4 Bonding

Question 1

distinguish between INTER- and INTRA-molecular forces

Answer 1

INTER: between molecules
INTRA: within a molecule, between atoms

Question 2

ionic bond is

Answer 2

electrostatic attraction;
between oppositely charged ions

Question 3

ionic bond is formed by

Answer 3

transfer of electrons between metal and non-metal atoms

Question 4

covalent bond is

Answer 4

electrostatic attraction;
between nucleus and shared pairs of electrons

Question 5

coordinate/dative covalent bond

Answer 5

some molecules have a lone pair of electrons that can be donated to form a bond with an electron-deficient atom (which has an unfilled outer orbital)
→ both shared electrons are from the same atom

Question 6

as the number of shared electrons in a covalent bond increases…

Answer 6

greater attraction force between electrons
→ atoms pulled closer together
→ bond length decrease, bond strength increase

Question 7

metallic bond

Answer 7

strong electrostatic attraction btwn lattice of cations and sea of delocalized electrons

Question 8

explain why metallic compounds are malleable

Answer 8

apply force → layers slide → lattice not broken but changes shape

Question 9

Explain the boiling/melting point (volatility), electrical conductivity, and solubility of IONIC compounds in terms of their structure

Answer 9

• high BP/MP (low volatility): strong electrostatic attraction force between oppositely charged ions, need large energy to break the bond
• non-conductors when solid but conduct when liquid: ions can move when molten
• soluble in water: the attraction between polar water molecules and ions is strong enough to break the lattice and release hydrated ions

Question 10

Explain the melting point (volatility), electrical conductivity, and solubility of GIANT COVALENT compounds in terms of their structure

Answer 10

• high melting point (low volatility): strong covalent bond, need large energy to break the bond
• non-conductors: no delocalized electron (except graphite)
• insoluble in water or any other polar solvants: solvent molecules cannot
  interact with the non-metal atoms strongly enough to break up the giant
  structure

Question 11

Explain the boiling/melting point and electrical conductivity of METALLIC compounds in terms of their structure

Answer 11

• high BP/MP: strong electrostatic attraction between cations and delocalized electrons, need large energy to break the bond
• conducts electricity in all states: delocalized electron

Question 12

Explain the trend in melting points of metals across the period table

Answer 12

increase across period: increasing number of valence delocalized electrons + decreasing atomic radii (size) → stronger electrostatic attraction force → more energy needed to break the bond

**decrease down the group: increasing atomic radii (size) → further distance between valence delocalized electron and nucleus → weaker electrostatic attraction force → less energy needed to break the bond

Question 13

Describe the structures of diamond, graphite, fullerene

Answer 13

• DIAMOND: each carbon atom is covalently bonded to 4 other atoms in a tetrahedral arrangement; very strong covalent bond, no delocalized electron
• GRAPHITE: each carbon atom is covalently bonded to 3 other atoms in a layered structure (the layers are made of hexagons with a bond angle of 120˚); delocalized electrons occupy the space in between the layers; the layers are held together by weak LDF
• FULLERENE: contains 60 carbon atoms, each of which is covalently bonded to 3 other atoms by single covalent bonds; the 4th electron is delocalized but can’t jump between individual fullerness, so it is a semi-conductor

Question 14

bond polarity is resulted from

Answer 14

difference in electronegativity of the bonded atoms

Question 15

why do some molecule have polar bonds but are non-polar?

Answer 15

the polar bonds in the molecule are arranged in such way that the dipole moments cancel each other → symmetrical shape, no overall dipole moment

Question 16

Explain using diagrams why SO2 is a polar molecule but CO2 is a non-polar molecule.

Answer 16

[pic]
SO2 vector ↓
CO2 vector ←→

Question 17

the octet rule

Answer 17

atoms have the tendency to gain 8 electrons in outer shell = most stable

Question 18

the strength of intermolecular forces depends on:

Answer 18

• Bond Type - H > D-D > LDF
• Mass - more massive = stronger LDF
• Surface Area - larger SA = more likely to be in contact and induce a dipole in adjacent molecules = stronger LDF
• Number of electrons - more electron = stronger LDF
• Size - bigger → valence electrons farther → less tightly held → weaker D-D

Question 19

exceptions of the octect rule

Answer 19

• H is stable with 2 valence electrons
• Be is stable with with 4 valence electrons
• B, Al are stable with 6 valence electrons

Question 20

VSEPR theory

Answer 20

Valance Shell Electron Pair Repulsion:
1. electrons arrange themselves as far away from each other as possible
2. lone pairs repel more strongly than bonding pairs.
3. multiple bonds behave like single bonds

Question 21

POLAR BONDS have

Answer 21

1. Difference in electronegativity
2. Asymmetrical shape (electrons drawn to the more electronegative atom)

Question 22

relative strengths of intermolecular forces

Answer 22

Hydrogen > Dipole-Dipole > London Dispersion Force

Question 23

how do you compare the boiling points of covalent bonds?

Answer 23

BP => bond strength
1. Bond Type - H > D-D > LDF
2. Mass - more massive = stronger LDF
3. Surface Area - larger SA = more likely to be in contact and induce a dipole in adjacent molecules = stronger LDF
4. Number of electrons - more electron = stronger LDF

Question 24

how do you determine the solubility of covalent bonds?

Answer 24

1. like dissolves like: polar substances are soluble in polar solvents and non-polar substances are soluble in non-polar solvents.
   (eg. ethanol and is soluble in water)
2. as molecules get larger, their solubility can decrease as the polar bond becomes a smaller part of their structure
   eg. ethanol, C2H5OH, is soluble in water but (hexanol, C6H13OH, is not)

Question 25

explain the London Dispersion Force

Answer 25

electrons move around randomly → instantaneous uneven distribution of electrons in a molecule
→ instantaneous dipole → induces a dipole in nearby molecules → nearby molecule (with polarized electrons) is attracted to the partial charge

Question 26

explain the Hydrogen Bond

Answer 26

strong dipole-dipole attraction between between molecules that have an electronegative N/O/F atom directly bonded to an H atom

Question 27

why does water have high boiling point?

Answer 27

strong hydrogen bond (H2O)

Question 28

what increases the strength of ionic bond?

Answer 28

• increase in charge density - greater charge or smaller ion

Question 29

what increases the strength of covalent bond?

Answer 29

• decrease in bond length

Question 30

what increases the strength of metallic bond?

Answer 30

• increase in charge density of ion (increase charge or decrease radius)
• increase the number of delocalized electrons

Question 31

alloy

Answer 31

mixtures of metals (metals are mixed together PHYSICALLY but are not chemically combined)

note: it is possible to form alloys because of the non-directional nature of the metallic bonds!

Question 32

Why are alloys HARDER than pure metal?

Answer 32

regular lattice structure is disrupted by the presence of another element → harder for layers to slide over each other