1
Q
atomic radius
A
the distance from the centre of the nucleus to the outermost shell of electrons
2
Q
bonding atomic radius
A
- aka covalent radius
- half of the distance between the centre of the nuclei of 2 covalently-bonded atoms
3
Q
why is atomic radius (RnB) > covalent radius (Rb)?
A
- the covalent bond is formed by the overlapping of atomic orbitals
- the overlapping region becomes common ground
4
Q
periodicity
A
the tendency of similar properties of the element to recur at intervals, if they are arranged in order of increasing atomic number
5
Q
periodic trend of atomic radius
A
- across a period, atomic radius decreases due to increased nuclear charge, which pulls electrons closer to the nucleus
- down a group, atomic radius increases due to increased no. of shells and increased shielding effect
6
Q
ionic radius
A
the effective distance from the nucleus of the ion up to which it has an influence in the ionic bond
7
Q
characteristics of ionic radius
A
- radius of cation radius of parent atom
8
Q
electron affinity
A
- the amount of energy released when a mole of electrons is added to a mole of gaseous atoms
- represents the ability of an atom to hold an additional electron
- the higher the tendency, the higher the e.a.
9
Q
factors affecting electron affinity
A
- nuclear charge: higher nuclear charge = higher e.a.
- size of atom: the larger the atom, the lower the e.a.
- e. config.: stabler configs = lower e.a.
10
Q
periodic trends of electron affinity
A
- along a period, e.a. increases, because nuclear charge increases and atomic size decreases
- down a group, e.a. decreases, because although size and nuclear charge both increase, effect of increased size is much more pronounced
11
Q
electronegativity
A
a measure of the ability of atoms to attract a shared pair of electrons toward itself in a covalent bond
12
Q
polar covalent bond
A
bond in which electrons are shared unequally, resting in a partial polar charge
13
Q
periodic trends of electronegativity
A
- along a period, electronegativity increases due to increased nuclear charge, resulting in increased attraction between nucleus and electrons
- down a group, electronegativity decreases due to increase in atomic size, resulting in reduced attraction
14
Q
types of oxides
A
- acidic oxide
- basic oxide
- neutral oxide
- amphoteric oxide
15
Q
Group I trends
A
Down the group:
- reactivity increases
- m.pt/b.pt increases
- density decreases
16
Q
Group VII trends
A
down the group:
- m.pt/b.pt increases
- density increases
- reactivity decreases
17
Q
transition element
A
- element whose atom has an incomplete d sub-shell
- or can give rise to a cation with an incomplete d sub-shell
18
Q
properties of transition metals
A
PHYSICAL PROPERTIES
- high m.pt/b.pt
- high tensile strength
- high density
- malleable and ductile
- high thermal & electrical conductivity
CHEMICAL PROPERTIES
- form compounds with more than one oxidation no
- incomplete d-subshell
- form complex ions
- form colored compounds
- act as catalysts when either elements or compounds
19
Q
why does chromium show a large jump in IE for the 7th e-?
A
Chromium’s config: [Ar] 3d5
- because there’s not much diff between 3d and 4s orbitals
- but taking from 3p6 is harder (also because the orbital is full)
20
Q
first ionisation energy
A
- the amount of energy required
- to remove 1 mol of e-s from 1 mol of atoms in gaseous state
21
Q
what does group number tell you about electron placement?
A
number of outer shell/valence electrons
22
Q
what does period number tell you about electron placement?
A
number of occupied electron shells
23
Q
Why do halogen m.pts/b.pts increase down the group?
A
- coz their Mr also increases down the group
- thus van der waals/intermolecular forces also increase
24
Q
why do alkali metal m.pts/b.pts decrease down the group?
A
- metallic bonding weakens
- as radius increases, delocalized electrons are shielded more from the effective nuclear charge through electron shielding effect
25
why do halogen b.pts increase as Mr increases?
- coz their intermolecular attraction increases
| - due to temporarily induced dipoles
26
explain why there's an increase in m.pt from one side of a period to the other (cations only)
- cations become more positive
| - size and radius decreases, so charge density increases
27
explain why there's an increase in m.pt from one side of a period to the other (anions only)
- Mr increases
| - van der waals/london forces also increase
28
why does Si have a greater m.pt than Ar?
- Si has a giant 3-D macromolecular covalent bonding structure
- while Ar is bonded by atomic van der waals/london/dispersion forces
29
is Zn a transition metal?
- it's a d-block element but not a transition metal
- zinc has a complete d subshell in both neutral atom and ion form
- it also only has 1 oxidation state
30
is Sc a transition metal?
- it has no d electrons
- so it's colorless in aqueous solutions
- but its atom has an incomplete d subshell
- and it has 2 oxidation states
- and the Sc2+ ion has a d e-
- Sc is a transition metal
31
why do transition metals have variable oxidation numbers?
