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Chemistry OCR AS Level > Electrons, Bonding and Structure > Flashcards

Electrons, Bonding and Structure Flashcards

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1
Q

What is the principal quantum number?

A
  • The principal quantum number, n, indicates the electron shell/energy level that the electrons occupy.
  • Different shells have different principal quantum numbers.
  • The larger the value of n, the further the shell is from the nucleus and the higher the energy level.
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2
Q

The first four shells hold a different number of electrons - what are they?

A
  • First Shell - 2
  • Second Shell - 8
  • Third Shell - 18
  • Fourth Shell - 32

Each shell holds up to 2n² electrons, where n is the prinicipal quantum number

I dont understand why its not 8 rather than 18, because the outershell should have 8 electrons right?. Elements react to ge 8 electrons in outer shell.

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3
Q

Define electron shell?

A

An electron shell is a group of atomic orbitals with the same principal quantum number, n.

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4
Q

In the past, scientists believed that electrons orbit the nucleus, but electrons are not solid particles and thus this planetary model has been replaced. What are electrons now believed to exist in?

A

Electrons exist in atomic orbitals.
Each electron shell contains a group r of atomic orbitals. Each atomic orbital can hold a maximum of two electrons with opposite spins,

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5
Q

Define orbital.

A

An orbital is a region of space where electrons may be found.

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6
Q

What are the different types of orbitals?

A

There are four different types of orbitals - s orbital, p orbital, d orbital, f orbital and each has a different shape (you don’t need to know about the f orbital in detail)

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7
Q

Describe (and draw) the shape of an s-orbital.

A

An s-orbital has a spherical shape.

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8
Q

How many s-orbitals are in each shell and what is the max number of s-orbital electrons found in each shell?

A

From the first shell (n = 1) upwards, each shell contains one s-orbital.
This means that there is a maximum of 2 s-orbital electrons in each shell from the first shell upwards.

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9
Q

Describe (and draw) the shape of a p-orbital.

A

A p-orbital has a 3D dumb-bell shape.

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10
Q

How many p-orbitals are in each shell and what is the max number of p-orbital electrons found in each shell?

A

From the second shell (n = 2) upwards, each shell contains three p-orbitals (px, py, pz) at right angles to one another (different planes)
This means that there are a maximum of 6 p-orbital electrons in each shell from the second shell upwards.

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11
Q

Structures for d-orbital and f-orbital are complex - do not need to learn them.
How many d-orbitals are in each shell, and what is the max number of d-orbital electrons found in each shell?

A

From the third shell (n = 3) upwards, each shell contains five d-orbitals.
This means that there is a maximum of 10 d-orbital electrons in each shell from the third shell upwards

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12
Q

How many f-orbitals are in each shell, and what is the max number of f-orbital electrons found in each shell?

A

From the fourth shell (n = 4) upwards, each shell contains seven f-orbitals.
This means that there is a maximum number of 14 f-orbital electrons in each shell from the fourth shell upwards

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13
Q

Summary: State the max number of each type of orbital in the fourth shell, and the maximum number of electrons in each type of orbital.

A

1 s-orbital - 2 s-orbital electrons
3 p-orbitals - 6 p-orbital electrons
5 d-orbitals - 10 d-orbital electrons
7 f-orbitals - 14 f-orbital electrons

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14
Q

How are electrons in orbitals represented?

A

Using box diagrams where each box represents a single orbital that can hold a maximum of two electrons.

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15
Q

Why do the two electrons in an orbital not repel each other?

A
  • Whilst both are negatively charged, electrons do not repel each other, because they have an additional property called spin.
  • The two electrons in an orbital must have opposite spins and we represent that using arrows, either ‘up’ or ‘down’.
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16
Q

What are sub-shells?

A
  • An electron shell/energy level is a group of atomic orbitals with the same principal quantum number.
  • Within each shell, orbitals of the same type can be further categorised into sub-shells.
  • There are s, p, d and f sub-shells.
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17
Q

What are the sub shells found at n = 1?
Draw the orbitals (represented as squares) at n = 1?
What are the max number of electrons in the first shell?

A

Sub-shell: 1s
Orbital: ☐
Electrons: 2 (Max number of electrons at n = 1 is 2)

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18
Q

What are the sub shells found at n = 2?
Draw the orbitals (represented as squares) at n = 2?
What are the max number of electrons in the second shell?

