S2 Flashcards
1
Q
chemical bond def
+what effect they have on electron configurations of atoms involved
A
= strong forces of attraction that hold atoms or ions together in a substance
create more stable electron configuration for the atoms involved
2
Q
ionic bond def
A
electrostatic attraction between electric charges of a cation and an anion
3
Q
anion
def, created by, naming
A
negatively charged ion (more electrons than protons)
created by reduction (accepting electrons)
naming: parent atom + “-ide”
4
Q
cation
def, creation, naming
A
positively charged ion (more protons than electrons)
created by oxidation (giving away electrons)
naming: same as parent atom
5
Q
polyatomic ions def
A
ions composed of multiple atoms held together by covalent bonds
6
Q
name NH4+
A
ammonium
7
Q
name OH¯
A
hydroxide
8
Q
name HCO3¯
A
hydrogencarbonate
9
Q
the octet rule
A
= the tendency of an atom to have a valence shell with a total of 8 electrons
10
Q
strength of the ionic bond depends on:
A
- ionic charge: larger ionic charge ⇒ stronger electrostatic attraction
- ionic radius: smaller ionic radius ⇒ stronger electrostatic attraction
11
Q
naming ionic compounds
A
cation name + anion name (does not reflect ratio of ions)
12
Q
structure of ionic compounds
A
lattice, composed of 3D repeating units of ions
no separate ionic compound species
13
Q
coordination number X/Y
A
- X … how much of the anion is around one cation
- Y … how much of the cation is around one anion
- ratio of elements in the empirical formula is proportional to the coordination number
14
Q
lattice enthalpy
def, equation, value, larger H effect on ionic bond
A
- ΔH°lattice
- = the standard enthalpy change that occurs on the formation of gaseous ions from one mole of the solid lattice
- greater than 0 (since breaking bonds is endothermic)
- MX(s) → M+(g) + X¯(g)
- larger lattice enthalpy ⇒ stronger ionic bond
15
Q
characteristics of ionic compounds
SATP, mp/bp, conductivity, volatility, hardness, malleability,solubility
A
- solid under SATP
- ionic bonds are non-directional
- high melting and boiling points (takes more E to break the lattice)
- low volatility
- conduct electricity in molten and aqueous states, low/don’t in solid/gas due to freely moving electrons
- bad thermal conductors due to no free e-
- hard ⇒ form crystals
- brittle
- most are soluble in water (those with extremely high lattice enthalpies are not)
16
Q
explain the brittleness of ionic compounds
A
a stress is applied ⇒ layers within the lattice shift in response to that stress, ions become adjacent to like charges, which experience an electrostatic repulsion, causing layers to separate from one another, which forces the lattice to break into pieces
17
Q
what causes a greater ionic character of a bonds
A
greater difference in electronegativity between two elements in a compound ⇒ greater ionic character of the bond
18
Q
difference in electronegativity ⇒ type of bonding:
A
0 = pure covalent
0.1 – 0.4 = non-polar or weakly covalent
0.5 – 1.7 = polar covalent
> 1.8 = ionic
19
Q
ionic liquids
A
- solvents, electrolytes
- usually large, irregularly shaped ⇒ static hindrance ⇒ lower boiling/melting points
- usually contain organic compounds
- stable, like all ionic compounds
- eutectic solvents
20
Q
static hinderance
A
= the slowing of chemical reactions due to amount of space that a group of atoms takes
21
Q
covalent bond def
A
= electrostatic attraction between a pair of electrons and positively charged nuclei
22
Q
organic vs inorganic compound
A
Organic compound ⇒ carbon is bonded to hydrogen
Inorganic compound ⇒ carbon may be present, but it is not bonded to hydrogen
23
Q
diatomic molecules
A
halogens, H, O, N
24
Q
formation of covalent bonds
A
two atoms approach each other close enough for their electron clouds to overlap ⇒ repulsion between the electrons of both atoms, repulsion between the nuclei, attraction between the electron cloud of one atom and the nuclei of another atom
25
what happens when optimum distance is achieved in a covalent bond
