Thermochemical Principles NCEA Level 3

Aufbau Principle

'electrons fill orbitals from the lowest energy level 1 first and build filling each set of orbitals in turn'

Hund's rule

'electrons will fill each orbital of a sublevel before pairing'. An atom can have a full energy level, a full orbital, or a half full orbital (no paired electrons)

Atomic number

number of protons in the nucleus of an atom

Mass number

the number of protons + neutrons in the nucleus

Isotopes

are atoms of an element with different mass numbers

Electron configuration

places the electrons for any atom or ion into the energy levels, sublevels and orbitals they are found in.

Principal energy

levels known as K,L and M shells are numbered 1,2,3...n etc

Orbital

the region of space in which an electron is found

S orbital

consists of one orbital

P orbital

consists of three orbitals

D orbital

consists of five orbitals

F orbital

consists of seven orbitals

Atomic radius

the relative size of an atom is estimated by measuring the distance between the nuclei of two atoms which are joined in a covalent bond, half this distance is called the atomic radius

Positive ion

will be smaller than the corresponding atom (there is one less occupied energy level once the valence electrons are removed, with no change in the nuclear charge)

Negative ion

will be larger than the corresponding atom (there are more electrons in the valence level, so there is greater electron-electron repulsion, with no change in the nuclear charge)

Ionisation energy

IE (measured in kJmol-1) is the minimum energy needed to remove one electron from each atom in a mole of atoms in the gaseous state. The process is endothermic (/\H = +)

Electronegativity

is a measure of the attraction between a nucleus and a bonded pair of electrons. When there is an electronegativity difference between the atoms in a covalent bond, the bond becomes polar. Electronegativity increases from left to right across the periodic table and decreases going down the periodic table vertically

Covalent radius

a measure from the nucleus to the outermost electron in a covalent bond

Polyatomic ions

groups of atoms with an overall charge

Bond dipoles

occur when there is an electronegativity difference between the atoms in a covalent bond this creates a polar bond. Dipoles cancel out if they are equal in size but opposite in direction

Melting point

Melting point is a measure of the strength of forces between particles in a solid. A high melting point indicates that a large amount of energy is required to disrupt the solid lattice and so strong forces must exist between the particles in that solid

Boiling point

is a measure of the strength of inter-particle forces. A high boiling point indicates that a large amount of energy is required to break the strong forces between particles

Polar substances

will usually dissolve in water or other polar solvents due to similar attractions between the polar molecules or ions and the polar water molecules

Non polar substances

are generally soluble in non-polar solvents, such as hexane and cyclohexane, but are not soluble in water

Ionic bond

forms when two oppositely charged ions come together. The bonding force is due to the very strong electrostatic attraction between the positive ions (cations) and the negative ions (anions)

Covalent bond

the sharing of at least one pair of electrons between two atoms

Covalent network

solids the covalent bonds hold the atoms in a rigid 3-D structure. These substances have very high melting and boiling points, since a large amount of energy is required to break the many strong covalent bonds

Metallic solid

strong electrostatic forces of attraction, high melting and boiling point, malleable and ductile, delocalised electrons, hard and conducts electricity

Ionic solubility

will dissolve in polar solvents such as water due to the charged nature of their ions, and are usually insoluble in non-polar solvents

Molecule

is a group of atoms which can exist as an independent particle

Attractive forces

between molecules are known as intermolecular forces (or van der Waals)

London forces

or Temporary dipole-dipole attraction, exist between all molecules whether they are polar or non-polar, attractive forces are weak. An asymmetric electron distribution creates a temporary dipole with a partial negative charge on one side of the molecule and a positive partial charge on the other side.

Permanent dipoles

have permanent regions of positive and negative charge

Hydrogen bonding

occurs when a hydrogen atom is covalently bonded to another atom which has a high electronegativity so forms a very strongly polar bond to H, it also has to have a lone pair of electrons in the 2p orbital. Hydrogen bonds are the strongest of the intermolecular forces but they are still much weaker than covalent bonds within molecules

Polar molecules

are generally soluble in water and other polar solvents such as ammonia NH3 or ethanol C2H5OH. This is because there are similar attractive forces between the solute molecules and the solvent (water) molecules. Have an overall separation of the charge within the molecule called a dipole

Non polar molecules

such as iodine, carbon dioxide are generally insoluble in water, there is little attraction between the non-polar molecules and the polar molecules. Non-polar molecules will dissolve in non-polar solvents such as hexane because only the weak intermolecular forces exist between solute molecules.

Enthalpy change

= total heat content of products - total heat content of reactants

/\rH = sum H (products) - sum H (reactants

= heat absorbed or released during a reaction'

Exothermic

negative values for /\rH show an exothermic reaction. Energy is released to the surroundings, but lost from the substance, energy needed to break the bonds in the reactants is less than that given out when bonds form in the products

Endothermic

positive values for /\rH show an endothermic reaction;

Energy is taken in from the surroundings and gained by the substance, energy needed to break the bonds in the reactants is greater than that given out when bonds formed in the products

Spontaneous reactions

do not require input of heat energy to proceed, they are usually exothermic and occur because the loss of enthalpy drives the reaction'

Endothermic reactions

result in a gain of enthalpy which means that spontaneous endothermic reactions are unlikely

Entropy S

is a measure of the disorder in the reactants or products, change in entropy is /\S

Enthalpy of fusion

/\fus H (degree) the amount of heat required to change one mol of a solid to liquid at its melting point, is known as the enthalpy of fusion of a substance

Sublimation

the change of state from the solid state to the gas state without melting to form a liquid is called sublimation.

