S3.1 Flashcards
periodic table
arrangment of elements in the periodic table
- elements are arranged according to their atomic number (Z)
- vertically arranged into groups (1-18) ⇒ elements with the same number of valence electrons, hence, chemical and physical properties
- horizontally arranged into periods (1-7) ⇒ number of principal energy levels
- elements are arranged into 4 blocks associated with the 4 sublevels (s, p, d, f)
groups of the periodic table
group 1 = alkali metals
group 2 = alkaline earth metals
groups 3-14 = transition metals
group 15 = pnictogens
group 16 = chalcogens
group 17 = halogens
group 18 = noble gases
group 1
- alkali metals
- extremely reactive (reactivity increases down the group)
- soft
- Li, Na, K have lower density than water
- low melting and boiling points (decrease down the group)
- react with water ⇒ create hydroxides ⇒ stored in oil to prevent reactions with oxygen and moisture in the air
group 17
- halogens
- reactive non-metals
- melting and boiling points increase down the group
- diatomic molecules
- reactivity decreases down the group
- a halogen will react with a solution of halide ions, formed by a less reactive halogen (more reactive halogen molecule + less reactive ion) ⇒ the halogen displaces the ion
properties of metals
good conductors of electricity and heat
malleable, ductile, lustre
act as reducing agents
most often form cations
properties of non-metals
do not conduct electricity
act as oxidising agents
most often form anions
properites of metalloids
metallic and non-metallic properties
semiconductors ⇒ their conductivity relies on temperature
“tightness of hold” on electrons - trend of the periodic trend
Across the period (to the right) ⇒ higher effective nuclear change with the same number of principal energy levels → electrons are held tighter
Down the group ⇒ growing number of principal energy levels → larger orbitals, electrons held looser, shielding effect
effective nuclear charge
the net positive charge experienced by outer valence shell electrons from the force of attraction between protons in the nucleus and the force of repulsion between inner core electrons
atomic radius
- = the distance from an atom’s nucleus to the outermost orbital of one of its electrons
- usually measured as half of the distance between the nuclei of two of the same atoms bonded together (metals - half the distance between two atoms adjacent to each other in the crystalline lattice, noble gases - half the distance between two nuclei of atoms in a frozen gas)
- the electron cloud has no definite edge
atomic radius is influenced by
- energy level ⇒ higher energy levels are further away
- charge on nucleus ⇒ more charge pulls electrons closer to nucleus
- shielding effect ⇒ core electrons in inner non-valence levels of the atom reduce positive nuclear charge experienced by valence electrons
net charge experienced by an electron
Zeff = Z - S
Z … atomic number, actual charge
S … shielding constant
trends in atomic radius across the periodic table
down a group ⇒ each atom has an additional energy level → atoms get bigger
across a period ⇒ electrons are in the same energy level, larger effective nuclear charge → outermost electrons are closer → smaller radius
ionic radius of cations and anions
Cations ⇒ smaller than the original atom (as a result of losing electrons)
Anions ⇒ larger than the original atom (as a result of gaining electrons)
less electrons => smaller
ionic radius of ions of transition elements
ions of non-transition metals have noble gas configurations (isoelectronic with them)
down a group ⇒ ionic radius of ions increases
across a period ⇒ nuclear charge increases → ionic radius decreases, even with energy level differences
ionic radius of isoelectronic ions
positive ions with the most protons are smaller
losing electrons is an endo or exothermic reaction?
endothermic
do filled and half-filled orbirals have higher or lower energy? how does this affect IE?
filled and half-filled orbitals have lower energy ⇒ lower IE
IE down a group, across a period
down a group ⇒ electron is further away from the nucleus, more shielding → decreased first IE
across the period ⇒ atoms have the same energy level and shielding, nuclear charge is increasing ⇒ increased first IE
electron affinity
def, reaction, exo/endothermic, unit
= energy change associated with the addition of a mole of gaseous electrons [kJ/mol]
X (g) + e– → X– (g) ⇒ exothermic reaction
trend in electron affinity across the period and down the group
across the period ⇒ greater nuclear charge → atoms are smaller → increased electron affinity
down the group ⇒ electrons are in higher energy levels → atoms are larger → decreased electron affinity
what is the electron affinity of nobel gases?
