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A Level Chemistry > Topic 2: Bonding and Structure > Flashcards

Topic 2: Bonding and Structure Flashcards

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1
Q

What is an ionic bond?

A
  • an ionic bond is the strong electrostatic attraction between two oppositely charged ions
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2
Q

What two things affect the strength of the ionic bond?

A
  • ionic charges
  • ionic radii
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3
Q

How do ionic charges affect the strength of the ionic bond?

A
  • in general, the greater the charge on an ion, the stronger the ionic bond
  • therefore, higher melting/boiling point
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4
Q

How do ionic radii affect the strength of the ionic bond?

A
  • smaller ions can pack closer together than the larger ions
  • electrostatic attraction gets weaker with distance, so small, closely packed ions have stronger ionic bonding than larger ions
  • therefore, ionic compounds with small, closely packed ions have higher melting and boiling points
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5
Q

What affects the size of an ion?

A
  • the number of electron shells
  • its atomic number
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6
Q

Explain the trend in ionic radii as you go down a group

A
  • the ionic radius increases as you go down a group
  • as you go down the group, the ionic radius increases as the atomic number increases
  • this is because more electron shells are added
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7
Q

Describe isoelectronic ions and its trends

A
  • isoelectronic ions are ions of different atoms with the same number of electrons
  • the ionic radius of a set of isoelectronic ions decreases as the atomic number increases
  • as you go through a series of ions, the number of electrons stay the same, but the number of protons increase
  • this means that the electrons are attracted to the nucleus more strongly, pulling them in, so the ionic radii decreases
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8
Q

What structures do ionic compounds form?

A
  • giant ionic lattice structures
  • a lattice is a regular structure
  • the structure is giant because the basic unit is repeated
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9
Q

Why do ionic compounds form giant ionic structures?

A
  • each ion is electrostatically attracted in all directions to ions of the opposite charge
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10
Q

What are the physical properties of ionic compounds?

A
  • hard, brittle crystalline substances
  • high boiling and melting points: ions are held together by strong electrostatic attraction.
  • are often soluble in water and other polar solvents, but insoluble in non-polar solvents: this tells us that the particles are charged, as ions are pulled apart by polar molecules e.g. water
  • don’t conduct electricity when solid but do when molten or dissolved: ions are fixed in a position by strong ionic bonds in a solid, but are free to move as a liquid, or in a solution
  • ionic compounds can’t be shaped: if you pull layers over each other, there will be strong repulsion between the like-charges, so they are brittle
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11
Q

What is a covalent bond?

A
  • the strong electrostatic attraction between the two positive nuclei and the shared electrons in the bond
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12
Q

What is bond length and what affects it?

A
  • the distance between the nuclei of the two atoms that are covalently bonded
  • in covalent molecules, the positive nuclei are attracted to the area of electron density between the two nuclei
  • however, the two positively charged nuclei repel each other, as do the electrons
  • the distance between the two nuclei is the distance where the attractive and repulsive forces balance each other
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13
Q

What is a discrete molecule?

A
  • an electrically neutral group of two or more atoms held together by chemical bonds
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14
Q

What is dative covalent bonding?

A
  • the shared pair of electrons in the covalent bond come from only one of the bonding atoms
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15
Q

Describe electron pair repulsion theory

A
  • the shape of a molecule or ion is caused by repulsion between the pairs of electrons, both bond pairs and lone pairs, that surround the central atm
  • the electron pairs arrange themselves around the central atom so that the repulsion between them is at a minimum
  • lone pair-lone pair repulsion > lone pair-bonding pair repulsion > bond pair-bond pair repulsion
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16
Q

What are the shapes and angles of different molecules?

A
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17
Q

What is electronegativity?

A
  • the relative ability of an atom in a covalent bond in a molecule to attract the bonding electrons in a covalent bond to itself
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18
Q

What are the properties of more electronegative elements?

A
  • higher nuclear charges (more protons in the nucleus)
  • smaller atomic radii
  • therefore, electronegativity increases across periods and up the groups
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19
Q

What type of compound will have a small electronegative difference?

A
  • purely covalent
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20
Q

What type of compound will have a very large electronegativity difference (> 1.7)?

A
  • ionic
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21
Q

What is a polar bond?

A
  • if the bond between two atoms have different electronegativities, the bonding electrons will be pulled towards the more electronegative atom
  • causes electrons to be spread unevenly, so there will be a charge across the bond
  • each atom has a partial charge
22
Q

What is a permanent dipole?

