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ocr chem > module 2 - foundations in chemistry > Flashcards

module 2 - foundations in chemistry Flashcards

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1
Q

relative mass and charge of a proton

A

mass - 1
charge - +1

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2
Q

relative mass and charge of a neutron

A

mass - 1
charge - 0

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3
Q

relative mass and charge of an electron

A

mass - 0 (negligible)
charge - -1

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4
Q

definition of an isotope

A

an atom of the same element with different number of neutrons and different masses

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5
Q

what is the atomic number

A

the number of protons and electrons

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6
Q

what is the mass number

A

number of protons and neutrons (take atomic number away from mass number)

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7
Q

how many protons, neutrons and electrons are there in the element Beryllium

A

protons - 4
neutrons - 5
electrons - 4

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8
Q

relative isotopic mass definition

A

mass of an atom of an isotope compared with 1/12 mass of carbon -12

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9
Q

relative atomic mass definition

A

weighted mean mass compared with 1/12 mass of carbon -12

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10
Q

how many protons, neutrons and electrons are in an ion of potassium

A

protons - 19
neutrons - 20
electrons - 18

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11
Q

relative atomic mass equation

A

[(% abundance x mass isotope 1) + (% abundance x mass isotope 2)] / 100

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12
Q

how can you guess the charge of an element

A

the group that the element is in

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13
Q

name the following ions:
NO3 -
CO3 2-
SO4 2-
OH -
NH4 +
Zn 2+
Ag +

A

nitrate ion
carbonate ion
sulfate ion
hydroxide ion
ammonium ion
zinc ion
silver ion

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14
Q

ammonium ion formulae

A

NH4 -

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15
Q

hydroxide ion formulae

A

OH -

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16
Q

nitrate ion formulae

A

NO3 -

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17
Q

nitrite ion formulae

A

NO2 -

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18
Q

hydrogencarbonate ion formulae

A

HCO3 -

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19
Q
A
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20
Q

carbonate ion formulae

A

CO3 2-

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21
Q

sulfate ion formulae

A

SO4 2-

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22
Q

sulfite ion formulae

A

SO3 2-

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23
Q
A
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24
Q

phosphate ion formulae

A

PO4 3-

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25
Q

difference between ammonia and ammonium

A

ammonia - NH3
ammonium - NH4

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26
Q

how to identify a compound containing metal + non metal + oxygen

A

ends in ‘ate’

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27
Q

how to identify a compound containing a metal + non metal

A

ends in ‘ide’

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28
Q

what does the bohr model tell us about atomic structure

A

electrons exist in energy levels around the nucleus

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29
Q

how many electrons are in the first 4 shells

A

2 , 8 , 18 , 32

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30
Q

how many electrons are in each subshell

A

s - 2
p - 6
d - 10
f - 14

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31
Q

what are the 4 orbitals that are within each subshell

A

s , p , d , f

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32
Q

what is an orbital

A

a region that can hold up to two electrons with opposite spins

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33
Q

what is the shape of s- orbitals

A

spherical

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34
Q

what is the shape of p- orbitals

A

elongated dumbells

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35
Q

electronic structure order?

A

1s,2s,2p,3s,3p,4s,3d,4p,4d,4f

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36
Q

electronic configuration of the atom Ni (28)

A

1s2,2s2,2p6,3s2,3p6,3d8,4s2

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37
Q

avagadro’s constant

A

3.02 x10 23

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38
Q

empirical formula definition

A

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

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39
Q

where is the s block

A

left side

40
Q

where is the d block

A

middle

41
Q

where is the p block

A

right

42
Q

where is the f block

A

bottom

43
Q

molecular formula definition

A

the number and type of atoms of each element in a molecule

44
Q

how to write an ionic equation

A

write full balanced symbol equation using state symbols
rewrite the equation ionically
identify and cancel spectator ions
rewrite ionic form of equation

45
Q

what compounds are soluble

A

nitrates
hydroxides if in group 1 or 2
sulphates if group 1 or 2
carbonates if lithium or sodium