=> forces are balanced, causing a release of energy
26
coordination (coordinate, dative) bond
both the electrons in the covalent bond come from the same atom
27
oxidation state
= hypothetical charge an atom would have if all the bonding electrons were assigned to the more electronegative element
28
formal charge
= hypothetical charge an atom would have if the bonding electrons were shared equally between atoms
= number of valence electrons – number of electrons assigned to the atom
29
which lewis structures are the most optimal based on formal charges
* **neutral molecules** ⇒ Lewis structure in which there are no formal charges is preferable
* Lewis structures with large formal charges are less plausible than those with **small formal charges**
* Lewis structures with similar distributions of formal charges ⇒ the most plausible structure is the one in which **negative formal charges are placed on the more electronegative atoms**
30
exceptions to the octet rule
* H ⇒ only 2 electrons in its outer shell
* odd electron species (NO)
* incomplete octet (Be, B, Al)
* expanded valence shell (central atom with principal quantum number n > 2)
31
elements with an incomplete octet
Be ⇒ 2 el. pairs
B ⇒ 3 el. pairs
Al ⇒ 3 el. pairs
32
bond length
= the distance (in pm) between the nuclei of two atoms joined by a covalent bond
inversely proportional to bond enthalpy
atoms vibrate ⇒ bong length is not a fixed value, average length can be measured however
33
bond order and bond length
triple bond < double bond < single bond
34
bond enthalpy
| +relationship to bond length
= enthalpy change required to break a particular bond in one mole of gaseous molecules
inversely proportional to bond length
35
bond order and bond enthalpy
triple bond > double bond > single bond
36
is it easier to break a bond between 2 larger atoms or smaller atoms?
| in connection to bond length
between two larger atoms, since bond length increases with the size of the atoms => smaller bond enthalpy => less E needed to break a bond
37
bond polarity
difference in electronegativity of bonded atoms
higher electronegativity ⇒ pulls electrons to itself ⇒ greater electron density at the atom with the higher electronegativity (δ¯)
38
electron density, δ¯, δ+
= the relative amount of electrons in a region of space
δ¯ ⇒ negative partial charge (higher electron density)
δ+ ⇒ positive partial charge (smaller electron density)
39
dipole moment
= **separation of charge** between two non-identical bonded atoms
vector μ, represented with with an arrow
40
non-polar covalent vs polar covalent bond
**non-polar covalent** ⇒ electrons are shared equally between two atoms, no partial charges
**polar covalent** ⇒ electrons are not shared equally between two atoms, partial charges
41
molecular polarity
net dipole moment, a measure of overall polarity of a molecule
symmetrical molecule ⇒ no net dipole moment
asymmetrical molecule ⇒ bond dipoles do not cancel out
42
delocalization
⇒ electrons are shared between more than two nuclei in a molecule
give greater stability to a molecule/polyatomic ion
43
resonance structures
in molecules in which there is more than one possible Lewis structure
44
bond order of an identified bond in a resonance structure
= sum of bond in all resonance structures ÷ number of resonance structures
45
formal charge of an identified atom in a resonance structure
= sum of formal charges in all resonance structures ÷ number of resonance structures
46
aromatic molecule
= cyclic planar molecules that are stabilized due to the presence of delocalized electrons
47
VSEPR theory
* = Valence Share Electron Pair Repulsion theory
* electron pairs repel each other ⇒ arrange themselves as far apart from each other as possible
* **non-bonding electron pairs occupy more space** than bonding pairs
* **double and triple bonds occupy more space** than single bonds
* electron pairs assume **orientation around an atom to minimise repulsions** ⇒ particular geometric shape of molecules
48
which repulsions are the greatest?