A
  • a dipole is a difference in charge between the two atoms caused by a shift in electron density in the bond
  • in a polar bond, the difference in electronegativity between the two atoms causes a dipole
23
Q

What is an overall dipole?

A
  • a dipole caused by the presence of a permanent charge across a molecule
  • polar bonds don’t always make polar molecules
  • in more complicated molecules, if polar bonds point in opposite directions, they’ll cancel each other out, making it non-polar
  • if polar bonds all point in the same direction, then it will be polar
24
Q

What are intermolecular forces?

A
  • intermolecular forces are forces between molecules
  • much weaker than covalent, ionic or metallic bonds
25
Q

What are the three intermolecular forces learnt?

A
  • London dispersion forces (instantaneous dipole-induced dipole bonds)
  • permanent dipole-permanent dipole bonds
  • hydrogen bonding
26
Q

Describe London dispersion forces

A
  • London forces cause all atoms and molecules to be attracted to each other
  • electrons in charge clouds are always moving quickly
  • at any moment, the electrons in an atom are likely to be more to on side than the other
  • at that time, the atom would have a temporary (instantaneous) dipole
  • this dipole can induce another temporary dipole in the opposite direction on a neighbouring atom
  • the two dipoles are attracted to each other
  • the second dipole can induce another dipole in a third atom
  • dipoles are constantly being created and destroyed to due constant movement
  • overall effect is that atoms are attracted to each other
27
Q

What factors affect the size of London forces?

A
  • the more electrons there are in the molecule, the higher the chance that temporary dipoles will form
  • this makes London forces stronger between the molecules
  • shape: long chain alkanes have a larger surface area of contact between molecules for London forces to form than compared to spherical shapes branched alkanes
28
Q

What are permanent dipole-permanent dipole bonds?

A
  • the partial polar charges on polar molecules cause weak electrostatic forces of attraction between molecules, know as permanent dipole-permanent dipole bonds
  • common polar bonds are: C-Cl, C-F, C-Br, H-Cl, C=O
  • as these bonds happen as well as London forces, molecules generally have a higher boiling and melting points
29
Q

Explain hydrogen bonding

A
  • hydrogen bonding is the strongest intermolecular force
  • it only happens when hydrogen is covalently bonded to fluorine, nitrogen or oxygen
  • F, N and O are very electronegative, so they draw the bonding electrons away from the hydrogen atom
  • the bond is so polarised and hydrogen has such as high charge density because its small, that the hydrogen atoms form weak bonds with lone pairs of electrons on the F, N, or O atoms of other molecules
30
Q

What is the bond angle of a hydrogen bond?

A
  • 180º because there are two pairs of electrons around the H atom involved in the hydrogen bond
  • these pairs of electrons repel to a position of minimum repulsion, as far as part as possible
31
Q

What molecules often form hydrogen bonds?

A
  • alcohols
  • carboxylic acids
  • amides
  • proteins
32
Q

Describe and explain how the boiling points of group 7 hybrids vary as you go down group 7

A
  • molecules of hydrogen fluoride form H bonds with each other
  • H bonding is the strongest intermolecular force, so intermolecular bonding in HF is very strong
  • a lot of energy is required to overcome these bonds, so has the highest boiling point
  • From HCl to HI, although the permanent dipole-dipole interactions decreases, the number of electrons in the molecule increases
  • so the strength of London forces increases, which overrides the decreases in strength of permanent dipole-dipole, so increase in boiling points
33
Q

Describe why boiling temperatures increase with longer unbranched alkanes?

A
  1. as molecular mass increases, the number of electrons per molecule increases and so the instantaneous and induced dipoles increase
  2. as the length of the carbon chain increases, the number of points of contact between adjacent molecules increases
    - London forces exist at each point of contact between molecules, so the more points of contact, the greater the overall London forces
34
Q

Why do branched chain alkanes have lower boiling temperatures than their unbranched isomers?

A
  • the more branching, the fewer points of contact between adjacent molecules
  • this leads to a decrease in the overall intermolecular force of attraction between molecules and a decrease in temperature
35
Q

Why does water have a relatively high boiling point?

A
  • it can form two H bonds per molecule because the electronegative oxygen atom has two lone pairs on it
  • so it can form stronger H bonds
36
Q

Why does ice float on water?

A
  • ice is an example of a simple molecular structure
  • in ice, the water molecules are arranged in rings of six, held together by H bonds
  • this creates large areas of open space inside the rings
  • the molecules are held further apart by H bonds than in liquid water
37
Q

Why are alcohols less volatile (have higher boiling points) than similar alkanes?