- due to patterns in successive ionization energies
- outer d and s subshells in transition metals are closer in energy compared to other elements
- so they don't experience the characteristic jump in IE from an s e- to a d e-
32
characteristic oxidation states shared by transition metals
- oxidation states correspond to the use of 3d and 4s e-s in bonding
- all transition metals can form ions of +2 and +3 oxidation states
- M3+ is the most stable oxidation state for Sc to Cr, but all other transition metals prefer M2+ due to increased nuclear charge making IE3 comparatively higher
33
highest possible oxidation state of a transition metal
+7 at Mn
34
can transition metals show covalent characteristics?
- yes
- at oxidation states > +3
- this is due to ions with higher charge having such a large charge density that they polarize cations and increase the covalent nature of the compounds
35
uses of compounds with high oxidation states
oxidizing agent
| e.g. K2Cr2O7
36
complex ion
- transition metal ions in aqueous solutions that form coordinate bonds with water molecules
- due to the metal ions' high charge density
37
ligands
- the species surrounding a central ion in a complex ion
| - all ligands have at least 1 atom with a lone pair of e-s
38
how is a complex ion formed?
- forms when a central ion is surrounded by molecules/ions (ligands) with a lone pair of e-s
- the ligands attach via coordinate bonding
39
coordination number
the number of coordinate bonds formed by the central ion
40
polydentate ligand
- ligands that form more than 1 coordinate bond
| e. g. as a hexadentate ligand, EDTA4- can form 6 coordinate bonds, and thus is equal to 6 monodentate ligands
41
chelate
compound containing a ligand that has occupied more than one bonding site (polydentate ligand)
42
importance of chelates
- can remove transition metal ions from solutions
| - thus inhibiting enzyme catalyzed oxidation reactions
43
heterogenous catalysts
the catalyst is in a different state from the reactants
44
why do transition metals make effective heterogenous catalysts?
- they can use their 3d and 4s e-s to form weak bonds to reactants
- thus enabling them to collide with the correct orientation
45
examples of transition metals as heterogeneous catalysts
- Fe in Haber's Process
- Ni in alkenes --> alkanes
- Pd and Pt in catalytic converters
- MnO2 in the decomposition of hydrogen peroxide
- V2O5 in Contact Process
46
homogeneous catalysts
catalyst in same state as reactants
47
why do transition metals make effective homogeneous catalysts?
- transition metal ions have variable oxidation states
| - thus they are particularly effective catalysts in redox reactions
48
examples of transition metals as homogeneous catalysts
- Fe2+ in heme (haemoglobin)
| - Co3+ in Vit B: 5 of the 6 sites are occupied by N, and the 6th site is used for biological activity
49
diamagnetism
- common to all materials
- as the orbital motion of electrons is what produces the diamagnetic effect
- it produces weak opposition to applied magnetic field
50
paramagnetism
- common to all substances with unpaired e-s (includes all transition metals and their compounds)
- stronger than diamagnetism
- produces magnetization proportional to applied magnetic field and in same direction
51
ferromagnetism
- occurs in substances with long range ordering of unpaired e-s
- strongest form of magnetism
- produces magnetization greater than applied magnetic field
e.g. iron, nickel, cobalt
52
domains
- occurs in ferromagnetism
- domains are regions in which unpaired d e-s in a large no of atoms line up with parallel spins
- they are randomly oriented but can be more ordered when a magnetic field is applied
- the magnetism will remain even after the magnetic field is removed
53
why do transition metals appear colored?
- transition metal ions absorb specific colors from the color spectrum
- in the presence of a ligand's lone pair of e-s, their d orbitals split into 2 sub-levels
- upon absorbing light, a 3d e- becomes excited and passes into the higher energy level
54
factors affecting color of complex
- nuclear charge and identity of central ion
- charge density of ligand
- geometry of complex ion
- no. of d e-s present (i.e. oxidation no of central ion)
55
factors affecting color of complex: nuclear charge and identity of central ion
- the strength of the coordinate bond depends on the electrostatic attraction between the nuclear charge and lone pair of e-s
- higher nuclear charge = higher electrostatic attraction = higher energy wavelengths (i.e. shorter wavelengths) absorbed
56
factors affecting color of complex: charge density of ligand
- ligands with higher charge densities cause greater splitting of d orbitals
- greater splitting = greater energy of light (i.e. shorter wavelengths) absorbed
57
factors affecting color of complex: geometry of complex ion
the splitting in energy of d orbitals depends on the relative orientations of the ligands and d orbitals
58
factors affecting color of complex: no of d orbitals and oxidation state of central metal ion
the no of d orbitals and the oxidation state affect:
- strength of interaction
- amount of e- repulsion
between the ligands and the central ion
Chemistry IB HL
flashcards
Decks in class (11)
# Cards
-
Topic 1: Stoichiometric Relationships35
-
Topic 2: Atomic Structure67
-
Topic 3: Periodicity58
-
Topic 4: Chemical Bonding and Structure65
-
Topic 5: Energetics & Thermochemistry66
-
Topic 6: Chemical Kinetics28
-
Topic 7: Equilibrium26
-
Topic 8 - Acids & Bases81
-
Topic 9: Redox Processes34
-
Topic 10 - Organic Chemistry158
-
Topic 11 - Measurement & Data Processing42
