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Bonding Flashcards

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1
Q

Describe the ionic bond

A

The ionic bond is the electrostatic force of attraction between oppositely charged ions.
•The cations and anions are held closely together in a lattice
•A result of electron transfer from metal atom to non-metal atom in order to attain a noble gas electronic configuration

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2
Q

What is a covalent bond?

A

A covalent bond is the electrostatic force of attraction between the nuclei of two atoms and a shared pair of electrons.
•A result of electron sharing Between two or more atoms to form a molecule.

Molecule: two or more atoms held together by covalent bonds.

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3
Q

Why do some elements have more than 8 electrons in their out shell when forming covalent bonds? Give examples of such elements.

A

Elements in period 3 such as phosphorus, sulfur and chlorine (with oxygen and fluorine) can accommodate the extra electrons in their unfilled p or d orbitals in their 3rd principal quantum shell.

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4
Q

What is a coordinate bond. Give examples

A

Coordinate bond:
the sharing of a pair of electrons in which both the electrons in the bond come from the same atom.

Examples: NH4+, CO, Al2Cl6

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5
Q

What are the conditions for the formation of a coordinate bond?

A
  1. One of the atoms has a lone pair of electrons.

2. The other atom has an unfilled outermost orbital to accommodate the electrons; it is electron deficient

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6
Q

What is the bond energy and how does it affect reactivity?

A

Bond energy:
the energy required to break one mole of a particular covalent bond in the gaseous state.
Unit:kJ/mol

Influences the activation energy of the reaction:
the higher the bond energy the greater the energy required to break bonds in order for the reaction to occur.

Triple bonds are stronger then double bonds and double bonds are in turn stronger than single bonds.

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7
Q

What is the bond length and how does that of a double bond compare to that of a single bond?

A

Bond length:
the distance between the nuclei of two covalently bonded atoms.

Double bonds are generally shorter than single bonds beacause:
•Their is a greater quantity of negative electrons between the two nuclei.
•Hence, there is a greater force of attraction between the nuclei and the electrons; pulls the atoms closer together
This results in a stronger bond

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8
Q

Describe the formation of molecular orbitals

A

Molecular orbitals are formed when atomic orbitals overlap to form a combined orbital, containing two electrons.

Covalent bonds are formed when atomic orbitals overlap

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9
Q

What determines the strength of a covalent bond?

A

The amount of overlap of atomic orbitals: the greater the overlap the stronger the bond

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10
Q

What is hybridization?

A

The mixing of atomic orbitals into new hybrid orbitals with different shapes,energies etc.

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11
Q

Give examples of hybridization

A

Sp3, sp2 and sp hybrid orbitals

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12
Q

What are the properties of sigma bonds?

A
  1. Sigma bonds are covalent bonds formed From the end on overlapping of orbitals.
  2. The electron density lies on the axis connecting the two nuclei; electron density of each sigma bond is symmetrical about a line joining the two nuclei.
  3. The first bond between two atoms is always sigma, the rest are pi bonds.
  4. Rotation is possible around a Sigma bond
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13
Q

What are the properties of pi bonds?

A
  1. Pi bonds are formed from the sideways overlapping of adjacent P orbitals (above and below the sigma bond.)
  2. The electron density lies above and below the axes connecting the two nuclei.
  3. Rotation is not possible about a pi bond
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14
Q

Describe the hybridization in carbon in terms of electron promotion.

A

C (ground state): 1s2 2s2 2px1 2py1 2pz0

  • there is a small energy gap between the 2S and 2P orbitals in the ground state.
  • thus, it is easy to promote an electron from the 2S to the empty to 2P orbital to give four unpaired electrons.
  • The extra energy released when the bonds are formed compensates for the energy required for the promotion of the electron.

Carbon is said to be in the excited state:
C : 1s2 2s1 2px1 2py1 2pz1

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15
Q

Describe the SP3 hybridization in carbon with examples.

A

Example : Methane, CH4

  • Carbon makes FOUR single SIGMA bonds.
  • The 1S orbitals of each of the four hydrogen atoms overlap with a SP3 hybrid orbital to produce a new molecular orbital,containing a pair of electrons.

NOTE: shape of sp3 hybrid orbitals

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16
Q

Describe the SP2 hybridization in carbon with the aid of an example.

