Chap 4A - chemical bonding

Types of bonding and properties

What is a chemical bond

Chemical bonds are electrostatic forces of attraction between particles such as atoms, molecules, ions or electrons

What is electronegativity

• Of an atom is a measure of its ability to attract the electrons in a covalent bond to itself

1. In the hydrogen molecule , the bonding electrons are equally shared between the two hydrogen atoms as they have the same electronegativity

2. A more electronegative atom will attract the shared pair of electrons more strongly towards itself

Describe the periodic trends of electronegativity

1. Across a period 

• Nuclear charge increases while shielding effect remains relatively constant.

• Effective nuclear charge increases, hence electronegativity increases across the period

2. Down a period 

• Elements in the same group experience roughly the same effective nuclear charge as both nuclear charge and shielding effect increase down the group.

• However, as the number of quantum shells increases, atomic radius increases. Hence, electronegativity decreases down the group

Describe the relationship between electronegativity and chemical bonding

• Difference in electronegativity between two atoms can be used to predict the type of bond formed between them
  

1. Similar electronegativity between non metal atoms -> electron sharing -> covalent bonding 

2. Large difference in electronegativity between metal and non metal atoms -> electron transfer -> ionic bonding 

Define metallic bonding

Def. : Metallic bonds are electrostatic forces of attraction between a lattice of positively charged metal ions and the sea of delocalised valence electrons of the metal atoms

Why is metallic bonding non-directional?

Mobile electrons can move freely around the metal -> non-directional bonding (attraction between the metal ions and the delocalised electron cloud occurs in all directions)

Describe the factors of metallic bonding

1. Number of valence electrons per metal atom 

• The larger the number of valence electrons contributed per atom, the greater the number of delocalised electrons -> more extensive would be the forces between the cations and delocalised electrons -> the stronger the metallic bonding

2. Charge and radius of metal cation 

• The higher the charge and the smaller the radius of the metal cation, the higher its charge density -> stronger the attraction between cations and delocalised electrons ->  stronger the metallic bonding

• Charge density 

• Amount of charge per unit surface area of that ion 

• Charge density = ionic charge / ionic radius (ratio) 

Define lattice

Lattice (Def.) : Refers to a regular, repeating three-dimensional arrangement of particles (atoms,ions, or molecules) that form a crystalline solid. In a lattice, the particles are arranged in a way that maximises stability by minimising energy, often following a symmetrical and ordered pattern

Define giant metallic structure

Giant Metallic structure (Def.) : A giant metallic structure is a three-dimensional lattice of closely packed positive metal ions immersed in a "sea" of delocalised electrons. The metallic atoms are held together by strong metallic bonds, formed by the electrostatic attraction between the positive metal ions and the mobile electrons which move freely throughout the lattice

Explain why Mg has a higher temp than Na

1. Each Mg atom has two valence electrons while each Na atom has one valence electron. Hence, magnesium has a larger number of delocalised electrons.

2. Mg2+ (0.065 nm) has a larger charge and smaller ionic radius than Na+ (0.095 nm). Mg2+ has a higher charge density than Na+

3. More energy is needed to overcome the stronger metallic bonding in magnesium than that in sodium

Explain why boiling point is a better indicator of the strength of metallic bonding than the melting point

When a solid metal melts, the closely packed cations become further apart but the metallic bonding is not completely broken, unlike during boiling

Explain why metals are good electrical conductors and malleable

1. Good electrical conductivity

• Contains mobile charge carriers (delocalised electrons) under an applied potential difference 

• Charge carriers : can either be ‘electrons’ or ‘ions’. When they are free to move, they allow electrical current to be conducted in the presence of an electric field

2. Malleable

• When a large stress (e.g. hammering) is applied to a piece of solid metal, the layers of ions will slide over one another into new positions

• The overall shape of the metal changes (the metal deforms) but it does not break apart because the 'sea' of delocalised electrons prevents repulsion between the cations as they move past one another, and metallic bonding remains intact

Define ionic bonding

• Def. : electrostatic forces of attraction between oppositely charged ions which are formed when an atom transfers one or more of its electrons to another atom, resulting in the formation of positive and negative ions

• Ionic bonding is non-directional and extends indefinitely in all directions while a covalent bond is directional as one covalent bond can only be formed between two specific atoms

What is coordination number and state its factors

• Number of ions that surround another ion of the opposite charge in an ionic lattice

