Chemistry Key Definitions (incomplete)

Atom

Positively charged nucleus containing most of the mass, surrounded by atomic shells with orbiting electrons of negative charge and negligible mass.

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Atomic Number

The atomic number of the nucleus shows the number of protons or electrons in the nucleus.

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Mass Number

The number of protons added to the number of neutrons (Mass of a proton and a neutron are both 1, the mass of an electron, 1/2000, is negligible).

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Isotope

An atom with the same number of protons but a different number of neutrons. OR same atomic number, different mass number Example - Carbon also exists as Carbon-14, which has a mass number of 14. The number of protons and electrons are unchanged but t

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Relative Atomic Mass

Average weighted mass of an atom compared with 1/12th of the mass of an atom of Carbon-12.

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Relative Isotopic Mass

Mass of an atom of an isotope compared with 1/12th of the mass of an atom of Carbon-12.

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Avogadro's Constant

The number of particles per mole (6.02e23 mol^-1)

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Molar Mass

Mass per mole of a substance, has the unit grams per mole (g mol^-1) Example - Carbon dioxide, CO2, has a molar mass of 44 g mol^-1 as Carbon has a mass of 12 and Oxygen has a mass of 16, there are two Oxygen atoms -> 12 + (16*2) = 44.

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Species

Type of particle that takes part in a chemical reaction.

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Stoichiometry

Molar relationship between relative quantities of substance in a reaction.

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Empirical Formula

The simplest whole number ratio of atoms of each element present in a compound.

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Molecular Formula

The actual number of atoms of each element in a molecule Concentration - The amount of solute (in mol) per 1dm^3 of solution, units mol dm^-3.

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Acids

Proton donors, release H+ ions into solution.

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Bases

Proton acceptors, they take H+ ions and neutralise acids.

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Alkalis

Soluble base which releases OH- ions in solution.

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Salts

Chemical compound formed from acid +metal.

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Anyhdrous

Substance containing no water molecules.

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Hydrated

Crystalline compound containing water molecules.

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Water of Crystalline

Water molecules which form essential part of the crystalline structure of a (hydrated) compound.

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Redox Reaction

A reaction in which species are both reduced and oxidised.

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Oxidation

Loss of electrons, increase in oxidation number shows that a species has been oxidised. Thing that's oxidised= reducing agent.

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Reduction

Gain of electrons, decrease in oxidation number shows that a species has been reduced. Thing that's reduced= oxidising agent.

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Disproportionation Reaction

A reaction in which the same species is both reduced and oxidised.

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Oxidation Number

Number of electrons that an atom uses to bond with atoms of another element

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First Ionisation Energy

The enthalpy change when each atom in one mole of gaseous atoms loses an electron to form one mole of gaseous 1+ ions. ENDO

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Second Ionisation Energy

The energy required to remove one electron from the outer shell from one mole of gaseous 1+ ions to form one mole of gaseous 2+ ions.

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Enthalpy of Formation

The enthalpy change when one mole of a compound is formed from its constituent elements in their standard states. usually EXO

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Lattice Enthalpy

The enthalpy change when one mole of a solid ionic lattice is formed from its gaseous ions under standard conditions. EXO

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Enthalpy of Atomisation

The enthalpy change when one mole of gaseous atoms is formed from its elements in its standard state. ENDO

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Electron Affinity

The enthalpy change when one mole of gaseous atoms each gains one electron to form one mole of gaseous -1 ions. EXO

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Second Electron Affinity

The enthalpy change when one mole of gaseous -1 ions each gains one electron to form one mole of gaseous -2 ions. ENDO

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Enthalpy of Solution

The enthalpy change that takes place when one mole of a compound is dissolved in water under standard conditions. usually ENDO

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Enthalpy of hydration

The enthalpy change that takes place when one mole of isolated gaseous ions is dissolved in water forming one mole of aqueous ions under standard conditions. EXO

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Nuclear Charge

The attraction from the protons in the nucleus with electrons. Increases along a period. A higher nuclear charge makes it more difficult to remove an electron from the atom’s outer shell, causing the ionisation energy to increase.

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Electron Shielding

The number of shells of electrons between the nucleus and the outer shell of electrons. Increases down a group. Electrons repel each other so the more electrons repelling the electrons in the outer shell, the easier it’ll be to remove them, causing t

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Atomic Radius

The distance from the outer shell of electrons to the nucleus of the atom. Increases down a group. A greater atomic radius means the outer shell electrons are under less influence from the attraction from the nucleus, this makes it easier to remove t

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Atomic Orbital

A region that can hold up to two electrons of opposite spins (Up and down). Orbitals in an s-subshell are spherical, orbitals in a p-subshell are hourglass shaped.

