Chemical Bonding Flashcards
What is the definition of electronegativity?
The ability of an atom to attract (a) pair(s) of electrons in a covalent bond
What are the factors affecting electronegativity?
- Nuclear charge - a higher nuclear charge increases the electrostatic force of attraction between the nucleus and bonding pairs
- Shielding - having fewer inner electrons in shells and subshells means the bonding pair is less shielded from the nucleus so will be more attracted to it
- Atomic radius - a lower atomic radius decreases the distance between the nucleus and bonding pairs
- The ability to form bonds; for example, noble gases do not have electronegativity values at all as their complete octet means they do not gain or lose electrons
Shielding and nuclear charge affect atomic radius, so these factors are interconnected
What is the trend in electronegativity down a group and why does this trend exist?
- It decreases, despite atomic charge increasing
- This is because the number of principal quantum shells increases, causing there to be more shielding between the nucleus and any potential bonding pairs
- The increased number of shells also results in an increased atomic radius
What is the trend in electronegativity across a period and why does this trend exist?
- It increases, despite the number of valence electrons increasing
- This is because nuclear charge is increasing, causing there to be a greater electrostatic force of attraction between the nucleus and any bonding electrons
- There is no substantial increase in shielding as no new shells are formed
- The increased nuclear charge also results in a decreased atomic radius
What does the difference in electronegativity on the Pauling electronegativity scale need to be for bonds to be covalent, polar covalent and ionic?
- Covalent: >1.0
- Polar covalent: 1.0-2.0
- Ionic: >2.0
Bonds will only be completely non-polar if the two atoms bonding are of the same element
What is an ionic bond?
The electrostatic attraction between oppositely charged ions (positively charged
cations and negatively charged anions)
Ionic compounds form giant, regular, crystalline lattices
What is metallic bonding?
The electrostatic attraction between positive metal ions and the delocalised valence electrons of these metal ions in a giant metallic lattice
What is a covalent bond?
An electrostatic attraction between the nuclei of two atoms and a shared pair or pairs of electrons
What is it called when a bonded atom has more than eight electrons and what is it called when a bonded atom has fewer than eight electrons?
- More than eight - Expanded octet
- Fewer than eight - Electron deficient (with the exception of hydrogen)
What are some examples of molecules that contain atoms with an expanded octet?
- Mainly compounds with period 3 or higher elements like sulfur and phosphorus as they can use 3d orbitals in bonding
- SO₂ - Each oxygen forms a double bond with sulfur, leaving it with 10 valence electrons
- PCl₅ - Each chlorine forms a single covalent bond with phosphorus, leaving it with 10 valence electrons
- SF₆ - Each fluorine atom forms a single covalent bond with sulfur, leaving it with 12 valence electrons
Do not worry about how 3d orbitals are used in bonding
What is a dative or coordinate bond?
- The covalent bond formed when one atom donates a lone pair of electrons to an electron-deficient atom
- The two atoms share the two electrons as they would in a regular covalent bond
The movement of electrons pairs is denoted by arrows
How is an ammonium ion formed?
- The lone pair on the nitrogen in an ammonia molecule is donated to the empty orbital of a H⁺ ion
- A dative covalent bond forms between the hydrogen and nitrogen
- As a proton was gained by the molecule but no new electrons were added, the molecule adopts a 1+ charge
How is an aluminium chloride dimer formed?
- At high temperatures, aluminium chloride can exist as an electron-deficient monomer (AlCl₃); aluminium needs two more electrons to complete its octet
- Two AlCl₃ molecules come into contact with one another
- A lone pair of electrons on a chlorine atom in one monomer is donated to the aluminium in the other monomer; the same happens for the other monomer
- Al₂Cl₆ is formed with two coordinate bonds
What is a sigma (σ) bond?
- The direct overlap of orbitals containing single, unpaired electrons
- S or sp hybridised orbitals can form sigma bonds
- The result is a molecular orbital
- A single covalent bond is always a sigma bond
What is a pi (π) bond?
