2.1 Atoms And Reactions Flashcards

1
Q

Relative Isotopic Mass

A

The mass of an atom of an isotope compared to one-twelfth of the mass of an atom of C-12 isotope

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2
Q

Amount of substance

A

Number of particles, measured in moles

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3
Q

Avagadro’s constant, Na

A

Number of particles per mole of a substance, 6.02 x10^23 mol^-1

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4
Q

Molar mass

A

The mass per mope of a substance in gmol^-1 = RFM

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5
Q

Mole

A

Amount of any substance with as many particles as there are C-atoms in exactly 12g of C-12 isotope,
Moles = Mass / RFM

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6
Q

Isotopes

A

Atoms of the same element with the same number of protons but different numbers of neutrons, and different masses

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7
Q

Acid

A

Proton donor - release H+ ions (protons) in solution, eg H2SO4, HCl, HNO3, CH3COOH (weaker)

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8
Q

Base

A

Proton acceptor - readily accepts H+ ions from acid in aq solution

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9
Q

Alkali

A

A soluble base that releases OH- ions in aq solution, eg NaOH, KOH, NH3

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10
Q

Neutralisation reaction ionic equation and observations

A

H+ (aq) + OH- (aq) –> H2O (l)

Metal dissolves, fizzing- CO2 (bubble through limewater cloudy to test), -H2 (lighted splint squeaky pop to test), temp rise and pH change?

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11
Q

Acid and carbonate forms…

A

Salt and carbon dioxide and water

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12
Q

Acid and metal oxide forms…

A

Salt and water

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13
Q

Acid and alkali forms…

A

Salt and water

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14
Q

Acid and metal forms…

A

Salt and hydrogen

Redox!

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15
Q

Salt

A

Produced when H+ ion of an acid is replaced by a METAL ion or another positive ion eg NH4+ ion

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16
Q

Hydrated

A

A crystalline compound containing water molecules

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17
Q

Anhydrous

A

A substance that contains no water molecules (the form without water)

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18
Q

Water of crystallisation

A

Water molecules that form an essential part of the crystalline structure of a compound

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19
Q

What do you need to work out empirical formula of hydrated compound?

A

Mass of anhydrous salt, and mass of WofC

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20
Q

Oxidation Number

A

A measure of number of electrons that an atom uses to bond with atoms of another element

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21
Q

Oxidation number rules

A
  • -
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22
Q

Redox

A

Reaction where both oxidation and reduction occur

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23
Q

Oxidation

A

Loss of electrons, or increase in oxidation number

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24
Q

Reduction

A

Gain of electrons, or decrease in oxidation number

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25
Q

Relative Atomic Mass

A

Weighted mean mass of an atom of an element compared to one-twelfth of the mass of an atom of C-12

26
Q

List soluble salts

A
All nitrates, all sodium/potassium/ammonium salts, 
Most chlorides (except Pb/Ag) 
Most sulphates (except Pb/Ba/Ca)
27
Q

List of insoluble salts

A

Most oxides/hydroxides/carbonates (except Na/K/NH4+),
Lead/silver chloride
Lead/barium/calcium sulphate

28
Q

Molar gas volume

A

Volume per mole of a gas, in dm^3mol^-1,

Moles of gas = vol dm3/24

29
Q

Ideal gas law

A

pV=nRT

Pa x m^3 = JK^-1 x K

30
Q

Why is 100% yield rare

A
Reaction at equilibrium, 
Other side reactions- byproducts, 
Impure reactions, 
Some r/p left behind in apparatus 
Separation and purification- lose product
31
Q

Benefits for sustainability of high atom economy

A

More desireable product- less waste/toxic product (safer, less pollution),
Saves landfill and protects environment,
Expensive to dispose of waste,
Less separation costs,
Less expensive reactant wasted,
More sustainable (saves limited resources/materials/energy)

32
Q

Strong acid

A

Dissociates completely in solution into constituent ions; very little of reverse reaction happens, so nearly all H+ ions are released

33
Q

Weak acid

A

Only partially dissociates in solution into constituent ions; reverse reaction favoured so few H+ ions released

34
Q

Shell

A

Group of atomic orbitals with same principal quantum number; shells further from nucleus have greater energy level than those closer to nucleus

35
Q

Each shell holds up to … (electrons)

A

2n^2 electrons where n is the shell number

36
Q

Each subshell holds up to … (electrons)

A

s- 1orbital, 2electrons,
p- 3orbitals, 6electrons,
d- 5orbitals, 10electrons,
f- 7orbitals, 14electrons

37
Q

Subshell

A

Group of same type of atomic orbitals (s/p/d/f) within a shell

38
Q

Orbital

A

Region around nucleus that can hold up to 2 electrons with opposite spins

39
Q

First ionisation energy

A

Energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous 1+ ions

40
Q

Factors influencing ionisation energy

A

Shells, shielding, nuclear charge and nuclear attraction

41
Q

Why do successive IEs increase within each shell

A

Electrons are being removed from an increasingly positive ion- less repulsion among remaining ions so held more strongly by nucleus

42
Q

Why are there big jumps in IE

A

A new shell is being broken into - an electron is being removed from a shell closer to the nucleus

43
Q

Ionic bonding

A

Electrostatic attraction between oppositely charged ions (positive and negative ions)

44
Q

Describe giant ionic lattices

A

millions of ions, each surrounded by ions of opposite charges, ions attract each other strongly in all directions, always solid

45
Q

Properties of giant ionic lattices

A

high mp and bp - solid at room temp, lots of energy needed to break strong es forces between opp charged ions in sold lattice;
do not conduct as solids - ions in fixed positions, not free to move and carry charge;
conduct when molten/dissolved - ions free to move and carry charge;
soluble in polar molecules - eg water molecules surround each ion, which are pulled out of lattice, to form a soln (lattice breaks down)

46
Q

Covalent bond

A

Strong electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atoms; directional so only acts between the 2 atoms involved

47
Q

Lone pair

A

Outer shell pair of electrons not involved in chemical bonding; repels more than bp (2.5’ reduced due to extra repulsive effect)

48
Q

Bond pair

A

Pair of electrons shared between 2 atoms; when only bp’s, all pairs repel each other equally

49
Q

Dative covalent bond

A

Shared pair of electrons where both electrons have been provided by one of the bonding atoms only

50
Q

Electronegativity

A

Ability of an atom to attract the bonding electrons in a covalent bond; increases towards F in the periodic table

51
Q

Polar covalent bond (in terms of EN)

A

Bond between different molecules with a small difference in electronegativity

52
Q

Ionic bond (in terms of EN)

A

Bond between different elements with a large difference in electronegativity

53
Q

Electron pair repulsion theory CHECK

A

Shape of molecule is determined by number and arrangement of electron pairs in outer shell of central atom and hence the repulsion between them

54
Q

Bond angle in linear molecule

A

180’

55
Q

Bond angle in trigonal planar molecule

A

120’

56
Q

Bond angle in tetrahedral molecule

A

109.5’

57
Q

Bond angle in octahedral molecule

A

90’

58
Q

Bond angle in pyramidal molecule

A

107’

59
Q

Bond angle in non linear molecule

A

104.5’

60
Q

Describe simple covalent molecules

A

strong covalent bonds within molecules; weak IMF (induced DD) between molecules break when they change state

61
Q

Properties of simple covalent molecules

A

low mp and bp - induced DD;
don’t conduct - no charged particles to move/carry charge;
soluble in non polar solvents - form similar strength attractions (IDD) weakens lattice