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Module 2 Chemistry > Module 2 > Flashcards

Module 2 Flashcards

Basics in Chemistry

1
Q

what is atomic number and atomic mass?

A

number= number of protons
mass= number of protons and neutrons

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2
Q

charge, mass and position of subatomic particles?

A

proton= nucleus, +1 charge, 1 mass
neutron= nucleus, 0 charge, 1 mass
electron= shells/ energy levels, -1 charge, 1/1846 mass

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3
Q

state the forces present within an atom

A

electrostatic forces of attraction between positive nucleus and negative electrons in shells

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4
Q

what are isotopes?

A

atoms of the same element with different number of neutrons and different atomic masses

-can be seen as O^16 and O^17, where atomic number is ommited as is same in both.

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5
Q

what is relative atomic mass?

A

the weighted mean mass of an atom of an element compared to 1/12th the mass of an atom of Carbon-12

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6
Q

equation for calculating relative atomic mass?

A

sum of (relative mass x percentage abundance) / total abundance

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7
Q

how can you determine the percentage abundance of an isotope in a sample

A
  • using a mass spectrometer:
    1) place sample in mass spectrometer
    2) sample is vaporised and ionised to form positive ions
    3) ions or accelerated, where heavier ions move more slowly and are more difficult to deflect than lighter ions, so ions of each isotope are separated
    4) ions are detected on a mass spectrum as a mass to charge ratio, and add to signal
    5) greater the abundance, greater the signal
  • peak height/total height of all peaks x 100
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8
Q

what is relative isotopic mass

A

mass of an atom of an isotope relative to 1/12th the mass of an atom of carbon-12

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9
Q

why do isotopes still have the same chemical properties?

A
  • the number of electrons don’t change
  • it is electrons involved in chemical reactions (neutrons have no impact)
  • so different isotopes of the same element react in the same way
  • (may be small difference in physical properties, e.g. higher melting/boiling point and density for higher relative mass isotopes)
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10
Q

what are cations

A
  • positive ions
  • have fewer electrons than protons
  • e.g. 2+
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11
Q

what are anions

A
  • negative ions
  • have more electrons than protons
  • e.g. 2-
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12
Q

rules for naming compounds

A
  • if two elements bonded only (BINARY COMPOUNDS), metal goes first and ends with -ide
  • if element contains oxygen, must end in -ate
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13
Q

which elements are diatomic

A
  • gases (O2, H2, N2)
  • group 7 halogens (F2, Cl2, Br2, I2)
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14
Q

common formulas you should know

A
  • CO2 = carbon dioxide
  • CO = carbon monoxide
  • H2O = water
  • CH4 = methane
  • NH3= ammonia
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15
Q

what are polyatomic ions

A

ions made up of more than 1 element bonded together

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16
Q

why are showing ions of transition metals more difficult, and how are they displayed

A
  • transition metals can form several ions with different charges
  • show the ionic charge using roman numerals in brackets
  • e.g. Fe (III) = Fe3+
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17
Q

sulfate ion

A

SO4 2-

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18
Q

nitrate ion

A

NO3 -

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19
Q

hydroxide ion

A

OH-

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20
Q

carbonate ion

A

CO3 2-

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21
Q

ammonium ion

A

NH4 +

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22
Q

phosphate ion

A

PO4 3-

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23
Q

bicarbonate ion

A

HCO3 -

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24
Q

silver ion

A

Ag+

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25
Q

zinc ion

A

Zn 2+

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26
Q

phosphor and sulfur ions

A

P4, S8

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27
Q

how to balance equations with state symbols

A
  • must have same number of atoms of each element of either side
  • (s) = solid
  • (l) = liquid
  • (g) = gas
  • (aq) = aqueous, dissolved in water
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28
Q

what are ionic equations

A
  • show the reacting ions only
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29
Q

how do you make ionic equations

A

1) anything aqueous can be split into ions
2) ions that don’t change can be cancelled out
3) left with just reaction ions
4) REMEMBER STATE SYMBOLS

