CBI1 - Fundamentals of Chemistry I

Question 1

Four quantum numbers and their notations.

Answer 1

Principal quantum number (n).
Orbital or azimuthal quantum number (l).
Magnetic quantum number. (ml).
Spin quantum number (ms).

Question 2

Describe Schrodinger’s wave equation and it’s relevance to electrons.

Answer 2

Electrons have wave-like properties whose behaviour can be described using the Schrodinger equation. Solutions of the equation are known as wave functions. Defines the regions where electrons can be found.

Question 3

Discuss electronegativity trends and definition.

Answer 3

Ability to withdraw electrons towards itself. Electronegativity increases across period due to increasing nuclear charge. Electronegativity decreases down group due to increasing number of shells.

Question 4

Define electron affinity.

Answer 4

Ability to attract an electron.

Question 5

Discuss atomic radius trend.

Answer 5

Atomic radius decreases across a period due to increasing nuclear charge. Atomic radius increases down a group due to increasing shielding.

Question 6

Define and give examples of isotopes.

Answer 6

Atoms of the same element with differing mass numbers. Chemical properties unchanged, physical properties altered. E.g carbon-12 and carbon-13.

Question 7

Discuss subatomic particles of an atom.

Answer 7

Protons - positively charged, located in nucleus, atomic number (Z).
Neutrons - neutral, located in nucleus.
Electrons - negatively charged, smallest subatomic particle. Occupy regions around nucleus.

Question 8

Discuss ionisation energy definition and trend.

Answer 8

Energy required to remove one electron from an atom. Ionisation energy (generally) increases across period due to increasing nuclear charge and decreasing atomic radius. Decreases down group due to increasing atomic radius.

Question 9

Define principal quantum number.

Answer 9

Species the main shell. Larger value of n means orbital is larger, of higher energy and less tightly bound to nucleus.

Question 10

Define orbital quantum number.

Answer 10

Also known as azimuthal meaning shape. Defines shape of orbital.
S orbital: l=0
P orbital: l=1
D orbital: l=2
F orbital: l=3

Question 11

Define magnetic quantum number.

Answer 11

Determines possible orientations of the energy degenerate orbitals. 
S orbitals: 0 
P orbitals: -1, 0, +1
D orbitals: -2, -1, 0, +1, +2
F orbitals: -3, -2, -1, 0, +1, +2, +3

Question 12

Define spin quantum number.

Answer 12

Two electrons in one orbital which either spin up or down.
Spin up: +1/2
Spin down: -1/2

Question 13

Define Rutherford-Bohr model of an atom.

Answer 13

Positively charged nucleus consisting of protons and neutrons surrounded by electrons occupying regions of space around it.

Question 14

Define Aufbau principle.

Answer 14

Orbitals must occupy orbitals of lower energy first.

Question 15

Define Hunds rule.

Answer 15

Electrons fill degenerate energy orbitals singly first and then pair up to minimise repulsion. Unpaired electrons will have the same spin. Paired electrons will have opposing spin quantum numbers.

Question 16

Define Pauli exclusion principle.

Answer 16

No two electrons can have the same set of four quantum numbers.

Question 17

Define Madelung rule.

Answer 17

Rule of how electrons are filled by writing names of orbitals and drawing diagonal lines.

Question 18

What order to electrons fill when doing electron configuration.

Answer 18

1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d…

Question 19

Two main types of bonding and their differences.

Answer 19

Covalent bonding - shared pair of electrons where difference in electronegativity is less than 2.
Ionic bonding - electrostatic force of attraction between oppositely charged ions (as a result of transfer of electrons) where difference in electronegativity is greater than 2.

Question 20

Discuss polar covalent bonds.

Answer 20

Unequal distribution of electron cloud due to differing electronegativity of two atoms in a bond.

Question 21

Define and discuss VSEPR.

Answer 21

Valence shell electron pair repulsion theory. Allows prediction of geometry of a molecule around a central atom. Considers only outer electrons and takes into account lone pairs and bonding pairs.

Question 22

Order of repulsion for lone pairs and bonding pairs.

Answer 22

Lone pair-lone pair > lone pair-bonding pair > bonding pair-bonding pair.

Question 23

By how many degrees does a lone pair reduce the bond angle.

Answer 23

Between 2.5-3.

Question 24

Define an atom.

Answer 24

Smallest particle of an element.

Question 25

What are A and Z

Answer 25

Z = atomic number (number of protons) 
A= atomic mass (number of protons and neutrons)

Question 26

What is a Dalton or unified atomic mass unit.

Answer 26

Equal to 1/12 of a carbon 12 atom.

Question 27

Define groups and periods.

Answer 27

Groups = columns in periodic table. 
Periods = rows in periodic table.

Question 28

Is 14C stable or unstable.

Answer 28

Unstable as it is radioactive.

Question 29

Define nuclide symbols.

Answer 29

Notation present on the periodic table.

Question 30

What are wave functions.

Answer 30

Solutions to Schrodinger’s equation describing probability density of an electron.

Question 31

Describe orbital shapes.

Answer 31

S = sphere
P = drop shaped lobes on opposing sides 
D = lobed 
F = multiple sets of lobes

Question 32

What are the two exceptions when doing electron configuration. Why.

Answer 32

Cooper and chromium.

Increased stability with half and completely filled sub shells.

Question 33

Discuss covalent bonding with polarity.

Answer 33

Shared paired of electrons between two non metals. Difference in electronegativity less than 2, greater than 0.5. uneven distribution of electron cloud.

Question 34

Molecular geometry for 2-6 electron pairs and bond angle.

Answer 34

2 - linear - 180
3 - trigonal planar - 120
4 - tetrahedral - 109.5
5 - trigonal bipyramidal - 90 axial, 120 equatorial 
6 - octahedral - 90

Question 35

Hybridisation in double and triple carbon bonds.

Answer 35

Double bonds - sp2 hybridised with one remaining p orbital.

Triple bonds - sp hybridised with two remaining p orbitals.

Question 36

Define sigma bond.

Answer 36

End to end overlap of s, p, d of f orbitals.

Question 37

Define pi bond.

Answer 37

Parallel/sideways overlap of p/d/f orbitals.

Question 38

Why is there restricted rotation around C=C

Answer 38

Bond is planar so it is energetically unfavourable for rotation to occur.

Question 39

Define difference between overlap and hybridisation.

Answer 39

Overlap - atomic orbitals overlap and region of overlap is where electrons are found.
Hybridisation - atomic orbitals hybridised to form new molecular orbitals with some character from each atomic orbital that was hybridised.
Both from valence bond theory.

Question 40

Discuss what happens to electrons when hybridisation occurs,

Answer 40

Electrons are promoted or excited.

Question 41

Define molecular orbital theory.

Answer 41

Quantum mechanic based model of molecular bonding that describe molecular orbitals as regions of space where electrons are likely to be found in molecules.