• Depends on : 

  • Relative sizes :  a small ion typically has a lower coordination number as there will be less space for ions of the opposite charge to surround it

  • Relative charges : a cation with +2 charge needs twice as many anions with –1 charge, compared to a cation with +1 charge, to gain charge neutrality

Define lattice energy

• Def. : Energy released when one mole of solid ionic compound is formed from its constituent gaseous ions

• Measures strength of ionic bond 

  • Factors : charge and radius of ions 

• Larger the magnitude of lattice energy, the more heat is given out when the ionic bonds are formed, the stronger the ionic bonding

Define giant ionic structure

• (Def.) : Refers to a three-dimensional lattice of alternating positive and negative ions held together by strong electrostatic forces of attraction (ionic bonds)

• Every ion is attracted to oppositely charged ions in multiple directions without a preferred orientation -> ionic bonds are non–directional

Explain why MgO has a higher mp than NaCI

1. Mg2+ has a smaller radius than Na+ while O2- has a smaller radius than Cl- 

2. Mg2+ and O2- have higher charges than Na+ and Cl- , respectively

3. Lattice energy of MgO is larger in magnitude. More energy is needed to overcome the stronger ionic bonding in MgO

Describe the electrical conductivity of ionic compounds

1. In the solid state, the ions are held in fixed positions -> no mobile charge carriers and ionic compounds are unable to conduct electricity

2. In molten or aqueous state, the ions are free to move -> ions can act as mobile charge carriers enabling ionic compounds to conduct electricity under an applied potential difference

Explain why ionic compounds are hard and rigid but brittle

• In an ionic lattice, oppositely charged ions are held in fixed positions throughout the crystal lattice by strong electrostatic forces of attraction between oppositely charged ions 

• Moving the ions out of position requires large amounts of energy to overcome these bonds -> hard (do not dent or deform easily) and rigid (do not bend easily)

• If a strong enough force (by cutting or knocking) is applied, it will force ions of like charges to move next to each other

• Strong repulsion between ions of like charges will cause the lattice to shatter -> brittle (crack without deforming)

Define covalent bonding

(Def.) : Electrostatic forces of attraction between the shared pair of electrons and positively charged nuclei of the two bonded atoms

How many covalent bonds can period 1, 2 and 3 atoms form

1. Period 1 atoms : Cannot expand their octets because they have only the 1s orbital in their valence shell

2. Period 2 atoms : only have 2s and 2p orbitals, atoms lack energetically accessible d-orbitals, they cannot expand their octets beyond 8 electrons -> form maximum 3 bonds 

3. Period 3 and higher atoms : have 3d-orbitals -> d-orbitals allow them to accommodate more than 8 electrons -> enabling them to expand their octet (Eg. PCI5 with 5 covalent bonds) 

Describe polar and non-polar molecules

1. Polar 

• Electrons are shared equally between atoms with same electronegativity
  

2. Non polar

• Electrons are shared unequally between atoms of different electronegativity, leading to a partial positive charge on one atom and a partial negative charge on the other

Define dative bonds

Def. : A dative covalent bond, also known as a coordinate bond, is a type of covalent bond in which both electrons that form the bond come from the same atom

State some characteristics of dative bonds

1. Electron Donor and Acceptor

• Donor (often a species with a lone pair of electrons , Eg. N or O)

• Acceptor (usually an electron-deficient species, Eg. Metal ion or proton)

2. Bond Similar to Covalent Bond

• Once the dative bond is formed, it behaves like a regular covalent bond in terms of strength and properties

• Distinction lies in the origin of the shared electron pair
  

3. Notation

• Represented by an arrow (→) pointing from the donor atom to the acceptor atom

Explain how to identify a dative bond

1. Species that have lone pair 

• Common donors (Lewis bases): NH₃, H₂O, Cl⁻, OH⁻, CN⁻, PCl₃

2. Species that lacks electrons 

• Common acceptors (Lewis acids): H⁺, BF₃, AlCl₃, Fe³⁺, Cu²⁺

• These species have an empty orbital to accept electrons

3. If a new bond forms without breaking an existing one, it is likely a dative bond

Explain why atoms form dative covalent bonds using BF3 and NH3

1. Electron deficient (Eg. BF3 + NH3) 

• Boron atom forms regular covalent bonds with three fluorine atoms, using up all three of its valence electrons

