Chemistry Video 7

Question 1

Ionic bonds

Answer 1

Cations and anions interact via ionic bonds. Strong electrostatic interactions that allow for the formation of ionic compounds.

Occurs when the difference in electronegativity is greater than ~1.8 and electron transfer will occur

Between metal and non-metal

Question 2

Lattice energy

Answer 2

Ionic bonds; Anions and cations arranged in a. grid. Anions are arranged in a geometry that will allow for the maximization of electrostatic interactions between oppositely charged particles.

Stronger interactions = greater lattice energy = greater energy needed to disrupt the lattice and separate all the ions

Question 3

2 factors that influence lattice energy

Answer 3

1. Magnitude of charge on the ions. Greater magnitude of the charge = greater lattice energy.
2. Ion size (ionic radius). Smaller radius = Closer nuclei = stronger lattice attraction

Question 4

Covalent bonds

Answer 4

Share electrons between 2 atoms. Can be between atoms of the same element or atoms of different elements that do not have large difference in electronegativity.

Covalent compounds are neutral in charge.

Each electron feels the pull of its own nucleus and the pull of the other nucleus

Between 2 non-metals

Question 5

Ideal internuclear distance

Answer 5

The ideal covalent bond length between 2 atoms. The distance between the 2 protons of 2 atoms. There is a dip down in energy because if each electron is interacting with both protons, then it is a stabilizing and energetically favourable situation.

If the atoms are pushed too close together, it becomes energetically unfavourable because of proton-proton interaction

i.e. 1s orbitals of 2 atoms overlap until ideal bond length is formed.

Question 6

Nonpolar covalent bond

Answer 6

2 electrons shared precisely evenly between 2 atoms due to an electronegativity difference of less than 0.4

Question 7

Polar covalent bond

Answer 7

2 electrons shared unevenly between 2 atoms due to an electronegativity difference that is NOT 0

Partial positive and partial negative charges exist. These partial charges are not formal charges

EN difference of 0.4 to 1.8

Question 8

Dipole arrow

Answer 8

Points in the direction of excess electron density. Larger arrow indicates more polarity

Question 9

Metallic bond

Answer 9

Pure metals. Make bonds amongst itself.

The metal is positively charged metal cations, with the valence electrons delocalized. The atoms are regarded as being of neutral charge. Allows metals to conduct electricity

Question 10

Lewis dot structures

Answer 10

For individual atoms. Chemical symbol in centre, surrounded by valence electrons represented by dots

Add or delete electron dots to show ions and add the - or + symbol to the top right corner

Can be used for ionic or covalent compounds

Covalent bonds are represented as a line, instead of 2 electron dots. Lone pairs remain as dots.

Question 11

Octet rule

Answer 11

Certain elements (C, N, O) prefer to be surrounded by 8 electrons.

However, some elements with access to d orbitals can expand their octets and have more bonds.

Question 12

Violations of octet rule

Answer 12

Boron and aluminum have 3 valence electrons and can only form 3 covalent bonds.

Hypervalent elements are phosphorus that make 5 bonds and sulfur that make 6 bonds. This occurs because these elements can access d orbitals.

Xenon can make covalent bonds with select elements, such as with 4 Fluorine and 2 sets of lone pairs

Question 13

VSEPR Theory

Answer 13

Valence shell electron pair repulsion. Electron domains will repel each other.

Question 14

Linear electron domain geometry

Answer 14

2 electron domains 180 degrees apart. The central atom is sp hybridized

Question 15

Trigonal planar electron domain geometry

Answer 15

3 electron domains 120 degrees apart. The central atom is sp^2 hybridized, which means 1 s orbital and 2 p orbitals

Question 16

Tetrahedral electron domain geometry

Answer 16

4 electron domains 109.5 degrees apart. The central atom is sp^3 hybridized, which means 1 s orbital and 3 p orbitals

Question 17

Trigonal bipyramidal electron domain geometry

Answer 17

5 electron domains has both 90 and 120 degrees apart for bond angles. The central atom is sp^3d hybridized, which means 1 s orbital, 3 p orbitals and 1 d orbital

Axial and equatorial positions are not the same. Lone pairs prefer equatorial positions (120 degrees apart)

Question 18

Octahedral electron domain geometry

Answer 18

6 electron domains 90 degrees apart for bond angles. The central atom is sp^3d^2 hybridized, which means 1 s orbital, 3 p orbitals and 2 d orbitals

Question 19

Electron domain geometry

Answer 19

Describes atoms and lone pairs

Question 20

Molecular shape

Answer 20

Describes atoms only

Question 21

If molecule has 0 lone pairs

Answer 21

Electron domain geometry = molecular shape

Question 22

If molecule has some lone pairs

Answer 22

Electron domain geometry IS NOT EQUAL to molecular shape

Question 23

H2O angle

Answer 23

Angle is 104.5 degrees.