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30
Q

what is a shell

A

group of atomic orbitals having the same principal quantum number n

  • regarded as energy levels, and energy increases as the shell number does
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31
Q

what is an orbital

A

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

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32
Q

what is a sub-shell

A

group of the same type of atomic orbitals within a shell (e.g p subshell contains 3 p orbitals)

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33
Q

characteristics of each orbital

A

s - orbital: sphere shape, 1 orbital (2 electrons)
p - orbital: dumb-bell shape, 3 orbitals (px, py, pz) in sub-shell (6 electrons)
d - orbital: 5 orbitals in sub-shell, (10 electrons)
f - orbital: 7 orbitals in sub-shell, (14 electrons)

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34
Q

how does increasing shell number effect the location of the orbitals

A

the greater the shell number, n, :
- the greater the radius of its s-orbital
- the further its p-orbital is from the nucleus

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35
Q

what are the max number of electrons in each shell 1-4

A

2 (s orbital only)
8 (s and p orbital only)
18 (s,p and d orbitals only)
32 (2,p,d and f orbitals)

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36
Q

how do electrons fill shells

A
  • fill sub-shells in order of increasing energy
  • 4s is filled before 3d:
  • 4d subshell is at a lower energy level than 3d subshell
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37
Q

how to draw electrons in orbitals as boxes

A
  • 1 orbital = shown as box
  • electrons = shown as arrows (2 in same one are in opposite spins, as this helps to counteract the repulsion between the 2 negatively charged electrons)
  • electrons only pair when no empty orbitals left in subshell (prevents repulsion between paired electrons until there is no further empty orbital available at same energy level)
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38
Q

what are all the blocks of the periodic table

A
  • determined by the highest energy sub-shell where outermost electrons are
  • s block = group 1, 2 + hydrogen and helium
  • p block = group 3, 4, 5, 6, 7, 8
  • d block = middle block ( 4s BEFORE 3d)
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39
Q

what order are shells filled

A

1s2, 2s2, 2p6, 3s2, 3p6, 4s2, 3d10, 4p6

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40
Q

what are the two anomalies to electron structure

A

Chromium= 4s1, 3d5 (instead of 4s2, 3d4)
- 4s subshell gives electron to 3d subshell to give a fully 1/2 filled subshell, giving extra stability

Copper = 4s1, 3d10 (instead of 4s2, 3d9)
- 4s gives one electron to 3d subshell to give a full outer shell, providing extra stability

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41
Q

how do you write a shorthand electronic configuration

A
  • use noble gas before in brackets and add on additional subshell electrons
  • can also be done with noble gases themselves, using the one before
  • e.g. 11-Na : [Ne] 3s1
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42
Q

how do you write electron configuration for ions

A
  • consider the amount of electrons and do accordingly
  • for ions which lose electrons, remember to remove 4s subshell electrons first
    -e.g. chloride ion = [Ne] 3s2, 3p6 = [Ar]
  • zinc ion = [Ar] 3d10 (you take away the 2 on 4s first)
  • remove 4s first because the energies of 4s and 3d are close together, and once filled, 3d falls below 4s energy level
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43
Q

what is the meaning of isoelectronic

A
  • containing the same amount of electrons, like between an ion and atom
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44
Q

what is a mole

A

amount of any substance containing as many particles as there are carbon atoms in exactly 12g of Carbon-12

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45
Q

what is avogadro’s constant

A

the number of atoms per mole o the Carbon-12 isotope = 6.02 x 10^23

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46
Q

why is the term “particles” important when talking about moles

A

1 mole of H : 1 mole of H atoms
1 mole of H2 : 1 mole of H molecules (so x2 for number of atoms)

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47
Q

what is molar mass

A

mass per mole of a substance (gmol^-1)

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48
Q

equation linking mass, molar mass, moles

A

mass (g) = moles x molar mass (gmol^-1)