| lp-lp, lp-bp, bp-bp
lone pair-lone pair repulsion > lone pair-bonding pair repulsion > bonding pair-bonding pair repulsion
49
2 electron domains, 2 bonding, 0 non-bonding
| bond angle, geometry
linear, 180°
50
3 electron domains, 3 bonding, 0 non-bonding
| geometry, bond angle
trigonal planar, 120°
51
3 electron domains, 2 bonding, 1 non-bonding
| geometry, bond angle
bent, < 120°
52
4 el domains, 4 bonding
| geometry, bond angle
tetrahedral, 109.5°
53
4 el domains, 3 bonding, 1 non-bonding
| geometry, bond angle
trigonal pyramidal, < 109.5°
54
4 el domains, 2 bonding, 2 non-bonding
| geometry, bond angle
bent (v-shaped), < 109.5°
55
5 el domains, 5 bonding, 0 non-bonding
| geometry, angle
trigonal bypiramidal, 90° and 120°
56
5 el domains, 4 bonding, 1 non-bonding
| geometry, bond angle
seesaw, < 90° and < 120°
57
5 el domains, 3 bonding, 2 non-bonding
| geometry, bond angle
T-shaped, < 90°
58
5 el domain, 2 bonding, 3 non-bonding
linear, 180°
59
6 el domains, 6 bonding
octahedral, 90°
60
6 el domains, 5 bonding, 1 non-bonding
| geometry, angle
square pyramidal, < 90°
61
6 el domains, 4 bonding, 2 non-bonding
| geometry, angle
square planar, 90°
62
molecular orbitals
result from the combination of atomic orbitals
wave functions
63
combination of atomic orbitals
1. **constructively** (combination of non-opposite waves) ⇒ creates bonding molecular orbitals (larger), result of constructive interference
1. **destructively** (combination of opposite waves) ⇒ anti-bonding molecular orbitals (cancels out), result of destructive interference
64
what does the strength of a bond depend on?
| molecular orbitals
degree of atomic orbital overlap
65
origin of σ-bonds
end-on-end overlap along the bond axis
combination of 2 s orbitals, 1 s and 1 p orbital, 2 p orbitals
66
π-bonds
| origin, where
* origin: side-by-side overlap around the internuclear axis, can only form between p-orbitals
* double and triple bonds (around the sigma bond)
* their electron density is concentrated above and below the bond axis
* electrons in a π-bond are delocalized
67
which can rotate? σ or π-bonds?
only σ-bonds
68
which is stronger? σ or π-bonds? why?
**sigma bonds**: in only one electron-dense region on the internuclear axis, located close to the nuclei ⇒ stronger attraction between the electron pair and positive nuclei
**pi bonds**: electron density is spread across two regions: above and below the internuclear axis, found further from the positive charge of the nucleus
⇒ σ-bonds are stronger than π-bonds
69
which is easier to break? triple/double bonds or single bonds?
| sigma/pi bonds
π-bonds are weaker than σ-bonds ⇒ π-bonds are easier to break ⇒ easier to break triple and double bonds than single bonds
70
hybridization
= concept of mixing atomic orbitals to form new hybrid orbitals for bonding by the promotion of electrons from s and p orbitals (excitation)
71
excitation
= the process of changing from a ground state to an excited state by energy absorption
72
general rules on hybrid orbitals
| strength, energy, planes, which can form
* greater overlap ⇒ stronger bonds
* total energy of orbitals does not change, it is only redistributed
* all hybridised orbitals exist in only one plane
* only σ-bonds form hybridized orbitals, π-bonds cannot form unhybridized orbitals
73
2 el domains
| hybridization, geometry
sp, linear
74
3 el domains
| hybridization, geometry
sp2, trigonal planar
75
4 el domains
| hybridization, geometry
sp3, tetrahedral
76
5 el domains
| hybridization, geometry
sp3d, trigonal bipyramidal
77
6 el domains
| hybridization, geometry
sp3d2, octahedral
78
allotrope
= different structural forms of the same element
79
carbon allotropes
* diamond
* graphite
* graphene
* fullerene (C60)
80
silicon allotropes
* pure silicon
* silica (SiO2)
81
diamond
| structure, m/bp, conductivity, colour, hardness, uses
* carbon allotrope
* each carbon atom is covalently bonded to four others (tetrahedral, 109.5°)
* can theoretically go on forever
* high melting and boiling points
* no free electrons ⇒ bad electrical conductivity
* forms colourless and transparent crystals
* hard
* used in jewellery and ornaments, in drill bits and diamond saws
82
graphite
| structure, appearance, conductivity, uses
* carbon allotrope
* each carbon atom is covalently bonded to three others (trigonal planar, 120°)
* layers arranged in fused hexagonal rings
* layers held together by weak London dispersion forces ⇒ easy to separate ⇒ soft and slippery
* hexagonal rings contain delocalized electrons ⇒ move around freely between layers
* dark grey shiny solid
* good electrical conductivity due to delocalized electrons
