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Thermochemical Principles And The Properties Of Particles And Substances -…
Thermochemical Principles And The Properties Of Particles And Substances
Electron Configurations, Bonding and Periodic Trends
One mole of an element contains 6.023 x 10^23
Energy levels
First level has 1s orbital only
Second level has 2s and 2p orbitals
Third level has 3s, 3p and 3d orbitals
Fourth level has 4s, 4p, 4d and 4f orbitals
Aufbau principle - 'elections fill orbitals from the lowest energy level (1s) first and build filling each set of orbitals in turn'.
Hund's rule - 'electrons will fill each orbital of a sub level before pairing'.
Full energy level
A full orbital
Half-full orbital (no paired electrons)
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6
Remember that 4s electrons are lost before 3D electrons when transition metals ionise.
First Ionisation Energy
The energy needed in kJ mol^-1 to remove one electron from each atom in a mole of the element in its gas state.
Eg, As(g) —> As+(g) + e-
There is a decrease going down a group - the electron being lost is further away from the nucleus and has more shielding from the inner shells of electrons, so it is not held as strongly.
There is an increase going across a period - the number of protons increases and thus electrons are added in the same energy level and the effective nuclear charge (e.n.c.) acting on the outermost electron is stronger.
Filling an s sub shell orbital gives extra stability to elements He, Be, Mg, Ca, so their first ionisation energy is above the expected trend.
Half-filling a p sub shell gives extra stability to elements N and P, which thus have higher than expected values
Metals
Metals form metallic bonds, consisting of electrostatic attraction between the lattice of (positive) metal nuclei and the surrounding delocalised electrons (negative)
An increase in the number of valence electrons gives an increase in:
Strength of the metallic bond
Hardness
Density
Melting and boiling points
Electrons are able to slide past the nuclei and one another in liquid and solid states to make metals:
Malleable
Ductile
Electrically conductive
Metals and non-metals react to form ionic bonds.
Electrons are transferred from the metal and it forms a cation (positive ion); the non-metal that receives the electrons forms an anion (negative ion).
A strong electrostatic force of attraction occurs between the oppositely charged ions within a giant 3D lattice of the ions in a fixed ratio.
Ionic compounds have the following properties:
Greater solubility in polar solvents (eg, water) than in non-polar solvents
Hard but brittle salts
Conductors of electricity when in solution or molten (liquid)
High melting and boiling points
Transition metals
Transition metals are elements in groups 3-12 (d block). They always lose the 4s electrons first when forming ions.
Chromium has greater stability with half-filled 4s and 3d orbitals.
Copper has greater stability with half-filled 4s and full 3d orbitals.
Other properties
Transition metals form coloured ions - incomplete d orbital electrons absorb light of different wavelengths; colours seen are wavelengths not absorbed
Transition metals form complex ions - a central metal cation is bonded to a fixed number of ligands (small molecules (eg, H2O, NH3,) or ions (eg, Cl-, CN-) with lone pairs of electrons that can datively bond to the metal cation). Ligands often change the colour absorption of a transition metal.
Transition metals have variable oxidation states - due to having spare 3d electrons they can share (eg, chromium +3 and +6; manganese +2, +4; copper +1 and +2)
The pH of complex ions with water molecules is acidic as one of the waters acts as a proton donor:
[Fe(H2O)6]3+(aq) --> [Fe(H2O)5OH]2+(aq) + H+(aq)
Chromium(VI) exists as yellow chromate in alkaline solutions and orange dichromate in acidic solution:
2CrO4,2-(aq) + 2H+ --> Cr2O7,2-(aq) + H2O(l)
Non-metals
Non-metals form covalent bonds with each other by donating one (or more) electron(s) each to a shared bond pair of elections that help make up a full valence electron shell.
The strong force of attraction of the positive nuclei to the shared bond pair is greater than the repelling effect of the two positive nuclei.
Bond strength decreases down a group - the bond pair becomes further from the nucleus and shielding from inner election shells increases.
Double bonds are shorter and stronger than single bonds.
Electrons in a covalent bond between different atoms will be shared unequally due to a different in each atom's electronegativity (measure of an atom's attraction for a bond pair of electrons). The uneven distribution of the electrons creates a polar covalent bond between the atoms and a permanent dipole across the bond.
Electronegativity is a measure of an atom's ability to attract and hold an electron from a covalent bond.
Bonds
Covalent Bonds
Covalent bonds in a molecule create a dipole across the entire molecule unless they cancel each other out because they are arranged symmetrically.
Only non-metals
Intermolecular forces involve weak forces of attraction between molecules.
Temporary dipole-dipole attractions or dispersion forces
Temporary uneven distribution of electrons around an atom/molecule repel electrons in a nearby atom/molecule inducing a small dipole in both
More likely to occur in larger atoms/molecules that contain more electrons, meaning atoms with higher numbers of electrons will have more temporary dipole-dipole attractions
Permanent dipole-dipole attractions
Occur between the opposite ends of molecules with a permanent dipole across the molecule.