A

Sub-shell: 2s 2p
Orbital: ☐ ☐☐☐
Electrons: 2 222 (Max number of electrons at n = 2 is 8)

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19
Q

What are the sub shells found at n = 3?
Draw the orbitals (represented as squares) at n = 3?
What are the max number of electrons in the third shell?

A

Sub-shell: 3s 3p 3d
Orbital: ☐ ☐☐☐ ☐☐☐☐☐
Electrons: 2 222 22222 (This means the maximum number of electrons at n = 3 is 18)

As you can see, each shell gains a new type of sub-shell (s, then p and then d). For the fourth shell, on top of 4s, 4p, 4d, there would also be 4f. As there 7 orbitals, it will hold an additional 14 electrons.

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20
Q

The sub-shells that make up electron shells, have different energy levels themselves. What is the order of the sub-shell energy level, lowest to highest?

A

1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d

*CHECK THIS IS ALL WE NEED TO KNOW.

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21
Q

What is the energy of orbitals?

A

Orbitals in the same sub-shell are at the same energy level.

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22
Q

How do electrons fill/occupy sub-shells?

A

Electrons occupy the orbitals in sub-shells in order of increasing energy levels, filling the lowest available energy level first.
Each energy level must be full before the next, higher energy level starts to fill.
When a sub-shell is built up with electrons, each orbital is filled up singly before pairing. Each orbital containing paired electrons should have opposite spins
The 4s orbital is at a slightly lower energy level than the 3d orbitals. This means 4s will fill before 3d.

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23
Q

How do you write the electron configurations of atoms?

Write the electron configuration for oxygen?

A

They are written in the form nxʸ, where:

  • n is the shell number
  • x is the type of orbital
  • y is the number of electrons occupying the orbitals of the sub-shell.

Oxygen: 1s²2s²2p⁴ - You need to be able to deduce the electron configuration for the first 36 atoms.

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24
Q

Write the orbitals occupied by oxygen

A

1s²2s²2pₓ²2pᵧ¹2pz¹ (z should be subscript)

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25
Q

What happens when an atom is ionised?

A

During ionisation ( in this case, when an atom loses an electron), the electron in the highest energy level would be removed first

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26
Q

Show the electron configuration of Li and Li⁺?

Show the electron configuration of F and F⁻?

A

Li: 1s²2s¹
Li⁺: 1s²
F: 1s²2s²2p⁵
F⁻: 1s²2s²2p⁶

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27
Q

Why do elements react and bond together?

A

Elements react in order to become more stable by obtaining a full outer shell of electrons (noble gas electron configuration).
A noble gas is energetically stable.

Do we still assume that its 8 electrons or not? because now the third shell can hold up to 18 and the fourth can hold up to 36?

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28
Q

Types of chemical bonding?

A

Ionic Bonding
Covalent Bonding
Metallic Bonding

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29
Q

How does ionic bonding occur?

A
  • Occurs between a metal and a non-metal, where electrons are transferred from the metal atom to the non-metal atoms, to form oppositely charged ions which attract. The metal ion becomes positively charged and the non-metal ion becomes negatively charged.
  • Elements involved obtain a noble gas configuration.
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30
Q

Define an ‘ionic bond’?

A

An ionic bond is the electrostatic attraction between oppositely charged ions.

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31
Q

How does covalent bonding occur?

A
  • Occurs between two non-metals, where a shared pair of negatively charged electrons are attracted to the positive nuclei of both bonded atoms.
  • Elements involved obtain a noble gas configuration.
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32
Q

Define ‘covalent bond’?

A

The electrostatic attraction between shared pairs of electrons and the nuclei of the bonding atoms.

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33
Q

What are dot-and-cross diagrams used to show?

Draw the dot and cross diagram for the ionic bonding in Na₂O:

A

Dot-and-cross diagrams are used to show the origin of electrons in chemical bonding.
Drawing should show that one electron is transferred from each of the two sodium atoms to one oxygen atom with the formation of two Na⁺ ions and one O²⁻. Both elements are left with a full outer shell by gaining or losing electrons.
When drawing ions:
- Show the outer shell of the ions - where lost you can leave the shell empty
- Show charges on the ions by adding quare brackets around the ion and placing the charge outside the bracket.

