Semester 1 (spring 2024) final exam

Question 1

full and condensed electron configuration for silver

Answer 1

1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d9, OR [Kr] 5s2 4d9

Question 2

aufbau principle exceptions

Answer 2

chromium and copper

Question 3

chromium condesned electron configuration

Answer 3

[Ar] 3d5 4s1

Question 4

copper condensed electron configuration

Answer 4

[Ar] 4s1 3d10

Question 5

convergence limit for hydrogen is 3.28 × 10^15 s^−1. what is the first ionisation energy in kJ/mol?

Answer 5

using E=hf, E is equal to 2.17 x 10^-18 J (for a single hydrogen atom, in J). This is 1310 kJ/mol.

Question 6

what can you determine from a graph showing the successive ionisation energies of an element? (3)

Answer 6

1. how many energy levels it occupies: this is equal to the number of jumps/large increases on the graph
2. which group does the element belong to: how many electrons are removed before the first big increase? (count all electrons, including the one right at the bottom of the increase). This number is equal to the number of valence electrons, and therefore the group number.
3. what is the element: x-axis should be the number of ionisations. This is also the number of electrons, which is equal to the number of protons (z, atomic number) for a neutral atom.

Question 7

electromagnetic spectrum: list from highest to lowest frequency (7 wave types)

Answer 7

gamma rays, x rays, ultraviolet, visible, infrared, microwave, radio

Question 8

what is the max number of orbitals in each energy level (its a formula)

Answer 8

n^2

Question 9

how to tell if a bond is polar or nonpolar

Answer 9

if the difference in electronegativity is greater than 0.4, it is a polar bond

Question 10

sigma bond definition

Answer 10

strongest type of covalent bond. formed by head-on overlap between orbitals

Question 11

pi bond definition

Answer 11

covalent bonds, do not form single bonds (only double/triple), formed by the overlap of p orbitals on adjacent atoms, perpendicular to any sigma bonds between those atoms

Question 12

order of intermolecular forces (3) from weakest to strongest

Answer 12

LDF<dipole-diple<hydrogen bonding

Question 13

larger ionic radius: weaker or stronger bond?

Answer 13

weaker

Question 14

greater charge: stronger or weaker bond?

Answer 14

stronger

Question 15

coordination bond definition

Answer 15

covalent bond formed between two atoms where both electrons in the pair being shared have been provided by the same atom (same properteis are regular covalent bond)

Question 16

inter vs intra molecular

Answer 16

inter - solid, liquid, gas (ex. melting overcomes intermolecular forces - forces BETWEEN molecules
intra - bonds that hold the atoms together WITHIN the molecule

Question 17

london dispersion forces

Answer 17

the random and continuous movement of electrons creates asymmetrical electron distribution, which can attract nearby atoms/molecules (always present)

Question 18

dipole-dipole

Answer 18

Only exists for two asymmetrical molecules (those with permanent dipoles). Much stronger than LDF.

Question 19

what creates a dipole

Answer 19

If polar bonds aren’t arranged symmetrically, the molecule will have a permanent dipole

Question 20

Dipole-induced-dipole

Answer 20

for a very polar and nonpolar molecule. partial charges on polar molecule affects distribution of electron density on the other molecule, causing a TEMPORARY dipole formed on the non-polar molecule. (goes away when they aren’t next to each other)

Question 21

what affects strength of dipole-induced-dipole force

Answer 21

large molecules are more susceptible than smaller ones

Question 22

requirements for hydrogen bonding (2)

Answer 22

1. hydrogen bonded to a FON element (fluorine, oxygen, nitrogen)
2. The FON element must have a non-bonding pair of electrons

Question 23

are larger molecules more or less soluble in water

Answer 23

larger are less soluble

Question 24

dye vs pigment

Answer 24

dye - small, polar molecule that dissolves in water (transparent solution when dissolved)
pigment - large, nonpolar molecule that does not dissolve in water (opaque solution when mixed). A “binder” must be added to dissolve pigments.

