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AS Chemistry > Chemical Bonding > Flashcards

Chemical Bonding Flashcards

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1
Q

Define

ionic bond

A

The electrostatic attraction between oppositely charged ions.

This is formed when a metal atom transfers electrons to a non-metal atom, e.g. NaCl, MgO, CaCl2

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2
Q

Define

covalent bond

A

A bond formed by the sharing of pairs of electrons between two atoms.

In a simple covalent bond, one electron from each atom contributes to the pair.

The atoms are held together by the electrostatic forces between the electrons and the nuclei. e.g. H2, O2, HCl, CO2, CH4, C2H4

For a covalent bond to form, atomic orbitals containing unpaired valence electrons must overlap each other.

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3
Q

Define

dative covalent bond

A

A covalent bond in which both electrons in the bond come from the same atom.

The donor atom having a lone pair of electrons shares the electrons with the acceptor atom having an empty orbital.

e.g. NH4+

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4
Q

Explain dative covalent bonding in aluminum trichloride

A

The aluminium trichloride (AlCl3) molecule exists as a vapour with single covalent bonds above 750oC, but forms two dative bonds with another AlCl3 molecule at lower temperatures, forming the dimer molecule Al2Cl6.

  • Bond angle as AlCl3 = 120o
  • Bond angle as Al2Cl6 = 109.5o
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5
Q

Define and list characteristics of a

sigma (σ) bond

A

A single covalent bond formed by the ‘end-on’ or ‘head-to-head’ overlap of atomic orbitals

  • occurs with s and p orbitals, e.g. in H2, HCl and Cl2
  • electron density is concentrated symmetrically along the inter-nuclear axis between the two atoms
  • ∴ bond is strong, requires lots of energy to break
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6
Q

Define and list characteristics of a

pi (π) bond

A

A covalent bond involving the sideways overlap of atomic orbitals

  • occurs with p and d orbitals
  • overlap is symmetrical to axis
  • charge is distributed above and below the inter-nuclear axis between the two atoms
  • ∴ bond is weaker, requires less energy to break
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7
Q

Define

hybridisation of atomic orbitals

A

The process of mixing atomic orbitals so that each has some character of each of the orbitals mixed

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8
Q

Outline the hybridisation of carbon in forming:

  1. C-C single bond, e.g. in methane
  2. C=C double bond, e.g. in ethene
  3. C≡C triple bond. e.g. in ethyne
A
  • A small amount of energy is used to promote an electron from the 2s orbital to the empty 2p orbital to give 4 unpaired electrons
    • new configuration is 1s22s12p3
    • The extra energy released when the bonds form more than compensates for the initial input
  1. sp3 orbital is formed, with each C forming 4 σ bonds (1s22sp34)
  2. sp2 orbital is formed, with each C forming a σ bond for each H, and one σ and one π bond with the other C (1s22sp232p1)
  3. sp orbital is formed, with one σ bond and two π bonds between Cs (1s22sp22p2)
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9
Q

Explain why molecules have different shapes

A
  • Electron pairs will repel each other as far as possible
  • There are two types of electron pairs: bonding pairs and lone pairs
  • Lone pairs are more compact and provide more repulsion
  • Strength of repulsion: lone-lone>lone-bond>bond-bond
  • This causes different molecule shapes
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10
Q

List the number of lone pairs, bond pairs and the bond angle(s) in this type of molecular shape:

trigonal planar (e.g. BF3, CO3-)

A
  • 0 lone pairs
  • 3 bonding pairs
  • 120o bond angle
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11
Q

List the number of lone pairs, bond pairs and the bond angle(s) in this type of molecular shape:

linear (e.g. CO2, BeCl2)

A
  • 0 lone pairs
  • 2 bonding pairs
  • 180o bond angle
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12
Q

List the number of lone pairs, bond pairs and the bond angle(s) in this type of molecular shape:

tetrahedral (e.g. CH4, NH4+)

A
  • 0 lone pairs
  • 4 bonding pairs
  • 109.5o bond angle
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13
Q

List the number of lone pairs, bond pairs and the bond angle(s) in this type of molecular shape:

pyramidal (e.g. NH3, PCl3)

A
  • 1 lone pair
  • 3 bonding pairs
  • 107o bond angle
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14
Q

List the number of lone pairs, bond pairs and the bond angle(s) in this type of molecular shape:

non-linear / bent (e.g. H2O, SCl2)

A
  • 2 lone pairs
  • 2 bonding pairs
  • 104.5o bond angle
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15
Q

List the number of lone pairs, bond pairs and the bond angle(s) in this type of molecular shape:

octahedral (e.g. SF6, PCl6)

A
  • 0 lone pairs
  • 6 bonding pairs
  • 90o bond angle
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16
Q

List the number of lone pairs, bond pairs and the bond angle(s) in this type of molecular shape:

trigonal bipyramidal (e.g. PF5)