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49
Q

what is avogadro’s law

A

at room temperature and pressure, 1 mol of gas occupies 24dm^3

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50
Q

equation linking volume and moles

A

volume (dm^3) = 24 (mol dm^-3) x moles

24 = the molar gas volume (volume per mole of gas molecules)

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51
Q

what is concentration

A

quantity of a substance, in moles, in a given volume

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52
Q

what is an ideal gas

A

a gas with:
- no intermolecular forces
- particles of random motion
- particles with negligible size
- elastic collisions

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53
Q

what is the ideal gas equation

A

pV= nRt

p= pressure (Pa/ pascals)
V= volume (m^3 (x10^-6 to dm))
n= moles
R= ideal gas constant, (8.314 Jmol-2K-2)
t= temperature (K, +273 to C)

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54
Q

what is molecular formula

A

the number of atoms of each element in a molecule

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55
Q

what is empirical formula

A

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

  • useful for giant crystalline structures such as ionic structures, metals, giant covalent structures, where actual number of atoms would be VERY large
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56
Q

what is relative molecular/formula mass

A

compares the mass of a molecule/formula unit with the mass of Carbon-12
(add up all RAMs)

  • molecular = for covalent structures
  • formula = for (giant) ionic compounds, where you use the empirical formula
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57
Q

how do you calculate empirical formula

A

1) find the moles of each element present
2) divide by the smallest value of moles present
3) use to calculate the ratio of elements, therefore formula

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58
Q

what is hydrated salt

A

crystalline compound containing water molecules

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59
Q

anhydrous salts

A

salt containing no water molecules

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60
Q

water of crystallisation

A

-H2O: water molecules in hydrated salts

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61
Q

explain PAG 1: mass change in crucible to find water of crystallisation

A

1) weigh an empty crucible
2) add hydrated salt and reweigh
3) place crucible onto pipe-clay triangle on top of tripod
4) heat gently for 1 minute and strongly for 5 minutes
5) leave crucible to cool and reweigh
6) heat crucible again for 1 minute, leave to cool and reweigh
7) repeat until no mass change- CAREFUL TO NOT FOR TOO LONG, MAY DECOMPOSE SALT

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62
Q

what are the assumptions made in the water of crystallisation practical, and how does this effect the accuracy

A

1) you assume all the water has been lost:
- can use colour change in compound, but you can still only see the surface, and there may be extra water inside
- if similar colour compounds, very difficult
- SO, heat to constant mass

2) the salt may decompose further when heaters:
- could cause a colour change (copper (II) sulfate turns into black copper (III)
- very difficult to judge if no colour change occurs when decomposed

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63
Q

how are molar ratios used

A
  • used to work out limiting reagents (reactant not in excess which will be used up and stop the reaction)
  • always base calculation off limiting reagent
64
Q

equation for percentage yield

A

actual yield / theoretical yield

65
Q

what is theoretical yield, and what can it be prevented by

A

maximum possible amount of product:
- reaction not completed
- side reactions taken place instead
- purification causing loss of some product

66
Q

what is the limiting reagent

A

the reactant that is used up first and finishes the reaction (not in excess)

  • find out by finding moles of all reactants, and comparing to how many moles there SHOULD be present using equation
67
Q

what is atom economy and explain how it is used

A

a measure of how well atoms have been utilised

  • makes up part of argument when discussing which reactions to use, BUT, may still use low AE if
  • reactants are readily available and won’t require much energy to get
  • the percentage yield is high
68
Q

equation for atom economy

A
  • Mr of desired product/ Mr of all products
69
Q

what are the pros of high atom economy experiments

A
  • sustainable, as preserves raw materials by making best use of natural resources
  • more DP, and less WP, reducing waste and not needing to dispose of it if harmful
70
Q

equation for percentage uncertainty

A

error of equipment x number of readings/ amount measures x 100

71
Q

how can you reduce percentage uncertainty

A
  • take fewer readings
  • increase amount measured
  • use equipment with lower error
72
Q

what is an acid

A

-proton (H+) donor
-pH < 7
- contain hydrogen in formula, and release H+ ions in aqueous solutions