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34
Q

Write the Ionic Equations for the formation of Na₂O. Also write how each of their electron configuration changes.

A

2Na → 2Na⁺ + 2e⁻
1s²2s²2p⁶3s¹→ 1s²2s²2p⁶
O + 2e⁻ → O²⁻
1s²2s²2p⁴ →1s²2s²2p⁶

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35
Q

How do ionic compounds exist?

A

All ionic compounds exist as:

  • 3D giant ionic lattices in solid state
  • where each ion is fixed in position
  • and surrounded by oppositely charged ions.
  • The electrostatic attraction between the oppositely charged ions occurs from all directions.

(Remember this in 3D drawings - same element cannot be next to each other)

When in liquid state, the ions are free to move around.

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36
Q

Examples of other Ionic Compounds?

A 
Sodium Chloride (NaCl)
Calcium oxide (CaO)
Aluminium Fluoride (AlF₃)

DRAW THEIR DOT-AND-CROSS DIAGRAMS.

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37
Q

Ioninc Equation for the formation of AlF₃

A

Al → Al³⁺ + 3e⁻

3F + 3e⁻→ 3F⁻

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38
Q

Properties of ionic compounds due to their giant ionic lattices?

A

High melting and boiling points
Electrical conductivity when molten/in solution
Soluble in polar solvents

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39
Q

Why do giant ionic lattices have high melting and boiling points?

A

Large amounts of energy are required to break the strong electrostatic force of attractions that hold together oppositely charged ions (the ionic bonds) in the giant ionic solid lattice. For this reason, ionic compounds have high melting and boiling points.

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40
Q

Why does MgO have a higher melting than NaCl?

A

This is because the charges on the Mg²⁺ and O²⁻ are greater than those on Na⁺ and Cl⁻. The greater the charge, the stronger the ionic bonds/electrostatic force of attraction between the oppositedly ions. This means a greater amount of energy is required to break the ioninc lattice of MgO than NaCl during melting.

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41
Q

Why can’t ionic compounds conduct electricity as a solid?

A

Ionic compounds exist as:

  • 3D giant ionic lattices in solid state
  • where each ion is fixed in position unable to move
  • so cannot carry the charge.
  • Therefore, cannot conduct electricity.
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42
Q

Why can ionic cmpounds conduct electricity when molten or in solution?

A

When an ionic compound is molten or dissolved in water (in solution) the giant solid lattice breaks down and the ions are free to move, so can carry the charge. Therefore, it can conduct electricity.

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43
Q

Why does an ionic lattice dissolve in polar solvents, such as water?

A

This is because polar solvents contain polar bonds. A polar bond occurs between atoms that do not share the electrons in their covalent bond equally. This results in small charges on each of the atoms. These small charges within the polar solvent break down the giant ionic lattice by attracting the oppositely charged ions, surrounding each ion to form a solution. This disrupts the lattice and the ions are pulled out of it.

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44
Q

How to draw e.g. NaCl dissolved in e.g. water

A

Na⁺ attracts δ- charges on the O atoms of the water molecules
Cl⁻ attracts δ+ charges on the H atoms of the water moelcules
Positive ions are attracted to the lone pairs on the oxygen of the water molecules (with an δ- charge) and coordinate dative covalent bonds may form. The hydrogen of the water molecules (with an δ+ charge) form hydrogen bonds with negative ions and Cl⁻ - CHECK THIS SECOND BIT AND CLEAR UP ANY QUERIES.

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45
Q

Define single covalent bond?
Define double covalent bond?
Define triple covalent bond?

A

Where atoms are bonded by one shared pair of electrons between the nuclei. Written as H-H
Where atoms are bonded by two shared pairs of electrons between the nuclei. Written as O=O.
Where atoms are bonded by three shared pairs of electrons between the nuclei. Written as N≡N.

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46
Q

Some elements will always make the same number of covalent bonds when they bond. State the number of covalent bonds for carbon, nitrogen, oxygen and hydrogen.
*4 covalent bonds could mean, 4 single bonds, 2 double bonds, a double bond and two single bonds, a triple bond and a single bond etc.

A

Carbon will always make four covalent bonds.
Nitrogen will always make three covalent bonds.
Oxygen will always makes two covalent bonds.
Hydrogen always makes one covalent bonds.

Notice: Max number of covalent bonds that can form = number of electrons in the outer shell.