A
  • 0 lone pairs
  • 5 bonding pairs
  • 90o and 120o bond angles
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17
Q

What type of molecular shape is this and what bond angle(s) does it have? (give examples)

  • 0 lone pairs
  • 3 bonding pairs
A

trigonal planar (e.g. BF3, CO3-)

120o bond angle

18
Q

What type of molecular shape is this and what bond angle(s) does it have? (give examples)

  • 0 lone pairs
  • 2 bonding pairs
A

linear (e.g. CO2, BeCl2)

180o bond angle

19
Q

What type of molecular shape is this and what bond angle(s) does it have? (give examples)

  • 0 lone pairs
  • 4 bonding pairs
A

tetrahedral (e.g. CH4, NH4+)

109.5o bond angle

20
Q

What type of molecular shape is this and what bond angle(s) does it have? (give examples)

  • 1 lone pair
  • 3 bonding pairs
A

pyramidal (e.g. NH3, PCl3)

107o bond angle

21
Q

What type of molecular shape is this and what bond angle(s) does it have? (give examples)

  • 2 lone pairs
  • 2 bonding pairs
A

non-linear / bent (e.g. H2O, SCl2)

104.5o bond angle

22
Q

What type of molecular shape is this and what bond angle(s) does it have? (give examples)

  • 0 lone pairs
  • 6 bonding pairs
A

octahedral (e.g. SF6, PCl6)

90o bond angle

23
Q

What type of molecular shape is this and what bond angle(s) does it have? (give examples)

  • 0 lone pairs
  • 5 bonding pairs
A

trigonal bipyramidal (e.g. PF5)

90o and 120o bond angles

24
Q

Describe

hydrogen bonding (with examples)

A

The strongest type of intermolecular force, formed between molecules having a hydrogen atom bonded to one of the most electronegative elements (F, O or N)

The hydrogen attached directly to one of the most electronegative elements acquires a significant amount of positive charge. It is attracted to a lone pair on another higly electronegative atom bonded to a hydrogen. This is a type of dipole-dipole interaction.

25
Q

Describe the hydrogen bonding shown in ammonia, water and hydrogen fluoride

(also, compare the strength of hydrogen bonding between them)

A
  • Ammonia: two hydrogen bonds formed: one lone pair on nitrogen, one hydrogen attracted to another molecule (least strength as N is less electronegative than O or F)
  • Water: four hydrogen bonds formed: two lone pairs on oxygen, both hydrogens are attracted to one other molecule each (stronger than NH3 but weaker than HF)(in liquid state, these bonds aren’t fixed and can move about)
  • Hydrogen fluoride: four hydrogen bonds formed: three lone pairs on fluoride, one hydrogen attracted to another molecule (strongest as F is the most electronegative atom)
26
Q

Explain the effects of hydrogen bonding on the physical properties of water (liquid state)

A
  1. High b.p. as there are exactly the right numbers of δ+ hydrogens and lone pairs so that every one of them can be involved in hydrogen bonding (compared to 1. ammonia, where each nitrogen only has one lone pair so there aren’t enough lone pairs to go around to satisfy all the hydrogens, and 2. HF, where there aren’t enough δ+ hydrogens for the lone pairs
  2. High viscosity as hydrogen bonding reduces the ability of water molecules to slide over each other
  3. High surface tension as hydrogen bonds exert a significant downwards force at teh surface of the liquid
27
Q

Explain the effect of hydrogen bonding on boiling point

A
  • Hydrogen bonds are the strongest type of intermolecular force, requiring high amounts of energy to break
  • ∴ b.p. is higher than expected with substances having hydrogen bonds
  • e.g. though fluorine is at the top of the halogens group and so should form compounds with lower b.p.s, HF has a high b.p. due to hydrogen bonding
  • alcohols have higher b.p.s than alkanes due to the hydrogen bond formed at the OH group
28
Q

Define

electronegativity

and explain the factors affecting electronegativity

A

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

  1. Nuclear charge: +ve protons attract -ve e’s
  2. Atomic radius: more attraction closer to the nucleus
  3. Shielding: less subshells means less shielding so more attraction

Electronegativity decreases down a group, atomic radius ↑ shielding ↑ outweighs higher nuclear charge so attraction

Electronegativity increases across a period as nuclear charge ↑atomic radius outweighs higher shielding so attraction ↑

29
Q

Define

a dipole and a dipole moment

A

A separation of charge in a molecule.

A dipole moment is a vector quantity showing the direction of the dipole going from positive to negative.

30
Q

Define

bond energy

A

The energy needed to break one mole of a particular bond in one mole of gaseous molecules.

31
Q

Define

bond length

and describe the relation to bond energy and reactivity

A

The distance between the nuclei of two bonded atoms.

The shorter the bond length, the greater the bond energy (more orbital overlapping makes a stronger bond) and the lower the reactivity as a high amount of energy is needed to break the bond.