Enthalpy of sublimation

/\subH (degree) The amount of heat required to change one mol of a substance from solid to gas without the formation of the liquid state, is called the enthalpy of sublimation

Enthalpy of vaporisation

/\vap H (degree) the amount of heat required to change one mol of a substance from liquid to gas, at its boiling point, is called the enthalpy of vaporisation

Calorimeter

is how heat given out in a chemical reaction can be found, which also measures the change in temperature of a known volume of water.

Specific heat capacity

c. the specific heat capacity for water is 4.184 J g-1 (degree) C -1:

Heat energy (q) = mass of water (m) x heat capacity (c) x change in temperature (/\T)

Hess's law

the energy change in a chemical reaction is independent of the pathway taken

Standard state

of element is the normal state of the substance (s), or (l) or (g)

Standard enthalpy of formation

/\f H of a compound is the enthalpy change when one mol of the compound is formed from its elements, and the elements and compounds are each in their standard state

Standard enthalpy of combustion

/\c H is the heat involved when one mol of a substance is completely burned in oxygen, all reactants and products are in their standard state

Ionic radius

Distance from the center of an ion's nucleus to its outermost electron

Monoatomic

made up of only one atom

Isoelectronic

ions and atoms that have the same electron configuration

Periodicity

the similarities in characteristics and properties of elements based on their position in the periodic table

Lewis structures

are simple, two dimensional drawings showing how the atoms with their valence electrons are linked in a molecule

Linear

Shape: 2 regions of electron density on X, 2 bonds and no lone pairs. Bond angle 180. (e.g CO2, CS2, HCN). General formula XY2

Bent 120

Shape: 3 regions of electron density on X, 2 bonds and 1 lone pair. Bond angle 120. (e.g SO2) General formula XY2

Bent 109

Shape: 4 regions of electron density on X, 2 bonds and 2 lone pairs. Bond angle 109. (e.g H2O) General formula XY2

Trigonal Planar

Shape: 3 regions of electron density on X, 3 bonds and 0 lone pairs. Bond angle 120. (e.g H2CO, SO3) General formula XY3

Trigonal pyramid

Shape: 4 regions of electron density on X, 3 bonds and 1 lone pair. Bond angle 109. (e.g NH3) General formula XY3

T-shaped

Shape: 5 or 6 regions of electron density on X, 3 bonds and 2 lone pairs. Bond angle 90. (e.g ICl3) General formula XY3

Tetrahedral

Shape: 4 regions of electron density on X, 4 bonds and 0 lone pairs. Bond angle 109. (e.g CH4, CF4) General formula XY4

See saw

or distorted tetrahedron. Shape: 5 regions of electron density on X, 4 bonds and 1 lone pair. Bond angle 120 (e.g SF4) General formula XY4

Square planar

Shape: 6 regions of electron density on X, 4 bonds and 2 lone pairs. Bond angle 90. (e.g XeF4) General formula XY4

Trigonal bipyramid

Shape: 4 regions of electron density on X, 5 bonds and 0 lone pairs. Bond angle 120 and 90. (e.g PCl5) General formula XY5

Square pyramid

Shape: 6 regions of electron density on X, 5 bonds and 1 lone pair. Bond angle 90. (e.g ICl5) General formula XY5

Octahedral

Shape: 6 regions of electron density on X, 6 bonds and 0 lone pairs. Bond angle 90. (e.g SF6) General formula XY6

Diatomic molecule

will be non-polar if the two atoms have the same electronegativity and the bonded pair of electrons is shared equally.

Polyatomic

for molecules containing more than two atoms and polar bonds, the molecule will be non-polar if the bonds are arranged symmetrically around the central atom so that the bond dipoles cancel out

Metallic bond

is the force of attraction between all the valence electrons and the metal cations

Ductile

drawn out into a wire

Malleable

can be hammered and bent into shape

Ionic solid

3D lattice, high melting and boiling point, hard but brittle, greater solubility in polar solvents, conducts electricity when in molten or aqueous state but not in solid state

Polar covalent bond

A covalent bond between atoms that differ in electronegativity. The shared electrons are pulled closer to the more electronegative atom, making it slightly negative and the other atom slightly positive.

Polarisability

Indication as to an extent of which the electron cloud in a molecule can be distorted by a nearby electronic charge

Interatomic

forces between atoms

Intermolecular

forces between molecules

Interionic

forces between ions

Phase

of a substance is a recognisable form in which that substance exists

Allotropes

Different forms of the same element

Activation energy

Ea, is the amount of energy needed to initiate a chemical reaction

Catalyst

are substances which change the rate of a chemical reaction without themselves being used up, they provide an alternative pathway for the reaction