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47
Q

What is a lone pair? (Compared to a bonding pair)

A

A lone pair is an outer-shell pair of electrons that is not involved in chemical bonding.

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48
Q

What does a lone pair do?

A

A lone pair gives a concentrated region of negative charge around the atom. This can influence the chemistry of a molecle in many way.

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49
Q

What is average bond enthalpy?

A

The mean energy needed for 1 mole of a given type of gaseous bonds to undergo homolytic fission.

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50
Q

Different bonds have different amounts of energy. In general, which bonds are the strongest and weakest?

A

Triple bonds are the strongest, then double bonds and then single bonds.

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51
Q

What is a dative covalent bond/coordinate bon?

A

A dative covalent bond/coordinate bond is a bond formed from a shared pair of electrons that has been provided by one of the bonding atoms only.

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52
Q

How can a dative covalent bond be written?

A

A dative covalent bond can be written as A→B, where the direction of arrow shows the direction in which the electron pair has been donated. In the example above, A is donating a pair of electrons to B.

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53
Q

Substances with a dative covalent bonds?

A

The ammonium ion, NH₄⁺

The oxonium ion, H₃O⁺

54
Q

Draw the formation of NH₄⁺ from NH₃ and H⁺ as a dot and cross diagram. Explain it also.

Now draw the ammonium ion only as a displayed formula.

A

To form the NH₄⁺ ion, the lone pair around the nitrogen atom in the NH₃ molecule provides both of the bonding electrons when bonding with H⁺, to form a dative covalent bond.
Drawings on page 88.

55
Q

Can you tell the difference between the dative covalent bond and the other covalent bonds once formed?

A

No, you cannot tell which bond was formed from the electrons of one of the bonding atoms only.

56
Q

How are oxonium ions, H₃O⁺, formed?

Write the equation for the formation of HCl

A

When an acid is added to water, water molecules form oxonium ions, H₃O⁺. For example hydrogen chloride gas forms hydrochloric acid when added to water, which then dissociates into H₃O⁺ and Cl⁻ . In equations H₃O⁺ is often simplified to H⁺.
HCl(g) + H₂O(l) → H₃O⁺(aq) + Cl⁻(aq)

57
Q

Draw the oxonium ion as a dot-and-cross diagram and displayed formula.

A

In the oxonium ion, one of the lone pairs around the oxygen atom, in a H₂O molecule provides both the bonding electrons, when bonding with H+, to form a dative covalent bond.

58
Q

Is obeying the Octet rule always possible?

A

This is not always possible because there may not be enough electrons in the outer shell to reach an octet or more than four electrons may pair up in bonding, expanding the octet.

59
Q

When more than four electron pair up in bonding, what is it called?

A

Expansion of the octet.

60
Q

What is the octet rule?

A

The octet rule states that elements gain, lose or share electrons to attain 8 electrons in their outer shell (configuration of the nearest noble gas)

61
Q

Which elements do not have enough electrons to complete the octet?

A
  • Beryllium, Boron and Aluminium can form covalent compounds.
  • However, none of these elements have enough unpaired electrons in its outer shell to reach an octet

*unpaired refers to those the electrons that could be invovled in covalent bonding

62
Q

Describe the bonding in Boron Trifluoride, BF₃,

A
  • Boron has 3 electrons in its outer shell. Fluorine atom has 7 electrons in its outer shell
  • Boron forms three covalent bonds so each of its outershell electrons pair up/are involved in bonding, so there are now six electrons in boron’s outer shell. The central boron atom does not achieve the octet.
  • However, each of the three fluorine atoms now have eight electrons in the outer shell, attaining an octet.

DRAW IT!

63
Q

Which elements have expand the octet?

A

Elements in group 15-17 of the periodic table, after period 3.
Atoms of non-metals in group 15 (P, As) can form 3 or 5 covalent bonds, depending on how many electrons are used in bonding.
Atoms of non-metals in group 16 (S, Se, Te) can form 2, 4 or 6 covalent bonds, depending on how many electrons are used in bonding.
Atoms of non-metals in group 17 (Cl, Br, I, At) can form 1.3, 5 or 7 covalent bonds, depending on how many electrons are used in bonding.

