Semester 1 (spring 2024) final exam Flashcards
covers most of S1, S2, and S3, and R2.1, R3
full and condensed electron configuration for silver
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d9, OR [Kr] 5s2 4d9
aufbau principle exceptions
chromium and copper
chromium condesned electron configuration
[Ar] 3d5 4s1
copper condensed electron configuration
[Ar] 4s1 3d10
convergence limit for hydrogen is 3.28 × 10^15 s^−1. what is the first ionisation energy in kJ/mol?
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.
what can you determine from a graph showing the successive ionisation energies of an element? (3)
- how many energy levels it occupies: this is equal to the number of jumps/large increases on the graph
- 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.
- 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.
electromagnetic spectrum: list from highest to lowest frequency (7 wave types)
gamma rays, x rays, ultraviolet, visible, infrared, microwave, radio
what is the max number of orbitals in each energy level (its a formula)
n^2
how to tell if a bond is polar or nonpolar
if the difference in electronegativity is greater than 0.4, it is a polar bond
sigma bond definition
strongest type of covalent bond. formed by head-on overlap between orbitals
pi bond definition
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
order of intermolecular forces (3) from weakest to strongest
LDF<dipole-diple<hydrogen bonding
larger ionic radius: weaker or stronger bond?
weaker
greater charge: stronger or weaker bond?
stronger
coordination bond definition
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)
inter vs intra molecular
inter - solid, liquid, gas (ex. melting overcomes intermolecular forces - forces BETWEEN molecules
intra - bonds that hold the atoms together WITHIN the molecule
london dispersion forces
the random and continuous movement of electrons creates asymmetrical electron distribution, which can attract nearby atoms/molecules (always present)
dipole-dipole
Only exists for two asymmetrical molecules (those with permanent dipoles). Much stronger than LDF.
what creates a dipole
If polar bonds aren’t arranged symmetrically, the molecule will have a permanent dipole
Dipole-induced-dipole
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)
what affects strength of dipole-induced-dipole force
large molecules are more susceptible than smaller ones
requirements for hydrogen bonding (2)
- hydrogen bonded to a FON element (fluorine, oxygen, nitrogen)
- The FON element must have a non-bonding pair of electrons
are larger molecules more or less soluble in water
larger are less soluble
dye vs pigment
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.
resonance structure definition
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
bond order
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.
what is the geometry of each carbon atom in a benzene aromatic ring?
trigonal planar (120 degree bond angle)
formal charge (bookkeeping) - how do you find it?
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
Deduce which of the following ions has the smallest ionic radius: Mg2+, F-, Na+, O2- (and why)
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)
define atomic radius
total distance from the nucleus of an atom to the outermost orbital of its electrons
define ionic radius
the distance between the nucleus of an ion and its outermost shell.
atomic radius trend
increases down a group, decrease left to right across a period (highest in bottom left corner)
define ionisation energy
amount of energy required to remove an electron from an atom/molecule
ionization energy trend
decreases down a group, increases left to right across a period (highest in top right corner)
electron affinity trend
decreases down a group, increases left to right across a period, highest in top right corner (same as ionization energy)
define electron affinity
energy change that results when an electron is ADDED to a gaseous atom (opposite of ionisation energy)
define metallic character
tendency of an element to lose electrons and form cations
metallic character trend
increases down a group, decreases left to right across a period (highest in bottom left corner)
define non-metallic character
the tendency of an element to accept electrons and form negative ions or anions
non-metallic character trend
decreases down a group, increases from left to right across a period (highest in upper right corner, opposite of metallic character)
electronegativity definition
Electronegativity is a measure of the tendency of an atom to attract a bonding pair of electrons
electronegativity trend
decreases down a group, increases from left to right across a period (highest in top right corner)
what d-block element is not a transition metal?
zinc
describe metling points of transition metals and why
high melting point because of stronger metallic bond (more delocalised electrons from 4s and 3d orbitals)
describe hardness of transition metals and why
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)
describe electrical conductivity of transition metals and why
they are electrically conductive because of the delocalised electrons
define metallic bond
electrostatic attraction between a lattice of cations and delocalised electrons
describe how coloured complexes form
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.
ligand definition
ion or molecule that can attach to a metal through coordinate bonding (usually they donate the electron pair, but could go either way)
2 types of stereoisomers
cis-trans isomers, optical isomers
cis-trans isomerism - when does it occur
molecules with restricted rotation (either a c=c double bond or a ring)
chiral carbon atom definition
carbon atom that is bonded to 4 different atoms/groups (said to be optically active - can have optical isomers)
enantiomer definition
aka optical isomers, atoms arranged in 2 different ways that are non-superimposable (mirror images)
polarimeter used to distinguish between pairs of which kind of isomer?
optical isomer
how do optical isomers affect plane-polarised light
they rotate it by the same angle but in opposite directions