positive
electronegativity
= tendency of an atom to attract electrons to itself when chemically combined with another element (i.e. how strong electrons are attracted to the nucleus)
higher electronegativity ⇒ pulls electrons towards itself
atoms with higher electronegativity also have negative electron affinity
trends across a period and down a group with electronegativity
down a group ⇒ decreased electronegativity
across a period ⇒ increased electronegativity
trend in metallic character across the period, down the group?
across the period ⇒ decreases
down the group ⇒ decreases
lewis acids and baces
Lewis acid ⇒ accepts electron pair
Lewis base ⇒ donates electron pair
trend of increasing acidic character across a period
increasing acidic character across a period due to greater nuclear charge, pulling electrons closer (less likely to give them away)
amphoteric substance
able to react as both an acid and base (f.i. Al2O3)
classifications of acids + naming
- binary acids ⇒ hydro- + non-metal atom + -ic acid
- oxyacids ⇒ anion:-ate, acid: -ic; anion: -ite, acid: -ous OR -ate/-ic form with oxidation number
- organic acids
hydrofluoric acid
HF
binary acid
fluoride
F-
anion
hydrochloric acid
HCl
binary acid
chloride
Cl-
anion
hydroiodic acid
HI
binary acid
iodide
I-
anion
iodic acid
HIO3
oxyacid
iodate
IO3-
anion
hydrosulfuric acid
H2S
binary acid
sulfide
S(2-)
anion
nitrous acid
HNO2
oxyacid
nitrite
NO2-
anion
nitric acid
HNO3
oxyacid
nitrate
NO3-
anion
methanoic/formic acid
methonate, formate
HCOOH
HCOO–
ethanoic/acetic acid
ethanoate, acetate
CH3COOH
CH3COO–
sulfuric acid
H2SO4
sulfate
SO4(2-)
sulfurous acid
H2SO3
sulfite
SO3(2-)
carbonic acid
H2CO3
carbonate
CO3(2-)
phosphoric acid
H3PO4
phosphate
PO4(3-)
hypochlorous acid
HClO
hypochlorite
ClO-
chlorous acid
HClO2
chlorite
ClO2-
chloric acid
HClO3
chlorate
ClO3-
perchloric acid
HClO4
perchlorate
ClO4-
hydrocyanic acid
HCN
cyanide
CN-
oxidation state
= the number of electrons shared or transferred when forming a bond
oxidation state of hydrogen
hydrogen is combined with more electronegative elements (most non-metals) ⇒ +1
hydrogen is combined with less electronegative elements (metals) ⇒ -1
transition metals
characteristics, ions, property similarity, role in reactions
- have an incomplete d-subshell
- able to form more ions since they can lose ions from both s and d orbitals
- properties of transition metals are similar to each other, different from the properties of the main group elements due to similarities in valence electron configuration
- act as surface catalysts
surface catalyst
= reactant molecules are adsorbed onto a solid surface before they react with the catalyst to form the product
which elements of the d-block are not transition metals and why
Zn, Cd, Hg, Cn have full d orbitals ⇒ are not transition metals
properties of transition metals
- high melting points
- high density
- hard
- good electrical conductors
magnetism of transition elements
paramagnetic
ferromagnetic
diamagnetic
paramagnetism of transition elements
half-filled orbitals ⇒ weak attraction to external magnetic fields
half-filled orbitals have a net magnetic field
ferromagnetism of transition elements
only Fe, Ni, Co
have half-filled orbitals ⇒ strongly attracted to external magnetic fields and retain the direction in which these external fields act
diamagnetism of transition elements
have only filled orbitals ⇒ weakly repelled by external magnetic fields
electron configuration of transition elements
- first and second series ⇒ ns2 (n-1)dX
- third and fourth series ⇒ ns2 (n-2)f14 (n-1)dX
- f is filled before d - Ac and La are the exceptions – they have their d orbital filled before f sublevel, only after them is the f-sublevel filled
formation of ions of transition metals
lose the ns electrons first, then the (n-1)d electrons
(n - 1)d5 and (n - 1)d10 are the most stable configurations
what are the most stable electron configurations
- all orbitals are completely filled
- all orbitals are half-filled
- orbitals are empty
electron configuration of Cr
[Ar] 4s1 3d5
electron configuration of Cu
[Ar] 4s1 3d10
electron configuration of Mo
[Kr] 5s1 4d5
electron configuration of Ru
[Kr] 5s1 4d7
electron configuration of Pd
[Kr] 5s0 4d10
atomic radii of transition metals
very similar
small increase down a group
lanthanide contraction ⇒ third transition series atoms are about the same size as the second
lanthanide contraction
= decrease in expected atomic size for the transition series atoms that come after lanthanides
electrons in f orbitals have a smaller shielding ability than in other orbitals ⇒ a stronger pull on the valence electrons due to greater effective nuclear charge