A
  • occurs when carbon is bonded to THREE ATOMS.
  • ONE S and TWO p orbitals form the sp2 hybrid orbital, the other p orbital remains unhybridised.
  • Carbon make THREE SIGMA binds and ONE PI bond.

double bond:

Example in C2H4:

  1. Each carbon atom is bonded to ONE other Carbon atom and TWO hydrogen atoms.
  2. ONE SIGMA and 1 PI BOND is formed in the double bond between the carbon atoms.
  3. TWO other sigma bonds are formed with each of the remaining two hydrogen atoms.

Note:

  1. The sigma bonds are found by the SP2 hybrid orbitals.
  2. Is pi bond is formed by the unhybridized P orbital
  3. The hydrogen nuclei of embedded in the S-SP2 overlap.
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17
Q

Describe the SP hybridization in carbon and give an example.

A
  • Occurs when carbon is bonded to TWO atoms.
  • TWO SIGMA BONDS and TWO PI BONDS are formed.
  • One 2s and one 2p orbital form the sp hybridized orbital, the remaining two 2p orbitals remain unhybridised.

Note: only the sp hybridized orbitals form the sigma bonds the unhybridised p orbitals form the pi bonds.

Example: Carbon dioxide

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18
Q

What is the formula for the type of hybridization?

A

Hybridization= Number of SIGMA bonds + lone pairs

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19
Q

Use the hybridization theory to explain the bonding in PCl3.

A
  1. Phosphorus has five valence electrons And thereby needs three more electrons to attain a stable electronic configuration.
  2. Each chlorine atom needs one more electron to complete their outer shell and attain a stable electronic configuration.

Therefore in PCl3: 3 sigma bonds and 1 lone pair

Hybridization type= 3 + 1 = 4

Sp3

Therefore, The 3p orbital of each chlorine atom overlaps with the SP3 hybrid orbitals of phosphorus to form three molecular orbitals ( one lone pair)

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20
Q

Explain why nitrogen can form NCl3 but not NCl5

A

In NCl3:
•Type of hybridization= 3 sigma + 1 lone pair = 4 = sp3
• The sP3 hybrid orbitals overlap with the 3P orbitals of chlorine to three molecular orbitals.

In NCl5:
However, for nitrogen to form and NCL5, nitrogen has to promote one electron to the next principal quantum level (3S orbital). This requires a large amount of energy which is not recovered in bonding.

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21
Q

Describe the metallic bonding

A

Metallic bonding:
The electrostatic force of attraction between positive ions in a Lattice a sea delocalized electrons.

• in a metal, the atoms are packed closely together in a regular arrangement called a Lattice.
• These metal atoms tend to lose their outer shell electrons and become positive ions.
• The outer shell electrons occupy new energy levels and all free to move throughout the metal lattice
The electrons are said to be delocalized.

The electrostatic force of attraction acts in all directions.

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22
Q

What are delocalized electrons?

A

Delocalized electrons are electrons that are not associated with any particular atom.

In metals the electrons move throughout the metallic structure between the metal ions when a voltage is applied.

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23
Q

What are the factors determining the strength of a metallic bond?

A
  1. The size of the positive charge on the ions in the metal lattice:
    increasing positive charge on the irons increases the strength of the metallic bond
  2. The size of the metal ion in the lattice (ionic radius]:
    The smaller the ionic radius the stronger the metallic bond.
  3. The number of mobile electrons lost per atom:
    The greater the number of electrons delocalized per atom, the stronger the metallic bond.
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24
Q

Explain in terms of metallic bonding why aluminum has a higher Melting point than sodium.

A

Within the Al lattice, the atoms lose three electrons to form Al3+ ions. However, in Na, the atoms lose only one electron to form Na+ ions.
Hence the metallic bonding in Al is stronger than Na because:
1. The size of the positive charge on the ions in the lattice is greater.
2. The size of the metal ions is less
(Recall: ionic size decreases across a period)
3. The number of electrons lost to the sea of delocalised electrons is greater.

As a result, a greater amount of energy is required to overcome the electrostatic forces in Al than Na

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25
Q

Define electronegativity.

What is its purpose?

A

The power of a particular atom that is covalently bonded to another atom to attract the bonding pair of electrons towards itself.

The greater the value of the electronegativity, the greater the ability of an atom to attract the electrons in a covalent bond towards itself.

Used to understand how intermolecular forces work

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26
Q

What are the factors influencing electronegativity?

A
  1. Nuclear charge:
    atoms in the SAME PERIOD with a greater positive nuclear charge are more likely to attract the bonding pair of electrons.
  2. Atomic radius:
    Atoms in the SAME GROUP in which the outer electrons are farther from the nucleus are less likely to attract the bonding pair of electrons.
    •the pull of the positive nucleus on the electron pair is lower.
  3. Shielding:
    The greater the number of inner electron shells and subshells, the lower the EFFECTIVE NUCLEAR CHARGE on the bonding electrons.
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27
Q

Describe and explain in electronegativity across a period.