46
Q

number of particles equation

A

number of particles = avagradros constant (NA) x moles

47
Q

water of crystallization definition

A

amount of water inside a hydrated salt

48
Q

hydrated definiton

A

contains molecules of water

49
Q

hydrochloric acid formulae

A

HCl

50
Q

sulfuric acid formulae

A

H2SO4

51
Q

nitric acid formulae

A

HNO3

52
Q

acetic acid formulae

A

CH3COOH

53
Q

sodium hydroxide formulae

A

NaOH

54
Q

potassium hydroxide formulae

A

KOH

55
Q

ammonia formulae

A

NH3

56
Q

difference between strong acid and weak acid

A
  • strong acid fully dissociates in aq solution
  • weak acid partially dissociates in aq solution
57
Q

difference between alkali and acid in water

A

Alkali: base that dissolves in water releasing OH-
ions into solution
Acid: acid that dissolves in water releasing H+ ions into solution

58
Q

preparing standard solution

A
  • Solid weighed + dissolved in beaker using less distilled water than needed to fill volumetric flask
  • Transfer to volumetric flask + last traces rinsed into flask with distilled water
  • Add distilled water drop wise until bottom of meniscus matches up with mark
  • Flask inverted slowly several times to mix, if not titration results will be
    inconsistent
59
Q

acid-base titration procedure

A
  • use pipette to add set volume of unknown solution to conical flask
    add a few drops of indicator (methyl orange or phenolphthalein) to conical flask
  • fill burette with known solution and make sure it’s levelled off at exactly 0ml
  • slowly add solution from burette to unknown solution in the conical flask
  • swirl the conical flask whilst solution from burette is being added
  • stop adding when reaction is complete (when there is an appropriate colour change)
  • record final volume on burette (tells us the volume of the solution added – titre)
  • repeat experiment for accuracy
60
Q

percentage yield calculation

A

(actual yield/theoretical yield) x 100

60
Q

atom economy calculation

A

Mr of desired product/ Mr of all products

60
Q

why aren’t yields ever 100%

A
  • product can be lost during filtration, evapouration, transferring etc
  • reversible reaction may not go to completion
  • impure starting chemicals
  • reaction incomplete
60
Q

why is it better to use a reaction with a higher atom economy

A
  • it uses fewer natural resources, produces less waste, and is better for the environment
60
Q

rules for assigning oxidation numbers to elements

A
  • the oxidation number of an element is always 0 as it isn’t bonded to another element
60
Q

rule for assigning oxidation numbers to compounds

A

each atom in a compound has an oxidation number, and it must always cancel out each other

60
Q

rule for assigning oxidation numbers to ions

A

the oxidation number must always add up to give the overall charge of the ion

60
Q

what does it mean when the name of an compound has a roman numeral

A

the element has that charge
e.g sodium chlorate (V)
this means that the sodium has a charge of 5

60
Q

what is oxidation

A

the loss of electrons

60
Q

what is reduction

A

gain of electrons

61
Q

ideal gas equation

A

pV= nRT

62
Q

oxidation in terms of electron transfer and changes in oxidation number

A
  • loses electrons
    -oxidation number increases
63
Q

reduction in terms of electron transfer and changes in oxidation number

A
  • gains electrons
  • oxidation number decreases
64
Q

redox reaction of acid with metal

A

metal + acid = salt + water

65
Q

ionic bonding definition

A
  • electrostatic attraction between positive and negative ions
66
Q

explanation of solid structures of giant ionic lattices

A
  • oppositely charged ions held together by strong ionic bonds in a huge three-dimensional structure
67
Q

how does ionic bonding affect the physical properties of ionic compounds

A
  • high mpt and bpt, electrostatic forces holding ionic lattice together are strong so require lots of energy to break
  • conduct electricity when molten, ions free to move and have delocalised electrons and mobile charge carriers that can conduct
68
Q

covalent bonding definition

A

strong electrostatic attraction between a shared pair of electrons and the nuclei of bonded atoms

69
Q

bond angle and number of lone pairs/bonding pairs of a linear molecule

A

180
bp = 2
lp = 0

70
Q

bond angle and number of lone pairs/bonding pairs of a bent molecule

A

104.5
bp = 2
lp = 2

71
Q

bond angle and number of lone pairs/bonding pairs of a trigonal planar molecule

A

120
bp = 3
lp = 0

72
Q

bond angle and number of lone pairs/bonding pairs of a triangular pyramidal molecule