- The sideways overlap of adjacent unhybridised p orbitals
- They will be the additional bonds in double and triple bonds
- The interaction between the two p orbitals to form a molecular orbital occurs above and below the plane of the sigma bond
What is the process of forming hybridised orbitals?
- Only valence orbitals form hybridised orbitals
- S and p orbitals will combine to form hybridised orbitals
- One s orbital and one p orbital form two sp hybridised orbitals
- One s orbital and two p orbitals form three sp² hybridised orbitals
- One s orbital and three p orbitals form three sp³ hybridised orbitals
How is hybridisation used in bonding?
- Whenever multiple single bonds or valence lone pairs exist in a molecule, hybridised orbitals must form
- This is because p orbitals cannot form sigma bonds, but they can when given some s character
- The hybridised orbitals are then used to form sigma bonds or hold lone pairs
Pure p orbitals are able to hold lone pairs too
How is hybridisation used in N₂?
- In each nitrogen atom, the s and a p orbital combine to form two sp hybridised orbitals
- One is used to form the sigma bond with the other nitrogen, the other holds the valence lone pair
- The remaining two p orbitals form two pi bonds with the other nitrogen atom
How is hybridisation used in ethene?
- In each carbon, the s and two p orbitals will combine to form three sp² hybridised orbitals
- Two of these will be used to form sigma bonds with the s orbitals of the two hydrogen atoms and one of them will be used to form a sigma bond with the other carbon
- The remaining p orbital forms a pi bond with the other carbon
What is the definition of bond energy?
The energy required to break one mole of a particular covalent bond in the gaseous state
What is the definition of bond length?
The internuclear distance of two covalently bonded atoms
How is bond length correlated with bond energy and strength?
- As bond length increases, bond energy and bond strength decrease
- This is because bond length is determined by the electrostatic attraction of the nuclei to the bonding pair(s), so a higher bond length means a weaker attraction and less orbital overlap
- Triple bonds have the smallest bond length (as three molecular orbitals are formed), so the highest bond energy and strength
How are bond length and bond strength correlated with the reactivity of a compound?
- As bond length increases, reactivity increases
- As bond strength increases, reactivity decreases
- This is because the bond needs to be broken for the atoms to react, so a compound with a bond that breaks more readily will be more reactive
What are the rules of VSEPR theory used to determine the arrangement of electrons around a central atom?
- Only valence electrons are involved in determining shape
- The most stable shape is adopted to minimise electron-electron repulsion
- Lone pairs are more repulsive than bonding pairs (because the electron density is higher as they are not being pulled apart by two nuclei)
- Pairs of electrons or bonds are considered single areas of electron density
- Double and triple bonds are treated the same as single bonds
What is the name of the shape of a molecule with two bonding areas of electron density and what is the angle between them?
- Linear
- 180°
Carbon dioxide is an example of such a linear molecule as even though it has four bonding pairs, the double bonds are treated as single bonds
What is the name of the shape of a molecule with three bonding areas of electron density and what is the angle between them?
- Trigonal planar
- 120°
What is the name of the shape of a molecule with two bonding areas of electron density and one lone pair and what is the angle between them?
- Bent/angular
- ~117.5°
- The bond angle decreases by roughly 2.5° for every lone pair substituted for a bonding pair
- Only the angles between bonding pairs (so bonds) are considered, not the angles between lone pairs
What is the name of the shape of a molecule with four bonding areas of electron density and what is the angle between them?
- Tetrahedral
- 109.5°
The wedged line denotes a bond coming out of the plane towards you and the dashed line denotes a bond going into the plane away from you
What is the name of the shape of a molecule with three bonding areas of electron density and one lone pair and what is the angle between them?
- Trigonal pyramidal
- ~107°
What is the name of the shape of a molecule with two bonding areas of electron density and two lone pairs and what is the angle between them?
- Bent/angular
- ~104.5°
Water is an example of such a molecule
What is the name of the shape of a molecule with five bonding areas of electron density and what is the angle between them?
- Trigonal bipyramidal
- 90°/120°
- Here we have different angles depending on which bonding pairs the angle is being measured between
- Anything with five or more areas of electron density has an expanded octet
What is the name of the shape of a molecule with four bonding areas of electron density and one lone pair and what is the angle between them?
