Chapter 2 - Chemical Bonding Flashcards Preview

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Flashcards in Chapter 2 - Chemical Bonding Deck (56):
1

What is electronegativity?

Electronegativity is a measure of the tendency of an atom to attract a bonding pair of electrons. The larger its number, the larger its force of attraction.

2

What is the trend of electronegativity across a period?

It increases across the period.
Across the period, effective nuclear charge increases due to the increase in nuclear charge while shielding effect remains relatively constant. The forces of attraction between nucleus and outermost electrons are stronger and the nucleus has a greater attraction for electrons of another nucleus.

3

What is the trend of electronegativity down a group?

It decreases down the group.
The atom size increases down the group due to its increase in number of electron shells. There is a longer distance and weaker attraction between nucleus and outermost electron, thus the nucleus has a weaker attraction for electrons of another nucleus.

4

What are the 3 most electronegative electrons?

F > O > N

5

How does the difference in electronegativity predict the type of bond formed between 2 atoms? (2)

- Covalent bonds are formed between non metal atoms with similar electronegativity, resulting in sharing electrons
- Ionic bonding occurs between a metal and a non metal with a great difference in electronegativity, resulting in electron transfer
(metallic bonding is a different case)

6

What are the 2 factors affecting the strength of metallic bonding?

1) The larger the number of valence electrons contributed per metal atom into the sea of delocalized electrons, the greater the number of delocalized electrons, the stronger the metallic bonding
2) The higher the charge and the smaller the radius of the metal cation, the higher its charge density and thus the stronger the metallic bonding. Compare the ratio of charge/ionic radii between 2 cations.

7

Define charge density.

Charge density of an ion is the amount of charge per unit surface area of the ion.

8

What are 5 physical properties of metals?

1) High melting and boiling points due to the fairly strong attraction between delocalized electrons and cations. (Boiling point is a better guide to strength of metallic bonding as on melting, metallic bonding is only weakened. Boiling completely breaks metallic bonding)
2) Good electrical conductivity. 'sea' of delocalized electrons act as mobile charge carriers
3) Good thermal conductivity. Thermal energy is picked up by the delocalized electrons as additional kinetic energy, making them move faster. Energy is transferred throughout the rest of the metal by moving electrons, so heat flows quickly from the hotter part to cooler part of metal.
4) Usually hard: hardness depends on how easily we can move particles from their fixed position in the solid lattice (strength of metallic bonding)
5) Malleable and ductile: When a large stress is applied to a piece of solid metal, the layers of ions will slide over each other into new positions. Its overall shape changes but it does not break or shatter as the sea of delocalized mobile electrons prevents repulsion between cations, keeping metallic bonding intact.

9

What are alloys?

Alloys are mixture of metals involving incorporation of small quantities of other elements into the pure metal.

10

How does alloying make metals harder?

Atoms of the added element have a different size. This disrupts the orderly arrangement of main metal atoms such that they can no longer slide over each other easily when a force is applied.

11

In ionic bonding, how are individual ions collected together in a giant lattice?

- Oppositely charged ions attract each other while ions of like charges repel each other. Thus, each cation in an ionic lattice will be surrounded by a number of anions, and each anion will in turn be surrounded by a number of cations.
- The number of ions that surround another ion of the opposite charge is the coordination number of the central ion, which depends on the relative sizes of ions and their relative charge.
- If one ion is very small, there will not be room for many oppositely charged ions around it (max 4). If the cations and anions are nearly equal in size, one ion can be surrounded by 8 others. The intermediate case is 6.
- To gain electrical neutrality, a cation with +2 charge needs twice its number of anions with -1 charge.

12

What affects the strength of ionic bonding?

Lattice energy, which is the heat evolved when 1 mole of pure ionic solid is formed from its constituent gaseous ions. The larger its magnitude, the stronger the ionic bonding.

13

What are the 2 factors affecting lattice energy?

The higher the charge and the smaller the radii of ions, the larger the lattice energy.
Lattice energy= k charge of cation x charge of anion / radius of cation + radius on anion. Thus the charge of ions is a more important factor.

14

What are 3 physical properties of ionic compounds?

1) high melting and boiling points. Melting is due to the strong electrostatic forces of attraction between oppositely charged ions. On melting, there is still a significant level of attraction between mobile ions in liquid state, resulting in high melting points.
2) Good electronic conductors when molten or aqueous. Mobile charge carriers (ions) are only free to move in these states. The electrical conductivity of an aqueous solution increases when its concentration increases (more ions), and also when it contains highly charged ions. (Al3+ vs Na+)
3) Hard and rigid but brittle. In an ionic lattice, oppositely charged ions are held in fixed positions throughout the crystal lattice by strong ionic bonding. Moving the ions out of position requires large amounts of strength to overcome the bonding. However, if a strong enough shearing force is applied, it will force ions of like charges to move next to each other. Repulsion between them will cause the lattice to shatter apart, that's why it is brittle.