Question 24

Methane CH4 angle

Answer 24

Angle is 109.5 degrees

Question 25

Ammonia NH3 angle

Answer 25

Angle is 107 degrees

Question 26

Net dipole moment

Answer 26

Add up all the bond dipole moment vectors within a molecule.

If there is net dipole, the molecule is polar. If there is no net dipole, the molecule is nonpolar overall, but may have polar bonds

Geometry is a factor.

Question 27

Polarity influences…

Answer 27

Melting point, boiling point, solubility

Question 28

Atomic vs hybrid orbitals

Answer 28

Atomic orbitals are like 2s, 2p.

Hybrid orbitals are like sp, sp^2, sp^3 etc. which are equal in energy to each other and lower energy than p orbitals.

sp (2 electron domains) combines s and px. sp^2 (3 electron domains) combines s, px and py. sp^3 (4 electron domains) combines s, px, py and pz.

sp^3d is 5 electron domains and combines s, px, py, pz and 1 d.

sp^3d^2 is 6 electron domains and combines s, px, py, pz and 2 d.

Question 29

Sigma bond

Answer 29

Direct overlap of covalent orbitals on the internuclear axis.

Question 30

Pi bond

Answer 30

Lateral overlap of covalent orbitals perpendicular to the internuclear axis. Weaker than sigma bond. But adding pi bonds increases the overall bond strength.

Question 31

Node

Answer 31

The internuclear axis, no probability of electron density. In pi bonds

Question 32

Single, double, triple bond

Answer 32

Single = 1 sigma
Double = 1 sigma, 1 pi
Triple = 1 sigma, 2 pi

Question 33

Bond strength

Answer 33

Triple bond stronger than double bond stronger than single bond

Question 34

Resonance structures

Answer 34

Multiple ways of distributing pi electrons (unhybridized electrons in a pi orbital (AKA pi bonds) or lone pairs) within a molecule. The pi electron density is delocalized. The molecule does not flip between the different resonance structures

i.e. partial double bond character

Question 35

Formal charge

Answer 35

valence electrons normally - # valence electrons around the atom in the current molecule

Question 36

Composite resonance structure

Answer 36

[ ] Box brackets around entire molecule with dotted lines to show the partial (i.e. double) bond characters.

Partial charge (+ or -) is shown on the top right outside of the box brackets surrounding the molecule

Question 37

Comparing resonance structure stability

Answer 37

1. Fewer formal charges = more stable
2. Formal charges are better on specific atoms. More electropositive = better for positive charge. (C+ > N+ > O+). More electronegative = better for negative charge. (O- > N- > C-).
3. More full octets = more favourable. Avoid lack of octet on C/N/O.

Question 38

Molecular orbital theory

Answer 38

Combines atomic orbitals to make molecular orbitals that are delocalized around the entire molecule

Question 39

Valence bond theory vs molecular orbital theory

Answer 39

Valence bond theory: bonds are localized between 2 atoms, bonds are overlapping atomic/hybrid orbitals, generates sigma/pi bonds, predicts molecular shape

Molecular orbital theory: electrons are delocalized throughout the molecule, combines atomic orbitals to form molecular orbitals, bonding/antibonding interactions, predicts arrangement of electrons

Question 40

Constructive interference

Answer 40

Increase amplitude of waves.

Occurs when lobes of the same phase overlap.

Electrons are 3D waves that combine with in-phase waves to make a region with high probability of electron density.

Lower energy, sigma bonding orbital. Or pi bonding orbital with the node on the internuclear axis.

Electrons occupy the lowest energy orbital, thus bonding orbitals are filled first.

Question 41

Destructive interference

Answer 41

Decrease amplitude of waves

Occurs when lobes of opposite phase overlap.

Electrons are are 3D waves that combine with out-of-phase waves to make a region with no probability of electron density. Which is called nodes.

Higher energy, sigma star () antibonding orbital that has a node between the nuclei. OR pi star () antibonding orbital that has a node between the nuclei, which means that the electrons cannot be shared between them.

Question 42

Molecular orbital diagrams

Answer 42

Atomic orbitals of each atom are shown on either side of the molecular orbital diagram. Fill according to Aufbau Principle and Hund’s Rule.

Question 43

Calculating bond order

Answer 43

Bonding orbitals increase bond strength. Antibonding orbitals decrease bond strength.

Bond order dictates the number of pairs of bonding electrons between 2 atoms.

Bond order of 1 means single bond. Bond order of 2 means double bond. Bond order of 3 means triple bond.

Bond order = (# of bonding electrons - # of antibonding electrons) / 2

Higher bond order means more bonds and a stronger bond

If the value of the bond order is between 2 integers (AKA a decimal value), it indicates the presence of partial pi bonds. i.e. resonance.

If bond order is 0, it means that a stable bond between the 2 atoms will not form