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Subshell

The space an electron can occupy within each shell. An s-subshell has one orbital (And so can hold a total of 2 electrons), a p-subshell has 3 orbitals (6 electrons), a d-subshell has 5 orbitals (10 electrons), a f-subshell has 7 orbitals (14 electro

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Periodic Table 'Blocks'

The blocks in a periodic table show which subshell the outer shell electron lies in. s-block is groups 1 and 2 (With Helium), p-block is groups 3 to 0 (Without Helium), d-block is the transition metals.

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Ionic Bonding

Electrostatic attraction between oppositely charged ions, where electrons are given and taken. Cation - +ve ion Anion - -ve ion

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Covalent Bonding

A strong electrostatic attraction between a shared pair of electrons & the nuclei of the bonded atoms.

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Dative Covalent Bond

A bond formed when one of the bonding atoms gives both of a pair of electrons.

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Metallic Bonding

The attraction of positive ions to delocalised electrons (Metals have 'seas' of delocalised electrons allowing them to form many metallic bonds).

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Lone Pairs (of electrons)

An outer-shell pair of electrons that isn't involved in chemical bonding.

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Bond Angles

The angle between bonds in a molecule, which are defined by the number of bonding pairs of electrons and lone pairs of electrons as the electrons repel each other and space out evenly. Lone pairs repel more than bonding pairs.

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Bond Angles 2

CO2 2 bonded, 0 lone= 180 Linear. H2O 2 bonded, 2 lone=104.5 Bent. BCl3 3 bonded, no lone=120 Trigonal Planar. NH3 3 bonded, 1 lone=107 Trigonal Pyramid. CH4 4 bonded, no lone=109.5 Tetrahedral. SF6 6 bonded, no lone=90 Octahedral

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IntERmolecular Forces

The forces of attraction between neighbouring molecules. The strength or amount of intermolecular forces are what affect the substance's melting/boiling points - Stronger intermolecular forces require more (heat) energy to break.

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IntRAmolecular Forces

The forces of attraction within the molecule itself.

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Electronegativity

The ability of an atom to attract the bonding electrons in a covalent bond An atom's electronegativity is defined by the number of electrons it has compared to the number of electrons the bonded atom has. If an atom is more electronegative than anoth

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Diploe-Dipole Interactions

An intermolecular force. Molecules with permanent dipoles allow for weak intermolecular bonds to be formed between the molecules - The d+ of one atom attracts the d- of an atom in another molecule, and that molecule does the same thing to other molec

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Van der Waal's forces

Attractive forces between induced dipoles in neighbouring molecules. Van der Waals are formed when the movement of electrons unbalances the distribution of the charge in the shells, this causes an instantaneous dipole to form - This in turn attracts/

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Hydrogen Bond

A hydrogen bond is a strong dipole-dipole interaction between; an electron deficient hydrogen atom on one molecule and a lone pair of electrons on a highly electronegative atom (The hydrogen is d+ and the other atom must be a d-) Hydrogen bonds are t

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Giant Ionic Lattice

Formed by the attraction of oppositely charged ions, each ion is surrounded by the oppositely charged ions and the ions attract each other to form a giant lattice. High melting and boiling points as they are held together by strong electrostatic forc

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Giant Covalent Lattice

Three dimensional structure of atoms bonded together by strong intramolecular covalent bonds. High melting and boiling point as high energies are required to break the strong covalent bonds. Not conductors of electricity as there are no free charged

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Giant Metallic Lattice

Lattices which contain ionised atoms in fixed positions with delocalised outer shell electrons which spread, and can freely move, throughout the structure. High melting/boiling point as the attraction between ions and electrons is strong and so a lot

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Simple Molecular Lattice

Three dimensional structure of molecules bonded together by weak intermolecular forces. Low melting and boiling points as the weak forces between the molecules require little energy to break them. Cannot conduct electricity as there are no charged pa

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Thermal Decomposition

One chemical substance breaking up with heat into at least two others.

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Displacement Reactions

Reaction in which a more reactive element displaces a less reactive element from an aqueous solution of the latter's ions.

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Precipitation Reactions

Formation of a solid from solution during a chemical reaction.