64
Q

Describe the bonding in Sulfur Tetrafluoride, SF₆

A
  • Sulfur has 6 electrons in its outer shell. Fluorine has 7 electrons in its outer shell.
  • Sulfur forms 6 covalent bonds, so each of the outershell electrons pair up/are involved in bonding so there are now 12 electrons in its outershell. The central sulfur atom has expanded the octet.
  • Also, each of the six fluorine atoms now have eight electrons in its outer shell, attaining the octet.

DRAW IT!

65
Q

How can the octet rule be modified?

A

Better rule would be:

1) Unpaired electrons pair up
2) The maximum number of electrons that can pair up is equivalent to the number of electrons in the outer shell.

i.e. if there are 6 electrons in the outer shell, 6 bonding pairs can form as a maximum (resulting in 12 electrons in the outer shell).

66
Q

What are the 2 types of covalent structures?

A
  • Simple molecular lattice

- Giant covalent lattice

67
Q

Examples of element that form simple molecular lattices?

A

Ne, H₂, O₂, N₂, H₂O

68
Q

The structure of a (solid) simple molecular lattice?

A

1) The atoms within each molecule are held together by strong covalent bonds.
2) The different molecules are held together by weak intermolecular forces (such as London forces - ALWAYS NAME THEM!)

69
Q

Melting and boiling point of simple molecular lattices?

A
  • Simple molecular lattices have low melting and boiling points
  • because only a small amount of energy is needed
  • to break the weak intermolecular forces between the molecules
70
Q

Electrical conductivity of simple molecular lattices?

A
  • Simple molecular lattices do not conduct electricity
  • as they do not contain any charged particles
  • that are free to move
  • and carry the charge.
71
Q

Solubility of simple molecular lattices (which have a simple molecular lattice)?

A
  • Simple molecular lattices are generally soluble in non-polar solvents such as hexane.
  • This is because weak london forces can form between the simple covalent molecules and the non-polar solvents.
  • which helps break down the simple molecular lattice allowing the substance to dissolve.
72
Q

Examples of giant covalent structures?

A

Diamond, graphite, graphene and Si

73
Q

Structure of diamond?

A
  • Giant covalent lattice where each carbon atom forms four strong covalent bonds, with four other carbon atoms so all electrons are involved in bonding.
  • 3D regular tetrahedal structure-
74
Q

Structure of Graphite?

A
  • Giant Covalent Lattice where each carbon atom forms three covalent bonds with three other carbon atoms.
  • This means one electron from each carbon is not involved in bonding and is delocalised, free to move within its layer
  • The carbon atoms form layers with a hexagonal arrangement of atoms.
  • The layers are held together by the weak intermolecular forces, London forces, between them.
75
Q

Structure of Silicon (and silicon dioxide)?

A

Giant covalent lattice where each silicon atom forms four covalent bonds with four other silicon atoms so all electrons are involved in bonding.
- 3D regular tetrahedral structure
Siliconis however mostly found as SiO₂, where each silicon atom is covalently bonded to oxygen atoms and all electrons are involved in bonding

76
Q

Melting and boiling point of giant covalent lattices?

A

Giant covalent lattices have high melting and boiling points because large amounts of energy are needed to break the strong covalent bonds that hold the structure together.

77
Q

Eelctrical conductivity of giant covalent structures?

A

Most giant covalent structures cannot conduct electricity because there are no charged particles which are free to move and carry the charge.
However graphite and graphene contain delocalised electrons that can carry the charge within their sheets so can conduct electricity.

78
Q

Solubility of giant covalent structures?

A

Giant covalent lattices are insoluble in both polar and non-polar solvents because the covalent bonds (lattice) is too strong to be broken by the attraction between solvent molecules and carbon atoms.

79
Q

What is the shape of a molecule or ion determined by?

A
  • The number and type of electron pairs (lone pair or bonding pairs) in the outer shell surrounding the central atom determines its shape.
  • Electron pairs have a negative charge so will repel each other.
  • THE SHAPE ADOPTED WILL BE THE SHAPE THAT ALLOWS. ALL THE ELECTRON PAIRS TO BE AS FAR APART AS POSSIBLE.
80
Q

What are bonded pairs?

What are bonding regions?

A

Bonded pairs refer to the bonding electrons in a single covalent bond. Bonding regions refer the the bonding electrons in a single, double or triple covalent bonds.