* used in pencils, as lubricant, as electrodes and in electric motors
83
graphene
| structure, strength, conductivity, uses
* carbon allotrope
* single graphite layer
* strongest thin material
* synthetic allotrope
* high thermal and electric conductor
* high tensile strength
* sheets can be rolled into tubes (nanotubes) ⇒ use in the medical and pharmaceutical industry, electronics
84
fullerene
| structure, conductivity
* carbon allotrope
* C60: 60 carbon atoms bonded together in 20 hexagons and 12 pentagons (truncated icosahedron)
* each carbon atom is covalently bonded to 3 other carbon atoms ⇒ delocalized electrons
* poor electrical conductivity
85
why is fullerene a bad electrical conductor?
delocalized electrons can only move around each fullerene and not jump between two separate fullerenes ⇒ limited conductivity, less than in graphene and graphite, where delocalized electrons can move in along a linear path
86
pure silicon
* 3-D lattice (each silicon atom is covalently bonded to 4 others in a tetrahedral arrangement)
* no delocalized electrons ⇒ poor electrical conductivity
* strong covalent bonds ⇒ high melting and boiling point
87
silica
| structure, mp, bp, conductivity, uses
* lattice structure
* tetrahedral arrangement
* each silicon is bonded to 4 oxygen atoms, each oxygen to 2 silicon atoms
* Si — O — Si geometry ⇒ bent (v-shaped)
* strong covalent bonds between silicons and oxygens ⇒ high melting and boiling point
* does not conduct electricity in solid state
* quartz, glass
88
covalent vs molecular crystal/network
**covalent** => basic units are atoms, held together by covalent bonds (stronger than intermolecular forces)
**molecular** => basic units are molecules, held together by weak intermolecular forces
89
characteristics of covalent networks
| volatility, bp, mp, SATP, conductivity, solubility
* strong bonding ⇒ non-volatile, solid at SATP, high boiling and melting point
* do not conduct electricity; exceptions: graphite, graphene, silicone (semiconductor)
* bad thermal conductors due to lack of free e- (except for diamond, graphite, graphene)
* strong bonding ⇒ difficult to dissolve
90
molecular network characteristics
| volatility, bp, mp, conductivity, solubility
* weak intermolecular forces ⇒ **volatile**, **low boiling and melting** point
* do not conduct electricity
* bad thermal conductors (weak intermolecular forces, less efficient at transporting vibrations)
* solubility: **“like dissolves like”** ⇒ solutes dissolve best in solvents of similar polarity due to similar bonding
91
when are intermolecular forces present
* always present in solid and liquid states
* at their max. in solid state
* ideal gas ⇒ no intermolecular forces
92
intramolecular vs intermolecular forces
intramolecular forces ⇒ bonds between atoms (covalent, ionic, metallic)
intermolecular forces ⇒ between molecules of a substance
93
kinds of intermolecular forces
1. london dispersion forces
2. dipole-induced forces
3. dipole-dipole forces
4. hydrogen bonds
94
London dispersion forces
⇒ at any moment in time, electron pairs are not distributed evenly ⇒ temporary induced dipoles
between all molecules
95
polarizability
= the ease of distortion of the electron cloud by an electric field
higher polarizability ⇒ stronger dispersion forces
96
factors that affect the strength of London dispersion forces
1. **number of electrons**: more electrons ⇒ greater distance from nucleus, larger electron cloud ⇒ larger polarizability ⇒ stronger
1. **shape of the molecule**: larger molecule, larger surface area ⇒ larger polarizability
97
dipole-induced forces
⇒ polar molecule induces the shift of electrons in a nonpolar molecule to form a temporary dipole, due to the attraction between negative and positive charges
98
dipole-dipole forces
| def, when they are stronger
⇒ electrostatic attraction between the negative pole of one molecule and the positive pole of another ⇒ orientation of positive-negative pole
larger net dipole moment ⇒ stronger dipole-dipole forces
99
hydrogen bond
⇒ electrostatic attraction between a hydrogen atom, which is covalently bound to a highly electronegative element (N, O, F), and another electronegative atom bearing a lone pair of electrons
180° angle
100
why is ice less dense than water
hydrogen bonding in water molecules ⇒ structure of ice … large gaps ⇒ smaller density than in water
101
chromotography def
= separation technique of a mixture based on the affinities of its components for the mobile or stationary phase, which are based on the intermolecular forces of the components of the mixture, the stationary phase and the mobile phase
102
what happens in chromotography when the substance has strong intermolecular forces?