The more polar the molecule, the strong the attraction.
Hydrogen bonds
Hydrogen bonds are the strongest attraction force (but are still only about 10% of the strength of a covalent bond).
A molecule must contain a hydrogen atom bonded to a fluorine, oxygen or nitrogen atom.
Alcohols and carboxylic acids have hydrogen bonds due to the O-H bond
This bond has a very large dipole that is attracted to the opposite end of the same bond on a nearby molecule (eg, water has a much higher melting and boiling point than molecules of similar size and polarity).
Metallic bonds
Ionic bonds
The periodic table
Atomic number increases across a period - number of protons increases
Atoms get smaller across a period - effective nuclear charge experienced by electrons in the same shell increases, pulling them in closer.
Electrons being added to the same energy level causes an increased charge of attraction from the nucleus (effective nuclear charge) acting on the same energy level outer electrons which decreases the atomic radius
The increasing effective nuclear charge also means that more energy is required to remove the outmost electrons, and so the 1st ionisation energy increases moving across the period.
Atoms get larger down a group - number of electron levels (shells) increases; effective nuclear charge shielded from outer electrons by inner electrons.
Allows electrons to move further way as the protons in the nucleus have less attractive pull on them.
Ions formed by elements on the left are smaller than their atoms because they have lost electrons to empty an energy level (shell) and there is less repulsion between fewer valence electrons.
Ions formed by elements on the right are larger than their atoms because they have gained electrons so there is increases repulsion between valence electrons
electronegativity is a measure of an atom's ability to attract and hold an electron from a covalent bond. Moving down a group, electronegativity decreases due to the outermost bonding electrons being in progressively higher energy levels that are further from the attractive pull of the protons in an atom's nucleus and which also experience shielding of this attraction by inner electrons on the atom.
Molecules!
BrF3
T-shaped
PCl6,-
Octahedral
AsF3
Trigonal Pyramidal
AsF3 has 4 regions of negative charge density surrounding the central As atom, which will all repel in a 3D space to form a base tetrahedral shape. However, only 3 of these regions are bonded, so the shape seen is trigonal bopryamidal base shape. All 5 of these regions are bonded, so the shape seen is also teugonal bipyramidal.
The As-F bond is polar in both molecules due to the difference in electronegativity of the As and F atoms. However, in AsF3 the molecule is symmetrical in shape, so the effect of the bond dipoles is cancelled out and the overall molecule is non-polar.
AsF5
Trigonal Bipyramidal
SF4
See-saw
SF3-
T-shaped
ClF5
Square pyramidal
Shape is square pyramidal. This is because there are 6 regions of negative charge density surrounding the central atom, that all repel each other in 3D space; 5 of these regions are bonded and one lone pair, so the shape that is observed is square pyramidal. The Cl-F bonds are polar due to the difference in electronegativity of the atoms causing the bonding electrons to be shared unequally. The molecule shape is unsymmetrical, so the effects of the polar bonds cause the overall molecule to be polar too.
BrF5
Square pyramidal
XeF4
Square planar
the molecular shape is due to the number of regions of negative electron charge density repelling each other in #D XeF3 has 6 regions of electrons, 2 lone pairs and 4 bonded pairs. The 2 lone pairs occupy space around the central Xe atom but are (unseen). Thus the 4 electron pairs with bonding are seen as the square planar shape. XeF4 is non-polar. The electronegativity difference between Xe and F atoms causes the bonding electrons to be shared unequally - making the internal bonds of the molecule polar. However, the molecule shape is symmetrical, which causes the effect of the polar bonds to be cancelled, leaving the overall molecule non-polar.
ICl4,-
Square planar
ClF3
T-shaped
SeF6
Octahedral
SeF6 has six regions of negative charge equally repelling each other and surrounding the central atom. All six regions are bonded, so SeF6 is octahedral in shape
Each of the Se-F bonds is polar due to the difference in electronegativity between the atoms; however, the molecule is symmetrical, so the effect of the polar bonds is cancelled to leave a non-polar molecule. As a non-polar molecule, SeF6 would be unable to form attractions with polar water molecules that would otherwise enable the SeF6 to mix and dissolve in water, as so SeF6 is insoluble in water.
I3,-
Linear
The I3,- ion has five regions of negative electron density surrounding the central I atom, equally repelling each other in 3D. Two of theses regions are bonded to I atoms while the other three regions involve lone pairs of electrons, so the shape seen is linear.
IF5
Square pyramidal shape
Square pyramidal as six regions of negative electron density equally repelling in 3D; five of these regions are bonded. Six regions of negative electron density, only five of which represent bonding results in an unsymmetrical molecule shape. The I-F bond is polar due to a difference in electronegativity between the I and the F atoms, causing the bonding electrons to unequally be shared. In the unsymmetrical shape the effect of the polar bonds does not cancel out and thus they cause the molecule to also be polar.