15

What is covalent bonding?

Covalent bonding is the electrostatic forces of attraction between the positively charged nucleus of both the bonded atoms and their shared pair of electrons. A covalent bond forms when orbitals of 2 atoms overlap, where the overlap region is occupied by a pair of electrons.

16

How is a covalent bond formed?

With similar electronegativity, each non-metal atom holds on to its own electrons tightly and tend to attract other electrons as well. As 2 atoms come close to one another, they would initially repel (from electron of other atom). As they get closer, the electrons encounter an attraction to the opposite nucleus that is stronger than their mutual repulsion. When they get too close, they would repel each other again since both atoms are positively charged. Electron-nucleus is just balanced by electron-electron and nucleus-nucleus repulsions, minimizing potential energy and forming a covalent bond.

17

What is a sigma bond?

It is a bond formed by a "head-on" overlap of 2 orbitals, which is present in all bonds. There can be only 1 sigma bond between 2 atoms.

18

What is a pi bond?

It is a bond formed by a "side-on" overlap of 2 orbitals. For a pi bond to exist, a sigma bond must first be present.

19

What are the 2 factors determining the strength of covalent bonds?

1) smaller atoms form shorter and stronger bonds as overlap between valence orbitals is more effective than bigger atoms and the valence orbitals are close to the nucleus. Electrons in overlapping orbitals experience a stronger attraction to the nucleus. (Except O-O and F-F bonds possibly due to the repulsion of lone pairs on adjacent atoms)
2) Multiple bonds are stronger and shorter than single bonds: the more valence orbitals that overlap, the more electrons being shared, the stronger the attraction between the 2 positively charged nuclei and the shared electrons.

20

What is dative bonding?

It is when the donor atom, which must have a lone pair, provides both bonding electrons, to the acceptor atom, which must have an empty orbital in its valence shell. Once formed, it is no different from an ordinary covalent bond and can be either sigma or pi bond.

21

What are 4 exceptions to the octet rule?

1) Electron deficient species, usually the central atom, is yet to achieve 8 electrons.
2) Expansion of octet, usually the central atom. Elements from period 3 onwards can utilize their d orbitals in covalent bonding as they are energetically accessible (usable). They can form more than 4 covalent bonds to other atoms.
3) Radicals: unpaired electrons in the outermost shell which is very reactive. Most radical molecules have a central atom (usually Grp 15 or 17) that has an unpaired electron and does not have an octet structure.
4) Most cations do not actually have noble gas electronic configurations, where only Group 1, Group 2, scandium and aluminum form cations with noble gas electronic configuration.

22

What are the 2 main principles of the VSEPR theory?

1) the electron pairs (or groups) around the central atom of the molecule (or ion) arrange themselves as far apart as possible so as to minimize mutual repulsion.
2) the repulsion between electron pairs decrease in this order:
lone pair-lone pair > lone pair-bond pair > bond pair-bond pair

23

What are the 5 basic shapes and bond angles of electron pairs? (no lone pairs)

1) 2 electron pairs: linear, 180
2) 3 electron pairs: trigonal planar, 120
3) 4 electron pairs: tetrahedral, 109.5
4) 5 electron pairs: trigonal bipyramidal, 90&120
5) 6 electron pairs: Octahedral, 90

24

What happens to the molecular shape if there are lone pairs?

Although lone pairs are important in determining the shape of a molecule, they are "ignored" when the molecule's shape is described.

25

What happens to the bond angle if there are lone pairs?

A lone pair is attracted by only 1 nucleus why a bond pair is attracted by 2 nuclei. A lone pair electron cloud is thus less "elongated" than a bond pair. A lone pair takes up more space around the central atom and its electrons are closer to the nucleus of the central atom. Thus, a lone pair repels other electron pairs around it more strongly than a bond pair would. The bond pairs would then be pushed more closely together, resulting in slightly smaller bond angles.

26

What happens to the molecular shape if it contains double or triple bonds?

The 4 or 6 electrons in the double or triple bond would count as one group of electrons.

27

How does hybridisation of atomic orbitals work?

Atomic orbitals mix to form new orbitals called hybrid orbitals which involve in sigma bonding. Orbitals are identical in shape and identical in energy, intermediate between energies of s and p orbitals. A hybrid orbital has more p character and thus it tends to form a longer and weaker bond.

28

What is the sp^3 hybridisation of carbon?

Mixing one s orbital with three p orbtials give 4 sp3 hybrid orbitals, which gives carbon 4 single bonds. They are directed at 109.5 from each other so as to minimize repulsion.

29

What is the sp^2 hybridisation of carbon?