• Leaves the boron atom with an incomplete octet configuration, as it only has six electrons in its outermost shell (electron-deficient molecule)

• Since boron has a vacant orbital (2p) which is energetically accessible, it can accept a lone pair of electrons from another molecule

• When ammonia (NH3) donates a lone pair of electrons from nitrogen to the boron atom, a dative covalent bond is formed

• This bond helps stabilise the BF3 molecule by enabling it to have a stable octet configuration

2. Unable to expand octet due to lack of energetically accessible vacant orbitals 

• N in NH3 can can donate its lone pair of electrons to form a dative covalent bond with an atom that has an empty orbital, like boron in BF3, without violating its octet rule

Describe the 2 types of orbital overlaps

1. Sigma (σ) bond is formed by 'head-on' overlap 

• Occurs along the nuclear axis connecting the two nuclei of the bonding atoms

1. Pi (the math one) is formed by ‘sideway’ overlap

• Can only occur when atoms undergo multiple bonding 

• From two p orbitals above and below the plane of the atoms involved

• Occurs in regions parallel to the nuclear axis, rather than along it

• The electron density of a bond is concentrated above and below the nuclear axis.

• p orbitals can overlap both head-on and side-on ,  s orbitals can only overlap head-on to form sigma bonds

• Side-on overlap for a pi bond is poorer than head-on overlap for a σ bond, as the degree of overlap is smaller for the formed -> a pi bond is weaker than a σ bond -> pi bond continue less to overall bond strength 

• Between two atoms, there can only be one σ bond as it is not possible for a second pair of orbitals to overlap head-on as well

• For a pi bond to form, a σ bond must first be formed

1. A single bond consists of a σ bond

2. A double bond consists of one σ bond and one pi bond

3. A triple bond consists of one σ bond and two pi bonds

Describe bond energy

• Measures strength of covalent bond 

• Average amount of energy required to break 1 mole of a covalent bond in the gaseous state to form gaseous atoms

• Larger the bond energy, the more heat is needed to break the bond, the stronger the covalent bonding

Describe the factors of covalent bonds

1. Bond length

• The distance between the nuclei of the two bonded atoms 

• When atoms have a larger radius, their nuclei are farther apart in a bond -> increased distance reduces the strength of the electrostatic attraction between the nuclei and the shared electrons

• Larger atoms have more diffuse orbitals, leading to less effective orbital overlap -> weaker covalent bonds.

• Exception : The F–F bond is shorter but weaker than the Cl–Cl bond -> due to the repulsion between the lone pairs on the fluorine atoms which are in close proximity

2. Multiplicity of bond

• The greater the number of shared electron pairs, the stronger the electrostatic forces of attraction between the shared pair of electrons and the positively charged nuclei, the stronger the covalent bond

• The C=C bond consists of a sigma (σ) bond and a pi (π) bond -> σ bond has a greater degree of orbital overlap and is more effective than the π bond -> σ bond is stronger -> bond energy of a C=C bond is less than twice the bond energy of a C–C bond

3. Polarity of bond

• For covalent bonds with the same multiplicity involving atoms of comparable radii

• Polar bonds are formed by atoms of different electronegativity values

• The additional electrostatic attraction between the two partial charges causes the covalent bond to be stronger -> polar covalent bond is stronger than a non–polar covalent bond

• Increase in polarity of the covalent bond caused by greater difference in electronegativity values will increase the covalent bond strength

4. Proximity of lone pair electrons

• There can be significant repulsion between lone pairs of electrons on the two bonded atoms close together -> covalent bond to be weakened

• The bond energy of F–F is lower than that of Cl–Cl, even though F–F has a shorter bond length due to fluorine’s smaller atomic radius -> short bond brings the lone pairs of electrons on both fluorine atoms very close to each other, resulting in strong repulsive forces that weaken the bond

Describe Van der Waals radius

• When we measure the distances between the nuclei of two bonded iodine atoms in iodine solid, we obtain two different values

• Shorter distance : between two bonded atoms in the same molecule (bond length) 

  • Half of this distance : covalent radius (referred to as ‘single covalent’ in the Data Booklet)

• Longer distance : between two non-bonded atoms in adjacent molecules

  • Half of this distance : van der Waals radius

  • The van der Waals radius of an atom is always larger than its covalent radius. 