ionization energy trends of transition metals
- across a period ⇒ increased first IE
- first IE of the third series is higher than those of the first and second series due to lanthanide contraction and poor shielding
electronegativity trend of transition metals
- across a period ⇒ increases, except for the last element in the series (full d-shell)
- electronegativity slightly increases between the first and second series, the third is about the same as the second
oxidation states of transition metals
most commonly exhibit multiple oxidation states that vary by 1
3B-7B ⇒ highest oxidation state = number of the group the element is in
complex ion
def, type of reaction, type of bond
monatomic cation combines with multiple monatomic anions or neutral molecules
complex ion formation is a type of Lewis acid-base reaction (metal is the base, ligands are the acid)
formed with coordinate covalent bonds
ligand
= attached monatomic anions and neutral molecules to a complex ion
coordination compound, number
complex ion + counterions = coordination compound (complex ion made neutral)
coordination number = number of atoms donating lone pairs to the central atom
primary and secondary valence
primary valence = oxidation number of the central transition metal
secondary valence = number of ligands bonded to the metal (coordination number)
coordinate covalent bond
a bond formed when the pair of electrons is donated by one atom
chelate
= complex ion containing a multidentate ligand
the multidentate ligand in a complex ion is called a chelating agent (EDTA)
types of ligands (dentates)
- unidentate ligand ⇒ has one lone pair of electrons, creating only one coordinate covalent bond with the central atom
- bidentate ⇒ two lone pairs of electrons, can create 2 coordinate covalent bonds with the metal
- multidentate ligand ⇒ 2+ lone pairs, same number of possible bonds with the metal
geometry of complex ions
2 ligands ⇒ linear
4 ligands ⇒ square planar or tetrahedral
6 ligands ⇒ octahedral
what does bonding of transition metal complexes involve?
the d orbitals of the central metal atom ⇒ a set of 5 degenerate orbitals
how does the crystal field splitting energy occur?
bonds in coordination compounds form due to the attraction of the electrons on the ligand for the + charge on the metal cation → electrons on the ligands repel electrons in the d orbitals of the metal ion ⇒ energies of the d orbitals are split
the difference in energy of the orbitals ⇒ crystal field splitting energy … Δo = h
how is the energy difference in Δo created?
after electrons on the ligands repel electrons of the d orbitals of the metal ion, the energies split
where ligands overlap with orbital lobes ⇒ strong repulsions ⇒ higher energy
ligands come in between orbital lobes ⇒ weak repulsions ⇒ lower energy
size of Δo depends on
- the identity of the metal ion ⇒ down a group of metals in the same oxidation state → increased energy gap
- the types of ligands
- oxidation state of the metal ion ⇒ charge of the metal cation increases → increased energy gap
types of ligands and effect on Δo
strong ligands ⇒ Δo is larger
weak-field ligands ⇒ Δo is smaller
electron jumps between d orbitals in transition metal complexes
electrons can jump from a lower energy level (the orbitals between the x, y and z axes) to a higher energy level (lie along the cartesian plane) if the higher energy level orbitals are either empty, half-filled or one is half-filled
transition metal ions show many intense colours in crystals or solutions due to the energy, caused by the transitions of electrons between the split d sublevel orbitals
application of coordinate compounds
- extraction from metal ores
- use of chelating agents in response to heavy metal poisoning
- qualitative analysis of metal ions
- commercial coloring agents, inks, cosmetics, paints, …
- biomolecules (porphyrin rings that entrap the metal ions - hemoglobin Fe, chlorophyll Mg, drugs - cisplatin)
transition elements with unstandard configurations
- Cr
- Cu
- Mo
- Ru
- Pd
why do melting/boiling points of alkali metals decrease down the group?
greater size and mass => weaker metallic bonds => less energy required to break them
why does the reactivity of halogens decrease down the group?
atoms become larger, more shielding => incoming electrons are not able to “add on” to the orbitals as well => smaller tendency to gain electrons, react
why does reactivity of alkali metals increase down a group?
atoms become larger, more shielding => lower IE => higher tendency to give electrons, react
why do the melting/boiling points of halogens increase down the group?
greater mass => stronger intermolecular (dispersion, London) forces => more energy to break them