81
Q

If all the electron pairs around the central atom were bonded pairs/bonding regions…

A

…all these pairs would repel each other equally. Double and triple covalent bonds repel in the same way/by the same amount as a single covalent bond

82
Q

Why do double and triple covalent bonds repel by the same amount?

A

This because the double (or triple) bond is fixed in place between the central atom and the other bonded atom.

83
Q

What shape and bond angle is formed when there is 1 bonding region around the central atom only? Give an example of a molecule like this.

A

Shape: Linear
Bond angle: No angle
Example: H₂

84
Q

What shape and bond angle is formed when there is 2 bonding region around the central atom only? Give an example of a molecule like this.

A

Shape: Linear
Bond angle: 180°
Example: CO₂

85
Q

What shape and bond angle is formed when there is 3 bonding region around the central atom only? Give an example of a molecule like this.

A

Shape: Trigonal Planar
Bond angle: 120°
Example: BF₃

86
Q

What shape and bond angle is formed when there is 4 bonding region around the central atom only? Give an example of a molecule like this.

A

Shape: Tetrahedral
Bond angle: 109.5°
Example: CH₄

87
Q

What shape and bond angle is formed when there is 5 bonding region around the central atom only? Give an example of a molecule like this.

A

Shape: Trigonal bypiramid
Bond angle: 90° and 120°
Example: PCL₅

88
Q

What shape and bond angle is formed when there is 6 bonding region around the central atom only? Give an example of a molecule like this.

A

Shape: Octahedral
Bond angle: 90°
Example: SF₆

89
Q

NOW DRAW THE 3D SHAPES OF ALL THE DIFFERENT EXAMPLES AND LABEL THE BOND ANGLE.

A

Check page 92

90
Q

How to draw 3D shapes:

A

Use line, bold wedges and dotted wedges.
A normal ine shows a bond in the plane of the paper.
A bold wedge shows the bond coming out from the plane of the paper towards you
A dotted wedge shows the bond going out from the plane of the paper away from you

91
Q

Molecules can also have lone pairs on the central atom. How is alone pair different to a bonded pair and what does this mean?

A

A lone pair of electrons is slightly more electron-dense than a bonded pair. This means that a lone pair repels more than a bonded pair.

92
Q

What is the relative strengths of the repulsion between the different electron pairs?

A

Lone pair/Lone pair
Lone pair/Bonded pair
Bonded pair/Bonded pair

93
Q

How does the lone pair affect a bond angle?

A
  • Each lone pair reduces the bond angle around the central by 2.5°.
  • If there are four electron pairs around a central atom
  • the molecule is based of tetrahedral
  • with a bond angle of 109.5°
  • however this is assuming all of the electron pairs are bonding pairs.
  • However, if one of them was a lone pair, the bond angle decreases by 2.5° to give 107°.
  • This change in angle cause the shape of the molecule to change, in this case it would be pyramidal.
  • If another electron pair of the four electron pairs around the central atom was a lone pair,
  • the bond angle decreases again by 2.5° to give 104.5°.
  • The shape would now be non-linear.
94
Q

What shape and bond angle is formed when there is 1 lone, and 3 bonding apirs around the central atom only? Give an example of a molecule like this.

A

Shape: Pyramidal
Bond angle: 107°
Example: NH₃

95
Q

What shape and bond angle is formed when there are 2 lone, and 2 bonding apirs around the central atom only? Give an example of a molecule like this.

A

Shape: Non-linear
Bond angle: 104.5°
Example: H₂O

96
Q

Non-linear and pyramidal are the only shapes you need to know when thre is a lone pair (based of tetrahedral). However, you may have to predict the angle still even if based of other shapes such as octahedral or trigonal planar.

A
97
Q

The principles discussed about shape and bond angle can also be applied to molecular ions such as…

A

NH₄⁺ which has a bond angle of 109.5 and is tetrahedral as there are four bonded pairs around the central atom.

*The charge is distributed across the whole molecule - shown by square brakcets and a + sign.

98
Q

What is electronegativity?

A

Electronegativity is a measure of the attraction a bonded atom has for the shared pair(s) of electrons in its covalent bond.

99
Q

General trend of electronegativity in the periodic table?

A

Electronegativity increases towards fluorine (being the most electronegative) from all directions (so towards top right excluding noble gases)

100
Q

Which elements are most electronegative?

