared e- = greater pull

27
Q

How does shielding relate to electronegativity?

A

Larger nuclear charge for a given shielding effect = greater electronegativity

28
Q

Is this bond polar or non-polar?

F - F

A

Non-polar because both atoms are the same and therefore have the same electronegativity, so the electrons in the bond feel exactly the same attraction

29
Q

Is this bond polar or non-polar?

H - F

A

Polar because H and F have different electronegativities.

30
Q

In the bond H - Cl, which atom is δ+ and which is δ-?

A 
H = δ+
Cl = δ-

Cl is more electronegative than H which means that the negative electrons in the bond will be more attracted to the Cl atom

31
Q

Arrange the following covalent bonds in order of increasing polarity

H-O, H-F, H-N

A

H-N, H-O, H,F

Because fluorine is the most electronegative

32
Q

Arrange the 3 types of intermolecular forces from the weakest to the strongest

A

🔹Van der Waals
🔹Dipole - dipole
🔹Hydrogen bonding

33
Q

What is the difference between a polar bond and a dipole moment?

A

Polar bond = polarity of a particular bond within a molecule

Dipole moment = overall polarity of molecule

34
Q

Where do D-D forces act?

Give an example

A

In molecules that have permanent dipoles.

E.g hydrogen chloride
Cl = δ-
H = δ+

H—Cl H—Cl

35
Q

What causes polarity in dipole moments?

A

Electronegativity of atoms in the molecule

36
Q

What are van der Waal’s forces?

A

Weak electrostatic forces of attraction between protons and electrons I’m different atoms/molecules

37
Q

Are van der Waal’s forces permanent or temporary?

Why?

A

Temporary

Electrons are constantly so moving so the distribution of charge is always changing, creating temporary dipoles

38
Q

How do van der Waal’s forces work?

A

The temporary dipole in one atom/molecule will affect the electron distribution in those around it.

Electron dense area will repel electrons in nearby a/m and electrons will be attracted to the positive charge in a neighbouring nucleus

39
Q

What increases the strength of van der Waal’s forces?

A

More electrons present

40
Q

Why do the bpts of the noble gases increase as you move down the group?

A

There are more electrons present so van der Waal’s forces are stronger

41
Q

How does the length of a hydrocarbon chain affect it’s bpt?

A

Longer chain = higher bpt

Since more e- are present

42
Q

Which atoms are electronegative enough to form hydrogen bonds?

A

Oxygen, nitrogen, fluorine

43
Q

Describe a hydrogen bond

A

A hydrogen atom bonds with a very electronegative atom with a lone pair of electrons

The hydrogen atom becomes very electron deficient because the electronegative atom attracts the shared pair of e- towards itself

The d+ hydrogen atom will be attracted to the lone pair of e- in a neighbouring molecule

44
Q

Why is the boiling point of HF higher than that of Xe, considering that Xe has more electrons?

A

HF is an example of hydrogen bonding, which is much stronger than the van der Waal’s forces in Xeon

45
Q

What is stronger, a hydrogen bond or a covalent bond?

A

Covalent is about 10x stronger

46
Q

Place the following elements in order of the strength of van der Waal’s forces between the atoms (weakest first)
Ar, He, Kr, Xe

A

He, Ar, Ne, Kr

47
Q

Which one of these cannot have d-d forces acting between them?
H2O, HCl, H2

A

H2 - no atom is more electronegative than the other so there is no dipole moment

48
Q

Why is hexane a liquid at room temp whereas butane is a gas?

A

Hexane has a longer chain length
Longer chain = more electrons
More e- = stronger vdW forces = higher bpt

49
Q

What is hydrogen bonding used in?

A

🔹structure and density of ice (less sense when solid)
🔹ironing
🔹DNA structure

Chemistry (22 decks)

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