Mixing one s orbital with two p orbitals give 3 sp2 hybrid orbitals and one unhybrised p orbital in carbon, which gives it 2 single bonds and 1 double bond. They are directed at 120 from each other to minimize repulsion.

30

What is the sp hybridisation of carbon?

Mixing one s orbital with one p orbital gives 2 sp hybrid orbitals and 2 unhybridised p orbitals, which gives carbon 1 single bond and 1 triple bond, or 2 double bonds. They are directed at 180 from each other to minimize repulsion.

31

What are polar covalent bonds?

It occurs when the sharing of electrons is unequal due to difference in electronegativity, causing the electron cloud to be elliptical, as one atom attracts the electrons in the covalent bond more strongly to itself. It causes a partial charge separation (dipole)

32

What does it mean when a cation has high polarising power?

When cations have high charge density, they tend to attract the electrons of the anion. When the cation attarcts the electrons of the anion too strongly, the electron cloud will be distorted. This is when the cation polarises the electron cloud of the anion and thus the cation has high polarising power.

33

What does it mean when an anion has high polarisability?

It is when anions are large. A large anion's outermost electrons are further away from the nucleus, so the electron cloud is less attracted by the nucleus and can be easily "pulled away" by a cation, and becomes distorted. This is when an anion is polarised by a cation and thus the anion has high polarisability.

34

What in the overall dipole/net dipole?

It is the sum of all bond dipoles and it depends on the shape of the molecule and sometimes the magnitude of the individual bond dipoles.

35

How does net dipole determine the polarity of a molecule?

- if individual bond dipoles cancel out each other, the molecule has no overall net dipole and is therefore non-polar.
- if bond dipoles reinforce each other, a polar molecule is formed.
- the greater the difference in electronegativity, the greater the polarity of a covalent bond, and the greater its dipole.

36

Define intermolecular forces of attraction.

They are electrostatic forces of attraction which arise from interactions of the net dipole of simple covalent molecules.

37

What are the 5 types of intermolecular attractions?

1) dispersion forces
2) permanent dipole-permanent dipole attractions
3) hydrogen bonding
4) ion-dipole attractions
5) permanent dipole-induced dipole attractions

38

How are dispersion forces formed?

It occurs in all molecules but are particularly important for non-polar molecules. Since electrons in a molecule or an atom are in constant motion, an unequal distribution of electrons can occur at any moment. These fluctuations in electron distribution results in the non-polar molecules forming an instantatenous dipole, which induces a corresponding dipole in its neighbouring molecule, thus causing an attraction between 2 molecules. Since the dipole might change an instant later, dispersion forces are perpetuate, weak and short-lived.

39

Why can noble gas atoms be liquefied and even solidified at very low temperatures?

Let us use Ar as an example. An instantaneous dipole in one Ar atom induces a dipole in its neighbouring Ar atom, attracting the Ar atoms together. This process takes place throughout a sample of Ar, thus making liquefying and solidfying possible.

40

What are 2 factors affecting the strength of dispersion forces? (which affects boiling point)

1) The larger the number of electrons in a molecule or atom, the larger its electron cloud, and the more polarisable it is, since the electrons are further away from the nucleus. Instantaneous dipoles and induced dipoles can form more easily, resulting in stronger dispersion forces. Thus, dispersion forces are stronger for larger particles.
2) For non-polar substances with similar number of electrons, the strength of dispersion forces is influenced by molecular shapes. Shapes that allow more points of contact have more surface area which electron clouds can be distorted, so induced dipoles can form more easily, resulting in stronger dispersion forces.

41

What are permanent dipole-permanent dipole attractions and how are they formed?

It is a stronger form of intermolecular forces on top of dispersion forces in polar molecules. When polar molecules lie near one another (in liquids and solids), the positive dipole of one molecule attracts the negative dipoleof another. The electrostatic attraction between these oppositely charged ends of any 2 molecules with permanent dipoles is called pd-pd attractions.

42

What is the factor affecting the strength of permanent dipole-permanent dipole attractions?

For molecules with approximately the same size of electron cloud, its higher dipole moment means it is more polarisable. The positive dipole of a molecule gets more attracted to the negative end, causing stronger pd-pd attractions.

43

What is hydrogen bonding and how are such bonds formed?

It is a stronger type of pd-pd attraction.
When a hydrogen atom bonds to another atom, its sole electron is used up in the covalent bonding, In F-H, O-H and N-H bonds, the bonding electrons are pulled strongly towards the highly electronegative fluorine, nitrogen or oxygen atoms, giving hydrogen a very positive charge as it is only left with 1 proton which gives it a very significant positive charge. The highly electron deficient hydrogen atom experiences a very strong attraction from the lone pair of electrons on the highly electronegative F/O/N atom of an adjacent molecule, causing hydrogen bonding.