Down a group, bond length ↑ as atomic radius ↑ ∴ bond energy ↓

32
Q

Define

bond polarity

A

The property of a covalent bond being polar, which is where the two bonding electrons are not shared equally by the atoms in the bond. The atom with the greater share of the electrons has a partial negative charge δ-, and the other has a partial positive charge δ+.

i.e. the bond has a dipole moment due to the atoms in the bond having different electronegativities

33
Q

Describe the differences between polar and non-polar molecules

A
  • Polar covalent bonds
    • Bond formed with atoms of different electronegativity
    • The greater the difference in electronegativity of the two bonded atoms, the greater is the ionic character
  • Non-polar covalent bonds
    • Bond formed between:
      • Identical atoms: the electronegativity of both atoms is the same so pair of electron shared equally
      • Symmetrical polyatomic molecules: dipoles of bond exert equal & opposite effects hence cancel charge
    • Non-polar molecules have no overall charge
34
Q

Explain how ionic molecules can display covalent properties

A

If one atom in an ionic bond is highly electronegative (i.e. with a high positive charge density), it may be able to attract some electrons from the other atom, thus causing some covalent properties.

This is called ion polarisation, which is the distortion of the electron cloud on an anion by a neighbouring cation. The distortion is greatest when the cation is small and higly charged

35
Q

Explain how bond polarity influences chemical reactivity

A
  • Many chemical reactions are started by a reagent attracting one of the electrically charged ends of a molecule
  • The more polar a bond is, the more attraction there will be, ∴ more polar bonds cause a molecule to be more reactive
  • With non-polar bonds, there is no such attraction, so molecules such as alkanes are not very reactive
36
Q

Define

induced dipole-dipole forces (also referred to as van der Waals’ forces)

and explain how they are formed, and what factors influence their strength

A

The weak forces of attraction between molecules caused by the formation of temporary dipoles.

This occurs due to the constant random movement of electrons in the electron cloud: at an instant, a non-polar molecule develops poles due to the distortion of electron density giving rise to an instantaneous dipole, which is able to induce a dipole in the adjacent molecules.

Strength is affected by:

  • increasing number of contact points between molecules
  • increasing number of electrons (+ protons)

This is the weakest type of intermolecular force, and is always present between molecules (even noble gases and non-polar molecules e.g. Br2) in addition to other types of forces they may have.

As the number of electrons increases down a group, van der Waals’ forces increase and so m.p. and b.p. generally increase

37
Q

Define

permanent dipole-dipole forces

and explain how they are formed and what factors influence their strength

A

A type of intermolecular force between molecules that have permanent dipoles.

This occurs between neighbouring polar molecules as the δ- charge on one molecule and the δ+ charge on the other molecule cause a weak attractive force between them.

Highly polar bonds will have greater dipole moments, and therefore stronger permanent dipole-dipole forces. Non-polar molecules have no dipole moments and so no such forces.

For example, CHCl3 has a liquid state at r.t.p. compared to gaseous CH4

38
Q

Define

metallic bond

and list the factors that influence its strength

A

The electrostatic attraction between metal cations and delocalized mobile electrons.

Strength is affected by:

  • size of metal ion
  • number of outer electrons donated to the electron cloud
  • both of the above factors influence charge density; the greater the charge density, the stronger the bond
39
Q

List and explain the physical properties of ionic compounds

A
  • Has an ionic lattice (solid) structure at r.t.p. with high m.p. & b.p. as high amounts of energy needed to overcome strong electrostatic forces
  • Strong but brittle as distortion of lattice leads to similar ions being adjacent and repelling one another
  • Conducts electricity when molten or in solution as mobile ions conduct
  • Insoluble in non-polar solvents, soluble in polar solvents such as water, as the ion-dipole attractions formed between the water molecules and ions produce enough energy to overcome the electrostatic forces between ions
40
Q

List and explain the physical properties of covalent compounds

A
  • Low m.p. & b.p. as weak induced dipole-dipole attractions between molecules
  • Electrical insulator as no mobile ions or electrons
  • Non-polar compounds are insoluble in water as water molecules are not attracted to them
  • Some molecules are soluble in water as hydrogen bonds can be formed, e.g. ethanol
41
Q

List and explain the physical properties of metals

A
  • Has a metallic lattice (solid) structure at r.t.p. (except Hg), with high m.p. & b.p. as high amounts of energy needed to overcome strong electrostatic forces
  • Good electrical conductor as delocalised electrons move when a potential difference is applied
  • Good conductor of heat as delocalised electrons can transfer kinetic energy
  • Ductile & malleable as electrons move with cations so the bond is not broken
  • Hard due to strong bond
42
Q

Describe chemical reactions in terms of bonds

A
  1. Reactant bonds are broken - endothermic process
  2. Product bonds are formed - exothermic process

AS Chemistry (18 decks)

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