substance has strong intermolecular forces ⇒ remains adsorbed for much longer than if it has weak intermolecular forces
103
absorption vs adsorption
**adsorption** ⇒ atoms, ions, compounds cling to the surface of the molecule
**absorption** ⇒ atoms, ions, compounds, substances enter a liquid, solid
104
kinds of chromotography
1. paper chromotography
2. TLC
3. liquid column chromotography
4. gas liquid chromotography
5. HPLC
105
mobile and stationary phase in paper chromotography
mobile phase: liquid solvent
stationary phase: paper (hydrated cellulose = glucose)
106
retardation (retention) factor
RF = b/a
a … distance travelled by the solvent (to the solvent front)
b … distance travelled by a component of the mixture
107
mobile and stationary phase in TLC
stationary phase: plate made of glass/metal coated with silica/alumina (polar as they contain many hydroxyl groups on the surface)
mobile phase: organic non-polar solvent
| thin layer chromotography
108
liquid column chromotography
mobile phase: solvent
stationary phase: adsorbent column of SiO2
sample is placed on top of a column of adsorbent SiO2 ⇒ solvent is poured down the column ⇒ different components separate into bands
wait until bands reach the bottom ⇒ collect them for further analysis
109
gas liquid chromotography
* mobile phase: gas
* stationary phase: dense and viscous liquid
* stationary phase covers the inside of a tube, through which a gas is pushed through ⇒ components of the sample (in the gas) separate within the tube in a heated chamber
* separation based on the difference in volatility
110
HPLC
* = high performance liquid chromotography
* mobile phase: liquid
* stationary phase: usually solid
* pump pushes mobile phase through the column, which contains the stationary phase (usually solid) and liquid sample
* faster than paper chromatography
111
dye vs pigment
**dye** ⇒ small, polar molecules (dissolves in polar solvents (like water) ⇒ forming transparent solutions)
**pigment** ⇒ large, nonpolar molecules (does not dissolve in polar solvents, remains suspended ⇒ forming opaque solutions)
112
how are electronegativity and metallic character connected?