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Functional Group

The part of the organic molecule responsible for its chemical reactions.

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Homologous Series

A series of organic compounds with the same functional group but with each successive member differing by CH2 e.g. Alkanes, alcohols, esters etc.

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Aliphatic Hydrocarbons

Carbons joined together in a straight or branched chain.

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Alicyclic Hydrocarbons

Carbons joined together in a ring structure- but are not aromatic.

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Aromatic Hydrocarbons

At least one benzene ring in the structure.

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Structural Isomers

Molecules with the same molecular formula but with different structural formulas.

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Stereoisomers

Molecules with the same structural formula but with a different arrangement in space.

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E/Z Isomerism

Different groups attached to each carbon of a C=C bond may be arranged differently in space because of the restricted rotation of the double bond.

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Cis/Trans Isomerism

E/Z isomerism when the atoms are the same e.g. H2-C=C-Cl2

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Curly Arrow

Represents the movement of a pair of electrons in the breaking or formation of a covalent bond.

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Electrophile

A species that can accept a pair of electrons to form a covalent bond.

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Nucleophile

A species that can donate a pair of electrons to form a covalent bond.

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Heterolytic Fission

The breaking of a covalent bond with both of the bonded electrons going to one of the atoms, forming a cation & anion.

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Homolytic Fission

The breaking of a covalent bond with one of the bonded electrons going to each atom, forming two radicals.

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Radical

A species with an unpaired electron.

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Substitution reaction of Alkanes INITIATION

Cl2 -> Cl-Cl -> Cl¬ + Cl¬

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Substitution reaction of Alkanes PROPAGATION

1) CH4 + CL¬ -> ¬ClH3 + HCl 2) ¬CH3 + Cl2 -> CH3Cl +Cl¬ overall) CH4+Cl2 ->CH3Cl +HCl

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Substitution reaction of Alkanes TERMINATION

e.g Cl¬ + Cl¬ -> Cl2 or ¬CH3 + ¬CH3 -> C2H6 or ¬CH3 + Cl¬ -> CH3Cl

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simplest formulas- alkanes, cycloalkanes, halogenoalkanes, alkenes, cycloalkenes, halogenalkenes

A= CnH(2n+2), CA= CnH2n HA= CnH(2n+1) Ae= CnH2n, CAe= CnH(2n-2) HAe= CnH(2n-1)

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Primary alcohols

The -OH group is attached to a carbon with no alkyl group or is bonded to only one alkyl group.

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Secondary alcohols

The -OH group is attached to a carbon atom bonded to two alkyl groups.

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Tertiary alcohols

The -OH group is attached to a carbon atom bonded to three alkyl groups.

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Oxidation of alcohols -PRIMARY

Gentle heating produces an aldehyde. This must be distilled immediately. Strong heating produces an aldehyde, which is then immediately refluxed to form a carboxylic acid.

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Oxidation of alcohols -SECONDARY

Strong heating produces a ketone.

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Oxidation of alcohols -TERTIARY

Resistant to oxidation.

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Dehydration of alcohol

An alchol can be dehydrated to form an alkene. conditions: heated under reflux, acid catalyst.

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Halide substitution

general equation: ROH + HX -> RX +H2O haloalkane formed

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Nucleophillic Substitution

A chemical reaction in which an atom/ group of atoms is exchanged for a nucleophile.

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Hydrolysis

Chemical reaction in which water is the reactant.

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Alkali hydrolysis

Reactions where -OH is the reacting species.

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Organohalogen compound

Carbon chain with atleast one halogen atom attached. Uses- pesticides, solvents, refrigerants, flame retardents & making polymers.

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CFC's initiation

UV radiation initiates the breakdown. CFCl2 -> CF2+CL¬ + Cl¬

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CFC's propagation

1) O3 +Cl¬ -> ClO¬ +O2 2) ClO¬ + O -> O2 +Cl¬ overall) O3+ O -> 2O2

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CFC's termination

1) Cl! + Cl¬ -> Cl2 2) ClO¬ + ClO¬ -> Cl2O2 3) ClO¬ + Cl¬ -> Cl2O

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Nitrogen oxide radicals also...

...catalyse the breakdown of the ozone. Formed during lightning strikes and aircraft travel (heat&pressure) 1) O3+ NO¬ -> NO2¬ + O2 2) NO2¬ + O¬ -> O2 + NO¬

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Chemistry Key Definitions (incomplete)

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