Topic 2: Bonding and Structure Flashcards Preview

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Flashcards in Topic 2: Bonding and Structure Deck (52):
1

What is an ionic bond?

- an ionic bond is the strong electrostatic attraction between two oppositely charged ions

2

What two things affect the strength of the ionic bond?

- ionic charges
- ionic radii

3

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

- in general, the greater the charge on an ion, the stronger the ionic bond
- therefore, higher melting/boiling point

4

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

- 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

5

What affects the size of an ion?

- the number of electron shells
- its atomic number

6

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

- 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

7

Describe isoelectronic ions and its trends

- 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

8

What structures do ionic compounds form?

- giant ionic lattice structures
- a lattice is a regular structure
- the structure is giant because the basic unit is repeated

9

Why do ionic compounds form giant ionic structures?

- each ion is electrostatically attracted in all directions to ions of the opposite charge

10

What are the physical properties of ionic compounds?

- 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

11

What is a covalent bond?

- the strong electrostatic attraction between the two positive nuclei and the shared electrons in the bond

12

What is bond length?

- 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
- this is bond length

13

What is bond enthalpy?

- the higher the electron density between the nuclei (i.e. the more electrons in the bond), the stronger the attraction between the atoms, the higher the bond enthalpy and the shorter the bond length

14

What is dative covalent bonding?

- one atom donates both electrons to a bond

15

What do molecular shape depend on?

- depends on the number of pairs of electrons in the outer cell of the central atom

16

What type of electron pair repel more than others?

- lone pairs
- this means the greatest angles are between lone pairs of electrons
- bong angles between bonding pairs are often reduced because they are pushed together by lone pair repulsion

17

Give the bond angles for different types of electron pairs

no lone pair: 109.5 e.g. methane
one lone pair: 107 e.g. ammonia
two lone pairs: 104.5 e.g. water

18

What do you call molecules with 2 electron pairs around the central atom?

- linear

19

What do you call molecules with 3 electron pairs around the central atom?

- no lone pair: trigonal planar
- one lone pair: bent (119)

20

What do you call molecules with 4 electron pairs around the central atom?

- no lone pair: tetrahedral
- 1 lone pair: trigonal pyramidal
- 2 lone pairs: bent

21

What do you call molecules with 5 electron pairs around the central atom?

- no lone pairs: trigonal bipyramidal (90 and 120)

22

What do you call molecules with 6 electron pairs around the central atom?

- no lone pairs: tetrahedral (90)

23

describe the properties of giant covalent structres

- 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
- insoltuble: 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

24

Why can graphite conduct electricity?

- 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

25

Describe metallic bonding

- 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 lattic 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

26

What are the properties of metallic bonding?

- 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

27

What is electronegativity?

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

28

What are the properties of more electronegative elements?

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

29

What is a polar bond?

- 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

30

What is a dipole?

- 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

31

What is an overall dipole?

- a dipole caused by the presence of a permanent charge across a molecule
- polar bonds don't always make polar molecules

- in a 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

32

What are intermolecular forces?

- intermolecule forces are force between molecules
- much weaker than covalent, ionic or metallic bonds

33

What are the three intermolecular forces learnt?

- London dispersion forces (instantaneous dipole-induced dipole bonds)
- permanent dipole-permanent dipole bonds
- hydrogen bonding

34

Describe London dispersion forces

- 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, 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 are that atoms are attracted to each other

35

What makes London Dispersion Forces stronger?

- molecules with greater surface areas due to bigger exposed electron cloud
- larger molecules due to larger electron clouds

36

What are permanent dipole-permanent dipole bonds?

- the partial polar charges on polar molecules cause weak electrostatic forces of attraction between molecules, know as permanent dipole-permanent dipole bonds
- as these bonds happen as well as London forces, molecules generally have a higher boiling and melting points

37

Explain hydrogen bonding

- 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

38

Draw hydrogen bonding on water, ammonia and hydrogen fluoride

GOOGOO

39

What molecules often form hydrogen bonds?

- Organic molecules that often contain -OH and -NH groups
- e.g. alcohols and amines

40

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

- 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 a high 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

41

Describe water's boiling point

- it has a fairly high boiling point, despite being a small molecule
- its ability to form H bonds gives it high boiling point

42

Why does ice float on water?

- ice is an example of a simple molecular structure
- in ice, the water molecules are arranged so that there is a maximum number of H bonds
- the lattice structure formed this way 'wastes' space
- as ice melts, some of the H bonds are broken and the lattice breaks down, allowing molecules to fill the spaces
- ice is much less dense than water, so it floats

43

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

- 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

44

How do substances dissolve in another

- bonds in the substance have to break
- bonds in the solvent have to break
- new bonds have to form between the substance and the solvent

45

What are the two main types of solvent?

- 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

46

Why do ionic substances dissolve in water?

- water is 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

47

Why do alcohols dissolve in polar solvents?

- 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
- carbon chain part of alcohol isn't soluble so more carbon atoms, less soluble alcohol will be

48

Why don't halogenoalkanes dissolve in water?

- 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

49

What does like dissolves like mean?

- substances usually dissolve best in solvents with similar intermolecular forces

50

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

- 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

51

Learn the property table!11!!!

GOOOOOOOOOOO

52

What is metallic bonding?

- the strong electrostatic attraction between metal ions and the sea of delocalised electrons