A

107
bp = 3
lp =1

73
Q

bond angle and number of lone pairs/bonding pairs of a tetrahedral molecule

A

109.5
bp = 4
lp = 0

74
Q

bond angle and number of lone pairs/bonding pairs of a trigonal bipyramidal molecule

A

180 and 120
bp = 5
lp = 0

75
Q

bond angle and number of lone pairs/bonding pairs of a octahedral molecule

A

90
bp = 6
lp = 0

76
Q

electronegativity definition

A

atom’s ability to attract the bonding electrons in a covalent bond

77
Q

why does electronegativity increase across a period and decrease down a group

A

period - atomic radius decreases and charge density increases
group - shielding increases and atomic radius increases so charge density decreases

78
Q

hydrogen bonding definition

A

intermolecular bonding between molecules containing N,O,F or H

79
Q

anomalous properties of water

A
  • low density, open structure due to hydrogen bonded lattice
  • in liquid state, H bonds are constantly breaking and re-forming so a lattice structure isn’t created
  • bpt of water is high due to hydrogen bonds
  • hydrogen bonding = strongest intermolecular bonds
80
Q

why are simple molecular lattices solid structures

A
  • they are covalently bonded molecules attracted by intermolecular forces
81
Q

average bond enthalpy as a measurement of covalent bond strength

A

the amount of energy needed to break a covalent bond into gaseous atoms averaged over different molecules.

82
Q

how does electron pair repulsion affect shapes of molecules and ions

A
  • electrons are negatively charged and will repel other electrons when close to each other
  • in a molecule, bonding pairs of electrons will repel other electrons around the central atom, forcing the molecule to adopt a shape to minimise these repulsive forces
83
Q

how is a polar bond formed

A
  • in a covalent bond between 2 atoms of different electronegativities, bonding electrons are pulled towards more electronegative atom. this makes a polar bond
  • in a polar bond, difference in electronegativity between the two atoms causes a permanent dipole.
  • greater the difference in electronegativity, the more polar the bond.
84
Q

dipole definiton

A

= difference in charge between the two atoms caused by a shift in electron density in the bond.

85
Q

polar molecule and overall dipole in terms of permanent dipole and molecular shape

A
  • polar molecules have an overall dipole. arrangement of polar bonds in a molecule determines whether or not molecule will have a overall dipole.
  • in simple molecules, e.g hydrogen chloride, the one polar bond causes a single permanent dipole, which gives whole molecule an overall dipole.
  • more complicated molecules might have several polar bonds. shape of the molecule will decide whether or not it has overall dipole and so whether or not it will be polar. if polar bonds are arranged symmetrically so that the dipoles cancel each other out, e.g in carbon dioxide, then the molecule has no overall dipole and is non-polar.
  • if polar bonds are arranged so that they all point in roughly same direction, they won’t cancel each other out and charge will be arranged unevenly across whole molecule. molecule will have an overall dipole and so it will be polar.
86
Q

formation of temporary dipole-dipole forces

A

a result of the present of electrons in the molecule and the formation of temporary dipoles:
- sudden displacement of electrons in one atom to one side causes atom to have a temporary dipole
-Once temporary dipole is formed, it induces temporary dipoles in neighbouring atoms and molecules (like a domino effect).
-the two temporary dipoles are attracted towards each other by induced dipole- dipole forces also known as London or Van der Waals forces.

87
Q

effect of structure and bonding on the physical properties of covalent compounds with simple molecular lattice

A
  • simple molecules held together by weak intermolecular forces.
  • atoms within a molecule are held together by strong covalent bonds but individual molecules are held together by weak intermolecular forces of attraction.
  • compounds with simple molecular structure have low mpts and bpts.
  • when in its solid state, the simple molecules that make up the compound are arranged in a regular lattice held together by weak intermolecular forces.
  • as these interactions are very weak, not much energy is required to overcome them, which results in simple molecular structures usually being gaseous or liquid at room temperature.
  • can’t conduct electricity. This is because there are no mobile ions or electrons to carry the current.
88
Q
A

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