A

Nitrogen, oxygen, chlorine and fluorine.

101
Q

Why do diatomic molecules, with identical bonding atoms, have non-polar bonds?

A

If the two bonding atoms in a diatomic molecule are identical, they have equal electronegativities and thus equal attraction for the bonding electrons in the covalent bond. This makes the bond non-polar

102
Q

Why do diatomic molecules with different bonding atoms usually have polar bonds?

A
  • If the two bonding atoms in a diatomic molecule are not identical,
  • they are likely to have different electronegativities,
  • so one of the bonding atoms (the more electronegative one)
  • will have a greater attraction for the bonding electrons
  • so they will be held closer to that bonding atom.
  • This makes the bond polar.
103
Q

Draw a polar bond in HCL and a non-polar bond in H₂.

A

Pg.94

104
Q

If a bond is polar, how will the atoms be changed?

A
  • The difference in electronegativity causes the bonds to become polarised
  • which results in a permanent dipole - a difference in charge between the two bodning atoms.
  • The more electronegative bonding atom will have a slight negative charge,
  • whilst the less electrongeative bonding atom will have a slight positive charge.
105
Q

The greater the difference in electronegativity =

A

= the greater the permanent dipole (i.e. charge difference between the bonded atoms)

106
Q

Molecules with a polar bond may have an overall dipole across the molecules making the molecule polar overall. What determines if a molecule is polar or not?

A

The arrangement of polar bonds i.e. the symmetry, in a molecule determines whether or not the molecule will have an overall dipole, that would make it polar.

107
Q

What makes a molecule polar/have an overall dipole?

A

For a molecule to be polar:
1) it must have polar bonds
2) the molecule must not be symmetrical.
Being asymmetrical means that the dipoles in the molecule due to polar bonds do not cancel out, so that there is an overall dipole across that molecule, making it polar.
In the case of molecules such as HCl, there is only one polar bond - there needs to be two in order for it cancel out, making it a polar molecule with polar bonds.

108
Q

What makes a molecule non-polar/not have an overall dipole even if the bonds are polar?

A 
For a molecule to be non-polar:
1) it must not have any polar bonds
or
1) it has polar bonds but
2) is symmetrical 
Being symmetrical means that dipoles in the molecule due to polar bonds cancel out, so that there is no overall dipole across the molecule - making the molecule non-polar.
EXAMPLE: CCL₄
109
Q

Oxygen is one of the most electronegative elements. This means when bonded to other atoms, these bonds will be polar. However this not mean all molecules containing oxygen will be polar. Why not - use H₂O and CO₂.

A
  • H₂O is a polar molecule with polar bonds
  • CO₂ is a non-polar molecule with polar bonds
    In CO₂ the carbonyl bonds are polar but the molecule is symmetrical (linear shape), so any dipoles due to polar bonds cancel out, so no dipole exists across the whole molecule, making it non-polar.
    In H₂O, the OH bonds are polar but the molecule is asymmetrical (non-linear shape), so dipoles due to due to polar bonds may not cancel out, so an overall dipole exists across the whole molecule, making it polar.
110
Q

Which molecules are alway asymmetrical?

A

Trigonal bipyramid

Non-linear

111
Q

Which molecules are sometimes asymmetrical?

A

Linear
Trigonal planar
Tetrahedral

112
Q

What are intermolecular forces?

A

The attractive forces that occur between neighbouring molecules - no transferring or sharing of electrons

113
Q

What are intramolecular forces?

A

These refer to chemical bonds such as ionic and covalent bonds - involves transfer or sharing of electrons

114
Q

What are the two main types of intermolecules forces?

A
  • Hydrogen bonding

- Van der Waals forces

115
Q

Relative strength of the different forces? Strongest to weakest?

A

1) Ionic and covalent bond
2) Hydrogen bonds
3) Permanent dipole-dipole forces (any intermolecular forces, excluding hydrogen bonding, where a permanent dipole is invovled)
4) London (dispersion) forces

116
Q

What forces do the ‘Van der Waals Forces’ refer to?

A
  • Permanent dipole-induced dipole interactions
  • Permanent dipole-permanent dipole interactions
  • Induced dipole-induced dipole interactions (London forces)
117
Q

What is a permanent dipole-induced dipole interactions?