A

The electronegativity INCREASES across a period because:

  1. The nuclear charge on the atoms increases.
  2. Atomic radius decreases.
  3. Shielding effect of inner shell electrons remains constant.
    Hence, electrons are more strongly attracted to the nucleus.
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28
Q

Describe and explain in electronegativity down a group.

A

The electronegativity DECREASES down a group because:

  1. As the number of electron shells increases, the shielding effect of inner electrons increases.
  2. The atomic radius increases
  3. The effect of the above factors outweigh the effect of the increase in nuclear charge.

Hence, electrons are less strongly attracted to the positive nucleus.

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29
Q

What is the most electronegative element and why?

A

Flurorine.

Top of its group and at the end of the period

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30
Q

Why can’t noble gases be ascribed an electronegativity value?

A

Noble gases cannot form bonds

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31
Q

How can the paulings electronegativity scale be used to predict the character of a bond?

A

•If the electronegativity difference between the atoms forming the bond is high (2.0 or more), the compound is IONIC
•if the electronegativity difference is low (below 1.0), the compound is COVALENT.
•A zero value indicates no ionic character in the bond
• some compounds are not entirely covalent or ionic and have some ionic character in them; these have Intermediate electronegativity differences.
Example: 1.0

32
Q

When is a molecule non-polar?

A
  1. When the Np values of the two atoms forming the covalent bond are the same- the pair of electrons is equally shared.
  2. A molecule may contain polar bonds but is itself a non-polar molecule.
33
Q

What happens when a covalent bond is formed between two atoms having different Np values?

A

A POLAR bond is formed:

The more electronegative atom attracts the bonding pair of electrons towards itself.

As a result:
1. The centre of the positive charge does not coincide with the centre of the negative charge in the molecule.
2. The electron distribution is asymmetric.
3. The two atoms are partially charged
• the less electronegative atom is denoted by: delta positive
•the more electronegative atom is denoted by: delta negative
• (A dipole is set up)

Example : HCl
• H delta positive
• Cl delta negative

34
Q

How is the degree of polarity of a molecule measured?

A

As a dipole moment:

An arrow that points to the delta negative end of the dipole

35
Q

Define bond polarity.

A

The partial separation of charge when two DIFFERENT atoms are joined by a covalent bond, which results in a unequal attraction of the bonding pair of electrons.

• Due to the difference of Np values of the different atoms.

36
Q

What must be taken into account when determining the polarity of a molecule containing MORE THAN TWO atoms?

A
  1. The polarity of each BOND- bond polarity.

2. The arrangement of the bonds in the molecule

37
Q

How is the lack of polarity of some molecules despite having polar bonds accounted for?

A
  • Despite having polar bonds, the arrangement of the said bonds are in such a fashion that the DIPOLE MOMENTS CANCEL OUT
  • I.e same type of bonds have dipoles pulling in opposite directions.

Example: Tetrachloromethane, CCl4

  • 4 polar C-Cl bonds point towards 4 corners of the tetrahedron
  • the dipoles in each bond cancel out each other, so the molecules is non polar
38
Q

Define polar bonds.

A

Polar bonds occur when the electron pair in the bond is drawn towards the atom with the larger electronegativity, making one end of the molecule slightly positive compared to the other.

39
Q

Describe the relationship between polarity and chemical reactivity

A

Many chemical reactions are started by a reagent attacking one of the electrically charged ends of a polar molecule.

For example:
Chloromethane, C2H5Cl, Is far more reactive than ethane as reagents such as OH- ions can attack the delta positive C atom of the polarized C-Cl bond.

Such attack is not possible in ethane because the C- H bond is VIRTUALLY NON-polar
This explains why alkanes are. It very reactive.

Thus polar molecules are more reactive than non-polar molecules.

40
Q

What are intermolecular forces?

A
  • forces of attraction between molecules/atoms
  • these forces must be overcome when a particular substance is melted or boiled.
  • responsible for the state of a substance.
41
Q

What are the three main types of intermolecular forces?

A
  1. Instantaneous dipole-induced dipole forces (Id-id forces); aka London dispersion forces.
  2. Permanent dipole-permanent dipole forces (pd-pd forces)
  3. Hydrogen bonding
42
Q

What are id-id forces?