- Sawhorse
- ~87.5°/~117.5°
- The angle between electrons in the axial plane and equatorial plane is smaller than the angle between electrons in the equatorial plane
- This arrangement therefore minimises repulsion as the number of 90° bonds is minimised; this rule, as well as the axial/equatorial arrangement, holds for molecules with a greter number of areas of electron density
What is the name of the shape of a molecule with three bonding areas of electron density and two lone pairs and what is the angle between them?
- T-shaped
- 90°
What is the name of the shape of a molecule with two bonding areas of electron density and three lone pairs and what is the angle between them?
- Linear
- 180°
What is the name of the shape of a molecule with six bonding areas of electron density and what is the angle between them?
- Octahedral
- 90°
What is the name of the shape of a molecule with five bonding areas of electron density and one lone pair and what is the angle between them?
- Square pyramidal
- ~87.5°
What is the name of the shape of a molecule with four bonding areas of electron density and two lone pairs and what is the angle between them?
- Square planar
- 90°
What is the name of the shape of a molecule with three bonding areas of electron density and three lone pairs and what is the angle between them?
- T-shaped
- 90°
What is the name of the shape of a molecule with two bonding areas of electron density and four lone pairs and what is the angle between them?
- Linear
- 180°
What is a dipole?
- A molecule with a partially positively charged component and a partially negatively charged component
- Caused by a difference in electronegativity in its constitutent atoms
- Partial positive charges are denoted by delta positive ( δ+) and partial negative charges are denoted by delta negative (δ-)
A dipole moment is a measure of the polarity or separation of charge
When might a molecule not be a dipole even when there are differences in electronegativity between its atoms?
- If the molecule is perfectly symmetric
- Carbon dioxide, with a linear shape, is a key example of this
- This is because the partial charges completely cancel each other out
- Molecules like methane where the difference in electronegativity is inconsequential are also generally not considered dipoles
These molecules must therefore rely on London dispersion forces
What are the four primary types of intermolecular forces in order of increasing strength?
- Instantaneous dipole-induced dipole (London dispersion) forces
- Permanent dipole-induced dipole forces
- Permanent dipole-permanent dipole forces
- Hydrogen bonds
- They are all weaker than ionic, covalent and metallic bonds
- Van der Waals’ forces is a generic term for all intermolecular forces
What is hydrogen bonding?
- A type of permanent dipole-permanent dipole attraction between a partially positively charged hydrogen and a partially negatively charged electronegative atom
- This electronegative atom must be oxygen, nitrogen or fluorine and must be bonded to a hydrogen
- The hydrogen itself must be bonded to one of these highly electronegative atoms
- Hydrogen bonds occur between molecules and are strong because the partial positive charges and partial negative charges are high
Hydrogen bonding confers water with high relative melting and boiling points, a high surface tension and a higher density than ice
What are permanent dipole-permanent dipole (pd-pd) forces?
- The electrostatic forces of attraction between δ+ ends of permanent dipoles and δ- ends of neighbouring permanent dipoles
- An example is the attraction between the oxygen atom in one molecule of a ketone and the carbon atom bonded to an oxygen atom in a neighbouring ketone
What are permanent dipole-induced dipole (pd-id) forces?
- The electrostatic forces of attraction between the partially charged atoms of a permanent dipole and the (opposite) partially charged atoms of a nearby molecule this permanent dipole has induced a dipole moment in
- An example is the interaction between hydrogen chloride and argon
What are instantaneous dipole-induced dipole (id-id) forces?
Also known as London disperion forces
- The electron clouds around atoms are constantly moving
- This causes temporary dipoles to arise around atoms, which can induce dipoles in nearby molecules
- As a consequence, there can be temporary electrostatic attractions between the instantaneous dipole and the induced dipole
- These attractions are constantly being created and lost and the dipole moment is low, making London dispersion forces very weak
- This is the only type of intermolecular force that is present in all molecules, even ones with atoms with no difference in electronegativity
- Id-id forces increase in strength with increasing total number of electrons as more and stronger instantaneous dipoles arise