  • e.g. HCl(g) + aq ===> H+(aq) + Cl-(aq)
  • where + aq shows that there is an excess of water
73
Q

what is a strong acid

A
  • acid that fully dissociates in water, releasing all H+ ions
  • e.g. HCl(aq) ====> H+(aq) + CL-(aq)
74
Q

what is a weak acid

A
  • acid that only partially dissociates in water, only releasing a small amount of H+ ions
  • e.g. CH3COOH(aq) ===> H+(aq) + CH3OOH-(aq)
75
Q

what is a base

A

proton (H+) acceptor, which neutralises acid to form a salt

76
Q

what is an alkali

A

a base that dissolves in water releasing hydroxide ions (OH-) into the solution
-e.g. NaOH(s) + aq ===> Na+(aq) + OH-(aq)

77
Q

what is a salt

A

a product from a reaction in which H+ ions from the acid are replaced with metal or ammonium ions

78
Q

what are example of bases

A

metal oxide, metal hydroxide, metal carbonates, alkali

79
Q

types of neutralisation reactions

A

acid + base ===> salt + water
acid + carbonate ===> salt + water + CO2
acid + metal ===> salt + hydrogen
acid + ammonia ===> ammonium salt

80
Q

what is the neutralisation ionic equation

A

H+(aq) + OH-(aq) ===> H2O(l)

81
Q

what are common acids

A

-HCl = hydrochloric acid
-H2SO4 = sulfuric acid
-HNO3 = nitric acid
-CH3COOH = ethanoic acid

82
Q

what are common alkalis

A

-NaOH (sodium hydroxide)
-KOH (potassium hydroxide)
-NH3 (ammonia)

83
Q

what do titrations tell us and what is their use

A
  • used to measure the volume of one solution that reacts exactly with another solution
    1) finding the concentration of a solution
    2)identification of unknown chemicals
    3) finding the purity of a substance
84
Q

what is a standard solution

A

a solution of a known concentration

85
Q

PAG 2: making a standard solution

A

1) weigh solid accurately (to find amount, use equations)
2) dissolve solid into a beaker with less distilled water than needed to fill the flask fully
3) move to VOLUMETRIC FLASK and rinse last traces of solution from beaker into flask also
4) fill flask carefully to graduation line so BOTTOM of meniscus touchers mark (if goes over line, start again)
5)invert flask a few times to mix solution thoroughly

86
Q

PAG 2: carrying out a titration

A

1) add one solution to conical flask using pipette
2) add second solution to burette and record initial reading
3) add few drops of indicator into flask
4) run solution from burette into flask, whilst swirling flask
5) stop when indicator colour changes, so reaches end point
6) record final reading and calculate titre (volume added from burette)
7) repeat and calculate mean with concordant results

87
Q

how do you take measurements from a burette

A
  • to nearest 0.05cm^3
  • if bottom of meniscus is in between two lines, use 0.05, and if on line use 0.00
  • read from top down of burette
88
Q

what are concordant titres

A
  • titres within 0.10cm^3 of eachother
89
Q

how do you calculate mean titre

A
  • only take concordant results and calculate
90
Q

whats the difference between oxidation number and charge number

A
  • oxidation number: can be thought of as the number of e- involved in bonding to another element
  • in oxidation, sign always goes first (+2 vs 2+)
91
Q

what are the rules of assigning oxidation numbers

A
  • all elements in their natural state have a number 0
  • sum of all oxidation numbers = charge of particle
92
Q

what is the order for assigning oxidation numbers

A

1) group 1,2,3 metals = charge number
2) fluorine is usually -1 (most electronegative)
3) hydrogen in usually +1 (except metal hydrides)
4) oxygen is usually -2 (except peroxides and flouride)
5) chlorine is usually -1