A
  • alcohols contain a polar hydroxyl group (OH) that has a partially negative charge on the oxygen atom and a partially positive charge on the H atom
  • this polar group helps alcohols form H bonds
38
Q

Why is the predominant bonding in alcohols not always hydrogen bonding?

A
  • as the chain length increases, the London forces will eventually predominate
  • it can be shown by the enthalpy change of vaporization as it measures the amount of energy required to completely separate the molecules into a liquid and converts it into a gas
39
Q

How do substances dissolve in another

A
  • it is a balance of energy required to break bonds in the solute and solvent against energy given out making new bonds between the solute and solvent
40
Q

What are the two main types of solvent?

A
  • polar solvents:
  • made of polar molecules, such as water
  • water molecules bond with each other with H bonds but not all polar solvents can form H bonds
  • non-polar solvents: e.g hexane, which its molecules bind to each other by London forces
41
Q

Why do ionic substances dissolve in water?

A
  • water is a polar solvent
  • when an ionic substance is mixed with water, the ions are attracted to the oppositely charged ends of the water molecules
  • ions are pulled away from the ionic lattice by the water molecules which surround the ions
  • this is called hydration
  • some ionic substances do not dissolve because the bonding between their ions are too strong e.g. Al2O3 because of high charge density
42
Q

Why do smaller alcohols dissolve in polar solvents?

A
  • alcohols can dissolve in water because the polar O-H bond in an alcohol is attracted to the O-H bonds in water
  • H bonds form between the lone pairs on the O and H atoms
  • the carbon chain part of alcohol isn’t soluble so the more carbon atoms, the less soluble alcohol will be
43
Q

Why don’t halogenoalkanes dissolve in water?

A
  • H bonding between water molecules is stronger than the bonds that would be formed with halogenoalkanes
  • but they do form permanent dipole-permanent dipole bonds
44
Q

What does like dissolves like mean?

A
  • substances usually dissolve best in solvents with similar intermolecular forces
45
Q

Why do non-polar substances dissolve best in non-polar solvents?

A
  • they form similar bonds with non-polar solvents
  • water molecules are attracted to each other more strongly than they are to non-polar molecules, so that don’t tend to dissolve easily in water
46
Q

What is metallic bonding?

A
  • the strong electrostatic attraction between the positive metal ions and the sea of delocalised electrons
47
Q

Describe metallic bonding

A
  • metal elements exist as giant metallic lattice structure
  • in metallic lattices, the electrons in the valence shell are delocalised and free to move, leaving a positive metal ion
  • the positive metal ions are electrostatically attracted to the delocalised negative electrons, forming a lattice of closely packed positive ions in a sea of delocalsied electrons
  • overall lattice structure is made of layers of positive metal ions, separated by layers of electrons
48
Q

What are the properties of metallic bonding?

A
  • melting points of metals are generally high: because of strong metallic bonding, with the number of delocalised electrons per atom affecting the melting point
  • the more electrons there are, the stronger the bonding will be and the higher the melting point
  • malleable and ductile: no bonds holding specific ions together. layers of metal ions can slide over each other without disrupting attraction between positive ions and electrons
  • good thermal conductors: delocalised electrons can pass kinetic energy to each other
  • good electrical conductors: delocalised electrons are free to move and can carry a current
  • insoluble: expect in liquid metals. due to strength of metallic bonds
49
Q

What three main factors affect the strength of metallic bonding?

A
  1. Number of protons:
    - the more protons, the stronger the bond
  2. Number of delocalised electrons per atom (the outer shell electrons are delocalised)
    - the more delocalised electrons, the stronger the bond
  3. Size of ion
    - the smaller the ion, the stronger the bond
50
Q

Why can graphite conduct electricity?

A
  • its carbon atoms forms sheets, with each carbon atom sharing three of its outer shell electrons with three other carbon stoms
  • the fourth outer electron in each atom is free to move between the sheets, making graphite a conductor of electricity
51
Q

Describe the properties of giant covalent structures

A
  • have very high melting and boiling points: a lot of energy is required to break very strong electrostatic bonds
  • very hard: due to very strong bonds through the lattice arrangement
  • good thermal conductors: travel easily through stiff lattices
  • insoluble: the covalent bonds means atoms are more attracted to their neighbours in the lattice than to solvent molecules. insoluble in polar solvents because no ions
  • cannot conduct electricity: there are no charged ions or free electrons
52
Q

Look at summary table of solid lattice properties

A

A Level Chemistry (19 decks)

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