A

Definition:

The weakest intermolecular forces of attraction resulting from temporary instantaneous dipoles induced in BOTH POLAR AND NON-POLAR molecules; also occur between individual atoms ( noble gases)

43
Q

How do Id-id forces arise?

A
  1. The electron charge cloud in a non-polar molecule ( or atom) are NOT STATIC and is in CONSTANT MOTION.
  2. As a result, at a given time the distribution of the electron charge clouds will not be EXACTLY SYMMETRICAL; one end of the molecule/ atom will have a SURPLUS OF NEGATIVE CHARGE than the other
  3. A TEMPORARY DIPOLE is set up
  4. This dipole can in turn Induce other dipoles in neighboring molecules
  5. As a result, forces of attraction(id-id) arise between the delta positive end in one molecule and the delta negative end of a neighboring molecule
44
Q

Why are the dipoles considered temporary?

A

The electron clouds are in constant motion.

45
Q

What are the factors affecting the strength of id-id forces?

A
  1. The number of electrons in the molecule/ atom:
    The greater the number of electrons the stronger the forces
  2. The number of contact points/surface area between the molecules:
    The larger the number of contact points, the stronger the forces
46
Q

How does an Increasing number of electrons increase the strength of id-id forces?

A

•Allows for a greater likelihood of of an electron charge cloud distortion occurring and thus the greater FREQUENCY AND MAGNITUDE OF THE TEMPORARY DIPOLES.

47
Q

How does an increase number of contact points between molecules increase the strength of the id-id forces?

A

•The larger the surface area of a molecule, the larger the number of contact points it will have with adjacent molecules.
• allows for a greater ability to induce a dipole in an adjacent molecule.
Thus, there are greater id-id forces

48
Q

Why does butane have a higher melting and boiling point than methylpropane despite both having the same number of electrons?

A
  • Butane has a larger surface area (more contact points) than 2-methylpropane, although they have the same molecular formula..
  • straight chain molecules have higher boiling points than their branched chain isomers because:
  1. The greater the degree of branching, the more spherical the molecule becomes.
  2. This reduces the surface area of contact (less contact points)
  3. Less id-id forces and thus the melting point decrease
49
Q

What are polar molecules and how are they formed?

A

Polar molecules are molecules with a PERMANENT dipole.

A molecule is polar and thus has a permanent dipole moment when:

  1. The bonds are polarized
  2. The dipoles of the polarized bonds do not cancel out each other (they are asymmetrical)
50
Q

What are p.d-p.d forces?

A

Attractive intermolecular forces which result from the permanent dipoles in molecules.

I.e attractive force between the delta positive charge on one molecule and the delta negative charge on a neighboring molecule.

51
Q

What is the ground state and excited state of electrons?

A

•The ground state of an electron, the energy level it normally occupies, is the state of lowest energy for that electron.

There is also a maximum energy that each electron can have and still be part of its atom.
•Beyond that energy, the electron is no longer bound to the nucleus of the atom and it is considered to be ionized.
•When an electron temporarily occupies an energy state greater than its ground state, it is in an EXCITED STATE . An electron can become excited if it is given extra energy

52
Q

Explain the promotion of electrons across orbitals.

A

Each orbital has a specific energy associated with it. For an electron to be boosted to an orbital with a higher energy, it must overcome the difference in energy between the orbital it is in, and the orbital to which is is going.

53
Q

Describe the conditions for a molecule to have pd-pd forces.

A
  1. The molecule has polar bonds that are asymmetrical- the molecule is polar.
54
Q

What are conditions in which a molecule will not have pd-pd forces?

A
  1. The molecule contains no polar bonds. (Only contains ID-ID forces)
  2. The molecule contains polar bonds that are symmetrical (The molecule itself is non-folder despite having polar bonds)
    I.e in the polar bones have dipoles of the same magnitude pulling in opposite direction.
    Can be seen in a numbe rod linear (CO2) , trigonal planar ( CCL4) and tetrahedral structures ( BF3)
55
Q

What is the effect of the PDPD forces on melting point of compounds that have IDID forces of similar magnitude?

A

It slightly increases the strength of the intermolecular bonding thereby causing the melting points of the compounds being slightly higher than expected from temporary dipoles alone.

E.g propanone (polar) butane ( non-polar)

56
Q

What is hydrogen bonding?

A

The strongest type of intermolecular force but weaker than covalent bonds. It is a strong type of PD-PD force.

57
Q

What are the conditions required for hydrogen bonding to occur between two molecules?