Lewis Diagrams, Shapes And Polarity
Apart from Si, P and S, the first 18 elements (as well as many other elements) of the periodic table obey the Octet Rule.
Si can have 12
P can have 10
S can have 10 or 12
Large atoms, especially sulfur and phosphorus, can expand their octet by using their empty d orbitals to accomodate extra electrons
I can have 10 and 12
As can have 10
Lewis diagrams of ions carry the charge on the central atom expect for oxyanions (eg, nitrate, carbonate, sulfate)
Shapes
Linear
2 sets of electrons
Approximately 180°
Trigonal planar
Trigonal planar
120°
3 sets of electrons
Bent (3 regions)
3 sets of electrons, 1 lone pair
120°
Tetrahedral
Tetrahedral
109°
4 sets of electrons, 0 lone pair
Trigonal pyramidal
4 sets of electrons, 1 lone pair
Approximately 109°
Bent (4 regions)
4 sets of electrons, 2 lone pairs
Approximately 109°
Trigonal bipyramidal
Trigonal bipyramidal
90° and 120°
5 sets of electrons
See-saw
5 sets of electrons, 1 lone pair
90° and 120°
T-shaped
5 sets of electrons, 2 lone pairs
90° for bonded
Octahedral
Octahedral
6 sets of electrons
90°
Square pyramidal
6 sets of electrons, 1 lone pair
90°
Square planar
6 sets of electrons, 2 lone pairs
90°
Equations
Hess's law: ΔcH°=ΣΔfH°(products)-ΣΔrH°(reactions). 'The energy change for a reaction is independent of the route taken'
q = mcΔT
where q = energy change, m = mass of water used in g, ΔT = change in temperature of the water, and c = the specific heat capacity of the liquid (for water c = 4.18 Jg^-1 °C^-1)
N = m/M
q = n*ΔrH°
ΔcH = q/N
Energy Changes
Enthalpy, ΔH
Change in enthalpy in chemical reactions is represented as change in heat energy or ΔH, measured in kJ mol-1
If energy is released, the reaction is exothermic and ΔH is a negative value.
If energy is absorbed, the reaction is endothermic and ΔH is a positive value.
Enthalpy of reaction (ΔrH) is the enthalpy change for any given reaction
ΔrH° is the standard enthalpy of reaction, used when the reaction is performed under standard conditions (1 atmosphere of pressure, 298K or 25°C).
The enthalpy does not include the activation energy for a reaction.
To change from a solid to a liquid (melt), or from a liquid to a gas (evaporate/boil), a substance must overcome the forces between the particles within the solid or liquid. At melting and boiling points, the heat energy being added is used to break the forces between the particles instead of increasing the temperature.
Types of substances
Ionic salts - Have a strong 3D lattice of ionic bonds between the ions - requires a large amount of energy to break; melting points and boiling points generally high.
Metals have a strong 3D lattice of nuclei mutually attracted to surrounding valence electrons within the metal structure - requires a large amount of energy to break, mps and bps are generally high.
Non-polar covalent molecules have very weak temporary dipole-dipole attractions between molecules - need very little energy to break, mps and bps are generally very low (but the bigger the mass the higher the melting point and boiling points).
Polar covalent molecules - energy needed to break the forces depends on the strength of dipole and intermolecular attractions; melting and boiling points vary, but are low compared with those of ionic salts and metals.
3D giant covalent molecules have very strong covalent bonds within the molecule - requires a large amount of energy to break, melting and boiling points are generally very high.
Enthalpy of sublimation, ΔsubH°
Energy needed to sublime 1 mole of solid to gas at its sublimation point.
Standard enthalpy of formation, ΔfH°
This is the energy change when 1 mole of a substance is formed from its constituent elements. The ΔfH° for any element (eg, H2, He) in its elemental state is zero.
A formation equation is written and balanced so that only 1 mole of the new substance is formed.
Standard enthalpy of combustion, ΔcH°
This is the energy change when 1 mole of a substance is combusted completely in oxygen.
A combustion equation is written and balanced so there is only 1 mole of the combusting substance.
Incomplete combustion makes the experimental less that the theoretical value
Entropy, S
Entropy is a measure of the level of disorder or randomness of the particles in a system.
Eg, I2(s) --> I2(g)
Entropy increases
A solid melts into a liquid
A solid dissolves into an aqueous solution.
A solid sublimes into a gas.
A liquid vaporises into a gas.
A reaction with more mol of products than reactants. There has been an increase in the total amount of particles during the reaction - this increases the level of disorder.
Entropy decreases
A gas condenses into a liquid
A liquid freezes into a solid
Ions in solution form a precipitate. Free-moving ions in solution have become locked in fixed positions in a solid, so there is less disorder.
A substance reduces in the temperature. The particle movement has decreased, so there is less random movement or less disorder in the system.
A reaction with more mol of reactants than products. There are fewer total particles in the system, so the level of disorder has decreased.
Both entropy and enthalpy need to be considered in the total energy change of a spontaneous reaction.
A net release in energy causes a reaction to be spontaneous.