smaller electronegativity ⇒ higher metallic character, smaller covalent character
113
metallic bond def
= electrostatic attraction between a lattice of positive ions and delocalized electrons
non-directional
114
strength of a metallic bond depends on:
1. **number of valence electrons** that become delocalized ⇒ more valence electrons → more electrons in the sea → stronger bond
1. **charge of the metal ion** ⇒ larger charge → more electrons in the sea and greater charge difference → greater electrostatic attraction → stronger bond
1. **ionic radius of the metal cation** ⇒ smaller radius → shorter distance between the positive nucleus of the cation and the surrounding delocalised electrons → larger attraction between nucleus and delocalized electrons → stronger metallic bond
115
metal properties
* conductors of electricity and heat
* ductile and malleable
* shiny
* sonorous
116
electrical conductivity def
= ability of charged particles to move through a region of space
117
thermal conductivity def
= a material’s ability to transfer heat as a result of a temperature difference
118
explain metallic electrical conductivity
potential energy difference is applied ⇒ delocalised electrons are repelled by the negative terminal and attracted to the positive terminal ⇒ electrical conductivity
119
explain metallic thermal conductivity
a region of the metallic lattice encounters an increase in temperature ⇒ the kinetic energy of the delocalised electrons in this region also increases ⇒ these delocalized electrons travel towards a region of the lattice which is at a lower temperature ⇒ transferring heat across the material
120
explain the general conductivity of metals
vibrations pass through easier with delocalized electrons ⇒ better conduction
121
ductility vs malleability
ductility ⇒ can be drawn out into a wire
malleability ⇒ can be reshaped on compression
122
explain metallic ductility and malleability
stress is applied ⇒ layers within the lattice shift in response ⇒ the cations in the lattice remain surrounded by delocalised valence electrons, meaning the metallic bonding also remains unaffected
123
why are metals shiny?
visible light hits delocalized electrons ⇒ electrons absorb light and vibrate ⇒ vibrations generate a second wave of light, which radiates from the surface => reflection of light appears as shine
124
explain why metals are sonorous
metallic lattice can move easily ⇒ incoming sound is spread easily throughout the material, since free electrons can move easily
125
which metals are more sonorous? less or more dense ones?
low density ⇒ more space between cations for the delocalised electrons to move ⇒ more sonorous
126
resistance
* ⇒ lowers conductivity, usually in the case of metalloids
* more vibrations (therefore energy) ⇒ interference with the clear path of electrons in the lattice = resistance
* 0 K (absolute zero) ⇒ no vibrations ⇒ resistance = 0
127
superconductors
resistance 0 at T ≠ 0 K, i.e. T > 0 K
128
which metals have higher conductivity?
more valence electrons ⇒ higher conductivity, except for group 3 (closer to metalloids, more collisions between electrons → loss of energy)
129
metallic bonds of transition metals
large number of valence electrons from both the s and d orbitals ⇒ greater electron density within the metallic lattice ⇒ increased strength of the metallic bond, higher conductivity
mostly d-orbital electrons are delocalized
130
division of materials
1. metals
1. polymers
1. ceramics (not conductive)
131
alloy def
= homogeneous mixture composed of 2+ metals or a metal + non-metal
132
properties of alloys
* greater strength
* greater resistance to corrosion
* enhanced magnetic properties
* less malleable, less hard
* lower melting temperatures
133
explain the decreased malleability and hardness of alloys
addition of another element disrupts the perfect arrangement of metal cations ⇒ not as easy to disrupt the lattice ⇒ less malleable, harder
134
explain the decreased melting/boiling points of alloys
differing cation sizes ⇒ weaker electrostatic attraction between cations and delocalized electrons (weaker metallic bond)
135
substitutional vs interstitial alloys
**substitutional alloys** = the element added to the base metal simply replaces the metal ions in the lattice
**interstitial alloys** = the element added occupies vacant space in the metallic lattice of the base metal
136
common alloys
* steel = iron + carbon
* bronze = copper + tin
* brass = copper + zinc
* commercial gold (XK) = gold + copper + silver
* sterling silver = silver + copper
137
what is rust?
iron oxide, a result of oxidation of iron
138
rust protection methods
barrier methods: painting, oiling
sacrificial methods (f.i. galvanization)
139
what is meant by sacrificial methods of rust protection
highly active metals that are used to prevent a less active material surface from corroding (in this example, iron from oxidizing)
140
galvanization
sacrificial method of rust protection by applying a protective zinc coating to steel or iron, to prevent rusting
141
6 el domains, 3 bonding, 3 non-bonding
| geometry, bond angle
T-shaped, < 90°
142
6 el domains, 2 bonding, 4 non-bonding
linear, 180°
143
molecular vs electron domain geometry
molecular => greater repulsions are taken into account
electron domain => assumes all electron domains are bonding