A
  • Some molecules have a permanent dipole, due to polar bonds, this means it has a slightly positive end (δ+) and a slightly negative end (δ-).
  • This means when it is near non-polar molecules, it’s dipole is able to cause electrons in the shells of the non-polar molecules to shift by being repelled by the δ- end or attracted to the δ+ end.
  • This causes the non-polar molecule to become slightly polar and then an attraction occurs.
  • The permanent dipole has induced a dipole in the other molecule.
118
Q

What is a permanent dipole- permanent dipole interaction?

A

Molecules with permanent dipoles will also be attracted to other molecules with permanent dipoles, where the oppositely charged ends attract one another e.g. δ- on one molecule is attracted to δ+ on another molecule.

119
Q

What is a London force?

A
  • A London force is the same as an induced dipole-induced dipole interaction.
  • Molecules that do not contain dipoles are non-polar, however non-polar molecules may still be attracted to one another due to London forces.
  • These are caused by the constant random movement of electrons in atom’s shells.
  • This movement unbalances the distribution of charge within the electron shells, so at any moment, there will be an instantaneous dipole across the molecule.
  • The instantaneous dipole induces a dipole in neighbouring molecules, which in turn induces further dipoles on their neighbouring molecules.
  • The small induced dipoles attract one another, causing London forces/induced dipole-induced dipole interactions to occur between them.
120
Q

How does the London forces between molecules get stronger?

A
  • The strength of London forces increases with increasing number of electrons.
  • The greater the number of electrons , the stronger the induced dipole and the greater the attraction between molecules.
121
Q

As you go down group 8, describe and explain the trend in boiling point?

A

The only intermolecular forces between the noble gases are London forces. As you go down the noble gases, the boiling point increases (the boiling points are quite low already - all are negative values). This is because as you go down the group, the number of electrons in their shells increases, so the strength of the London forces between them also increases, so more energy is required to break these intermolecular forces as you go down the group

122
Q

Which atoms must hydrogen be bonded to allow hydrogen bondign to occur and why?

A

Oxygen, nitrogen, and fluorine, because they are are all highly electronegative. This means, the dipoles formed are particularly strong.

123
Q

Draw the hydrogen bonding between:

A
  • two water molecules

- a water molecule and an ammonia molecule

124
Q

Hydrogen bodnign has effects on the property of water. Name them.

A
  • Ice is less dense than water.
  • Water has a higher than expected melting and boiling point
  • High surface tension and viscosity
125
Q

Why is ice less dense than water?

A
  • Ice (solid state) has an open lattice, unlike water, where the H2O moelcules(dont say water molecules) are held further apart than in water, and fixed in position by hydrogen bonds. When the ice melts, the rigid hydrogen bonds collapse, allowing the water molecules to move closer together.
126
Q

Why does water have a higher than expected melting and boiling point?

A

In order to melt or boil H₂O, the hydrogen bonds between them need to be broken. These are much stronger than other intermolecular forces, and thus requires a lot of energy to be broken.

127
Q

Why do other molecules compared to H₂O in group 16, have a low boiling point, such as H₂S, H₂Se, H₂Te?

A

This is due to the fact that only H₂O molecules form hydrogen bonds between each other, and these are much stronger than the London forces between the other molecules so require a lot more energy to break.

128
Q

Why does water have high surface tension?

A

Water has high surface tension because water molecules at the surface form strong hydrogen bonds between each other

129
Q

Why is graphite slippery?

A

This is because the layers in graphite are held together by weak intermolecular forces (London forces) that only require a small amount of energy to break, allowing the layers to slide over each other.

130
Q

Why does graphite have low density?

A

Layers are quite far apar compared to the length of the covalent bonds, so graphtie has low density - used to make strong lightweight equipment.

131
Q

Structure of graphene?

A
  • Giant Covalent Lattice
  • where each carbon atom forms three covalent bonds with three other carbon atoms.
  • This means one electron from each carbon is not involved in bonding and is delocalised, free to move within the lattice
  • The carbon atoms form a single layers with a hexagonal arrangement of atoms i.e. a single sheet of graphite
132
Q

Applications of graphene?

A
  • High strength, low mass and good electrical conductivity means is has applications in high speed electronics and aircraft tech.
  • Due to its flexibility and transparency, useful as touchscreens on electrical devices.

Chemistry OCR AS Level (9 decks)

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