A
  1. One molecule has to have a hydrogen atom covalently bonded to F , O or N ( The three most electronegative elements).

This causes the H atom to acquire significant amount of partial positive charge.

  1. A second molecule to have an fluorine, oxygen or nitrogen atom with an available lone pair of electrons
58
Q

Describe how the hydrogen bond arises.

A
  1. When a hydrogen atom is covalently bonded to a very electronegative atom, the bond is very highly polarized.
  2. The Delta positive charge on the hydrogen atom is high enough for a bond to be formed with a lone pair of electrons on the fluorine, oxygen nitrogen atom of a neighboring molecule.
59
Q

What is the condition for the maximum Bond strength?

A

The angle between the covalent bond to the hydrogen atom and a hydrogen bond is usually 180

60
Q

What does the average number of hydrogen buns fomed per molecule depend on?

A
  1. The number of hydrogen atoms attached to fluorine, oxygen or nitrogen in the molecule
  2. The number of lone pairs present on the fluorine, nitrogen or oxygen
61
Q

Describe the effect of hydrogen bonding on boiling point

A

Substances containing hydrogen bonds have much higher boiling points than would be predicted from van der waal forces alone

62
Q

Describe the Peculiar property of water in terms of enthalpy change of vaporization and boiling point

A

Water has a much higher enthalpy change of vaporization and boiling point than most other group 16 hydrides.

Explanation:

Water is extensively hydrogen bonded therefore this indicates that much more energy is required to overcome these bonds between the water molecules.

Note: rise is due to id-Id forces increasing because of increasing number of electrons

63
Q

Describe the peculiar properties of water in terms of surface tension and viscosity.

A

Water has a high surface tension and high viscosity

Explanation:
1. Hydrogen bonding reduces the ability of water molecules to slide over each other, so the viscosity of water is high

  1. Hydrogen bonds in water also exert a significant downward force at the surface of the liquid this causes the surface tension of water to be higher than most liquids
    • When many water molecules form hydrogen bonds with other water molecules, they form a lattice of water molecules, which is strong and flexible. This creates a high surface tension.
    The surface tension create a surface film, and is responsible for bubbles in the water droplets and capillary action
64
Q

Why are most solids denser than liquids?

A

This is because the molecules are more closely packed in the solid state

65
Q

However why is ice less dense than water?

A
  1. In ice, there is a three dimensional hydrogen bonded network of water molecules.
  2. This produces a rigid lattice in which each oxygen atom is surrounded by tetrahedron of oxygen atoms.
  3. This more open arrangement, due to the relatively long bond lengths of the hydrogen bonds, Allow the water molecules to be slightly further apart than in the liquid.
66
Q

Describe the solubility of substances in water

A

Most covalent compounds are not soluble in water while ionic compounds are.
However, some covalent compounds or server in water because they can form hydrogen bonds with Water.

67
Q

Describe the helical nature of DNA

A

Due to hydrogen bonding:

  1. Molecules of DNA contain and N-H bonds and also contains C=O bonds.
  2. Thus hydrogen atoms can form a hydrogen bond with this electro negative oxygen atom. This results in the molecule spiraling, as the C=O bones and the N-H bonds approach each other.
68
Q

Explain why the boiling point of HF is higher than that of NH3.

A

The H-F bond is more polar than the N-H bond, and there is stronger hydrogen bond in the HF than NH3.

69
Q

Check why H20 has a higher b.p than HF and NH3

A

.

70
Q

Check physical properties and bonding.

A

.

71
Q

Why are ionic compounds solid at room temperature and pressure?

A
  1. There are strong electrostatic forces holding the positive and negative ions together.
  2. The ions are regularly arranged in a lattice , With the oppositely charged ions close to each other
72
Q

Why are ionic compounds soluble in water?

A

Water molecules are polar and are attracted to the ions on the surface of the ionic solid. These attractions are called ion-dipole attractions. These attractions replace the electrostatic forces between the ions which go into the solution.

73
Q

Why are metals insoluble in water?

A

The electrostatic force of attraction between the ions and the delocalised electrons is too great for the water molecules to disrupt the structure and form bonds with the ions.

74
Q

Why are some simple structures insoluble in water?

A

The covalent molecules are non-polar thus water molecules are not attracted to them.
E.g iodine

75
Q

Why are some simple molecules soluble in water?

A

They can form hydrogen bonds with water
E.g ethanol

Some molecules undergo hydrolysis reactions with water
E.g HCL : H+ , Cl -
SiCl4 : .., SiO2

Chemistry (3 decks)

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