  • In the Data Booklet, the radii listed for the noble gases are the van der Waals radii since they do not form compounds

Define simple molecular lattice

(Def.) : A simple molecular structure consists of simple discrete molecules held together by weak intermolecular forces. These molecules consist of a small number of atoms joined by covalent bonds within the molecule itself

Explain why argon has a lower boiling point than hydrogen chloride

Argon has a simple atomic structure while hydrogen chloride has a simple molecular structure. Stronger permanent dipole-permanent dipole attractions (and dispersion forces) exist between polar HCl molecules.  Less energy is needed to overcome the weaker dispersion forces between argon atoms so it has a lower boiling point

Explain why simple molecules have a poor electrical conductivity

• Poor electrical conductors as there are no mobile charge carriers to carry electrical charges

• Some simple molecular substances can conduct electricity in aqueous state due to formation of mobile ions when the molecules interact with water (Eg. HCI) 

• The ions act as mobile charge carriers allowing electricity to be conducted under an applied potential difference

Describe the solubility and softness of simple molecules

Solubility	Soluble in organic solvents  Some soluble in water due to their ability to form hydrogen bonds with water
Softness	Very soft because their lattice structures are held by weak intermolecular forces of attraction

Define giant molecular lattice

(Def.) : A giant molecular structure consists of a large number of atoms joined by strong covalent bonds in a continuous, three-dimensional network

Describe the structure of diamond

• Each carbon atom is covalently bonded to four other carbon atoms in a tetrahedral arrangement 

• Arrangement is repeated throughout the whole molecule, giving rise to an extensive 3D network of covalent bonds

Describe the structure of graphite

• Carbon atoms are arranged in layers made up of hexagonal rings of carbon atoms

• The carbon atom is covalently bonded to three other carbon atoms to form an extensive network of planar hexagonal rings

• Each carbon atom contributes a valence electron to form a delocalised electron cloud above and below the plane of carbon atoms.

• The layers are stacked on top of each other held together by weak instantaneous dipole–induced dipole attractions

Describe the electrical conductivity of diamond and graphite

Electrical conductivity

Poor electrical conductors as there are no mobile charge carriers to carry electrical charges

Good electrical conductors in the direction parallel to the layers as delocalised electrons can act as mobile charge carriers under an applied potential difference. They are poor electrical conductors in the direction perpendicular to the layers as delocalised electrons cannot move across the layers

Describe the solubility of diamond and graphite in water

Generally insoluble in water and organic solvents because it consists of uncharged atoms so partial opposite charges not attracted to atoms

Describe the hardness and uses of diamond and graphite

Hardness	Very hard , strong and non-malleable due to extensive network of strong covalent bonds between carbon atoms (rigid structure)		Flaky and slippery because the weak forces of attractions (instantaneous dipole–induced dipole attractions) between the layers enable them to slide over one another easily, giving graphite its slippery feel
Uses	Diamond : used as drilling or cutting tool  Quartz : plays a crucial role in glass and electronics		Used as lubricant for hot machines , pencil leads

Describe the trend in melting point across period 3

1. From Na to Al, melting point increases as more energy is needed to overcome the stronger metallic bonding due to :

• Increase in number of valence electrons contributed per atom

• Increase in charge density of the cation (due to increasing charge and decreasing ionic radius/cationic size)

2. Si has a giant molecular structure and a very high melting point as a lot of energy is needed to overcome the many strong covalent bonds between the Si atoms during melting

3. Non-polar P4, S8, and Cl2 have simple molecular structures

4. Ar has simple atomic structure

• Little energy is needed to overcome the weak dispersion forces between the molecules or between argon atoms hence the low melting points

• The number of electrons and hence strength of dispersion forces decreases from S8 to P4 to Cl2 to Ar. Hence, energy needed to overcome the dispersion forces during melting and the resulting melting point decreases in the same order

Describe the trend in electrical conductivity across period 3

1. Na, Mg and Al have giant metallic structures with sea of delocalised valence electrons to act as mobile charge carriers

1. From Na to Al, the number of valence electrons contributed per atom increases. Therefore, electrical conductivity increases
   

2. Si has a giant molecular structure with covalent bonds between atoms. As a metalloid, it is a semiconductor
   

3. P4, S8, and Cl2 have simple molecular structures. Ar has a simple atomic structure. 

• In P4, S8 and Cl2, the valence electrons are held strongly in covalent bonds -> no mobile charge carriers in these molecules as well as in argon atoms so they do not conduct electricity