Question 25

resonance structure definition

Answer 25

indicated by a double-headed arrow or dashed lines, electrons are delocalized and spread over multiple possible positions. used when there is more than one valid lewis formula for electron distribution

Question 26

bond order

Answer 26

single bonds - bond order 1
double bonds - 2
triple bonds - 3
in resonance structures:
if there are 2 possible positions for electrons to occupy, bond order is 1.5. For 3 possible positions, it is 1.33. etc.

Question 27

what is the geometry of each carbon atom in a benzene aromatic ring?

Answer 27

trigonal planar (120 degree bond angle)

Question 28

formal charge (bookkeeping) - how do you find it?

Answer 28

formal charge = V-(B+L)
V - number of valence electrons (the group number)
B+L - number of non-bonding electrons plus number of bonding pairs

Question 29

Deduce which of the following ions has the smallest ionic radius: Mg2+, F-, Na+, O2- (and why)

Answer 29

Mg2+
Why: all have the same electron configuration, so we must look at atomic number. Magnesium has the highest atomic number (12) so the valence electrons are mostly strongly attracted to the nucleus (so smallest radius)

Question 30

define atomic radius

Answer 30

total distance from the nucleus of an atom to the outermost orbital of its electrons

Question 31

define ionic radius

Answer 31

the distance between the nucleus of an ion and its outermost shell.

Question 32

atomic radius trend

Answer 32

increases down a group, decrease left to right across a period (highest in bottom left corner)

Question 33

define ionisation energy

Answer 33

amount of energy required to remove an electron from an atom/molecule

Question 34

ionization energy trend

Answer 34

decreases down a group, increases left to right across a period (highest in top right corner)

Question 35

electron affinity trend

Answer 35

decreases down a group, increases left to right across a period, highest in top right corner (same as ionization energy)

Question 36

define electron affinity

Answer 36

energy change that results when an electron is ADDED to a gaseous atom (opposite of ionisation energy)

Question 37

define metallic character

Answer 37

tendency of an element to lose electrons and form cations

Question 38

metallic character trend

Answer 38

increases down a group, decreases left to right across a period (highest in bottom left corner)

Question 39

define non-metallic character

Answer 39

the tendency of an element to accept electrons and form negative ions or anions

Question 40

non-metallic character trend

Answer 40

decreases down a group, increases from left to right across a period (highest in upper right corner, opposite of metallic character)

Question 41

electronegativity definition

Answer 41

Electronegativity is a measure of the tendency of an atom to attract a bonding pair of electrons

Question 42

electronegativity trend

Answer 42

decreases down a group, increases from left to right across a period (highest in top right corner)

Question 43

what d-block element is not a transition metal?

Answer 43

zinc

Question 44

describe metling points of transition metals and why

Answer 44

high melting point because of stronger metallic bond (more delocalised electrons from 4s and 3d orbitals)

Question 45

describe hardness of transition metals and why

Answer 45

because they have both 4s and 3d electrons as delocalised valence electrons, there is more attraction between electrons and the metal ions in the lattice (stronger metallic bond - greater hardness)

Question 46

describe electrical conductivity of transition metals and why

Answer 46

they are electrically conductive because of the delocalised electrons

Question 47

define metallic bond

Answer 47

electrostatic attraction between a lattice of cations and delocalised electrons

Question 48

describe how coloured complexes form

Answer 48

As they absorb radiation from the visible light range to excite electrons from one location to another, transition elements generate colourful complexes. In the presence of ligands, d- orbitals split into two sets of distinct orbital energies.

Question 49

ligand definition

Answer 49

ion or molecule that can attach to a metal through coordinate bonding (usually they donate the electron pair, but could go either way)

Question 50

2 types of stereoisomers

Answer 50

cis-trans isomers, optical isomers

Question 51

cis-trans isomerism - when does it occur

Answer 51

molecules with restricted rotation (either a c=c double bond or a ring)

Question 52

chiral carbon atom definition

Answer 52

carbon atom that is bonded to 4 different atoms/groups (said to be optically active - can have optical isomers)

Question 53

enantiomer definition

Answer 53

aka optical isomers, atoms arranged in 2 different ways that are non-superimposable (mirror images)

Question 54

polarimeter used to distinguish between pairs of which kind of isomer?

Answer 54

optical isomer

Question 55

how do optical isomers affect plane-polarised light

Answer 55

they rotate it by the same angle but in opposite directions