44

What are 2 requirements for hydrogen bonding to occur between 2 molecules?

1) One molecule must contain a hydrogen atom bonded to a highly electronegative F/O/N atom.
2) the other molecule must contain an atom with a lone pair of electrons on a F/O/N atom.

45

What are 3 unusual properties of water caused by hydrogen bonding?

1) It would have a higher than expected boiling point due to the large amount of heat needed to overcome the strong intermolecular hydrogen bonding.
2) Liquid water has higher surface tension (large force required to break the layer of water molecules on its surface). The hydrogen bonded H2O molecules form an array across the surface of water, allowing objects that are expected to sink, to "float".
3) Although most substances have a higher density in solid than in liquid state, ice is less dense than liquid water. In solid state, H2O molecules are held at fixed positions and arranged in an orderly manner to form a regular lattice such that hydrogen bonding is maximised to 4 per molecule. In liquid state, the hydrogen bonding between water molecules in ice are positioned in a roughly tetrahedral shape around each O atom. This produces an open attice, with empty spaces between the H2O molecules. The more random arrangement of hydrogen bonding in liquid water results in H2O molecules packing much more closely, taking up less space. The lattice structure of ice occupies a larger volume for the same mass of liquid water, hence ice has a lower density than water.

46

How does hydrogen bonding induce dimerisation in carboxylic acids?

Dimer are 2 structurally similar monomers joined by covalent/intermolecular bonds, in this case hydrogen bonds.

47

Why is the boiling point of H2O higher than that of HF?

H2O contains 2 protonic H atoms and 2 lone pairs on the O atom; there is an average of 2 hydrogen bonds formed per H2O molecule. HF contains 1 protonic H atom and 3 lone pairs of F atom; there is an average of 1 hydrogen bond formed per HF molecule. Thus, there is a more extensive hydrogen bonding among H2O molecules, requiring a larger amount of energy to overcome the bonding.

48

What are 2 physical properties of simple covalent compounds?

1) Melting and boining point are usually lower than that of other compounds.
2) They are unable to conduct electricity as there are no mobile charge carriers; all electrons are not free to move throughout a sample of that substance and there are no mobile ions.

49

What are giant molecular compounds?

The atoms in the compounds are held together by covalent bonds that extend in 3 dimensions throughout the sample. Their physical properties reflect the strength of covalent bonds.

50

What is the structure of diamonds and what are its 4 physical properties?

In diamonds, each C atom is tetrahedrally bonded to 4 other C atoms, an arrangement that extents throughout the giant lattice.
1) diamond is very hard due to its giant lattice and strong covalent bonds
2) extremely high melting point
3) insoluble in any solvent
4) electrically neutral as all electrons are either held in covalent bonds or held by nuclei; not free to move

51

What is the structure of quartz (SiO2) and what are its 2 physical properties?

In quartz, each Si atom is covalently bonded to 4 O atoms, and each O atom is bonded to 2 Si atoms in a pattern that extends throughout the giant lattice.
1) very hard
2) very high melting point

52

Why are there 2 different values when measuring the distances between covalent molecules? What are their trends?

Using I2 as an example,
The shorter distance is between 2 bonded I atoms in the same molecule, known as the I-I bond length. Half its distance is the covalent radius. ("Single covalent" in data booklet)
The longer distance is between 2 non-bonded iodine atoms in adjacent molecules, known as the van der Waals distance. Half its distance is the van der Waals radius.("van der Waals" in data booklet)
van der Waals radius is always larger than its covalent radius. Just like covalent radii, van der Waals radii decreases across a period and increases down the group.

53

What is an ion-dipole attraction?

Ion-dipole attractions occur between an ion and the oppositely charged end of a dipole. Because it is very extensive, it releases a lot of heat when formed.

54

What are permanent dipole-induced dipole interactions?

They occur when a polar molecule distorts the electron cloud of a nearby atom or non-polar molecule. It is generally weak and releases little heat when formed. It is very similar to dispersion forces and is comparable to it in terms of strength.

55

How to explain solubility of two substances? (4)

1) identify type of interactions holding solute particles together
2) identify type of interactions holding solvent particles together
3) identify type of interaction that can occur between solute and solvent particles
4) compare their relative strengths. A solute is only soluble in a solvent if the energy released from solute-solvent interactions is greater than or comparable to the energy needed to overcome solute-solute and solvent-solvent interactions.

56

How do we compare the strength of different types of chemical bonding? (7)

In descending order, the strongest type of chemical bonding to the weakest:
1) ionic
2) covalent (nuclei to shared electrons)
3) metallic
4) ion-dipole attraction
5) hydrogen bonding
6) pd-pd
7) pd-id
8) dispersion forces