93
Q

how do you name ions which have different charges using oxidation numbers

A
  • use roman numerals to indicate oxidation state (number)
94
Q

how do you name polyatomic ions using oxidation number

A
  • always end in -ate, and bracket with roman numerals indicates oxidation number of the non-oxygen element present
  • used to be use of -ite as well, but erased now with just -ate
  • the bracket is usually emitted when using the common form of -ate
95
Q

what are redox reactions

A
  • reactions containing both reduction and oxidation
96
Q

what is oxidation

A
  • loss of electrons and increase in oxidation number
97
Q

what is reduction

A
  • gain of electrons and decrease in oxidation number
98
Q

how can you tell what has been oxidised and what has been reduced, and what is the pattern in numbers

A
  • work out the oxidation number of each ion present, and see the overall change between it as a product and as a reactant
  • the TOTAL reduction change will always equal the TOTAL oxidation change (total changes is oxidation number balance)
99
Q

how would you word if something has been oxidised or reduced

A

_____ has been oxidised/reduces because its oxidation number has increased/decreased from _____ to _______

100
Q

what is disproportionation

A

when an element is both oxidised and reduced in same reaction

101
Q

what is ionic bonding

A

electrostatic attraction between positive and negative ions (transfer of electrons)

  • ions formed often have same electron configuration as nearest noble gas
102
Q

what are cations

A

positive ions

103
Q

what are anions

A

negative ions

104
Q

how are ionic compounds structured?

A

giant ionic lattice: repeating pattern of alternatively charged ions, held by electrostatic forces of attraction between ions

105
Q

explain boiling point of ionic compounds

A

-lots of energy required to overcome the strong electrostatic attraction between +/- ions
-dependent on charge density of an ion (greater charge density increases boiling point):
1) smaller ionic radius (across period)
2)containing ions of greater charge

106
Q

solubility of ionic compounds

A

-dissolve in polar solvents, like water
-polar solvent molecules break down the lattice and surround each ion in solution

107
Q

trend of solubility in ionic compounds, and why it’s hard to predict

A

depends on: attraction of solvent itself and attraction with water molecules:
1) compound with high charge ions may have attraction too large for water to break down, so not very soluble
2) BUT if higher charge, will be able to attract water molecules more too

108
Q

explain electrical conductivity in ionic compounds

A
  • changes dependant on state
    -solid: ions held in fixed position in giant ionic lattice, no mobile charge carriers so does NOT conduct
    -liquid/molten: ionic lattice breaks and ions are free to move as mobile charge carriers, so DO conduct
109
Q

what is covalent bonding

A

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

110
Q

explain covalent bonding in terms of orbitals, and the type of attraction

A
  • the overlap of atomic orbitals, each one containing one electron and therefore giving the shared pair
  • the attraction is LOCALISED in covalent bonding, and solely between the shared pair of e- and the positive nuclei of bonded atoms - called a molecule
111
Q

what does covalent bonding form

A
  • simple molecular compounds
  • giant molecular structures
112
Q

what is the usual structure of the atoms once in a covalent bond

A

same electron structure as the nearest noble gas
-USUALLY, but cannot always trust the model, e.g. with BF3

113
Q

how do you draw dot and cross for covalent bonding

A
  • electrons in middle of shared shells
  • lone pairs shown as dots in displayed formula
  • atoms beyond fluorine (which have outer shell n=3 (can hold 18 electrons due to the d-subshell allowing expansion of the octect)) can hold up to 18 electrons in outer shell, all in different arrangements
114
Q

how do you show the displayed formula of covalent bonds

A
  • use lines to show bonds and show the relative positioning of all the atoms
  • use : for a lone pair of electrons (paired electrons that are not shared)
115
Q

what is multiple covalent bonds

A

electrostatic attraction between 2/3 shared pairs of electrons and the nuclei of bonding atoms

-e.g. O=C=O and O=O
-e.g. N≡N and H-C≡N

116
Q

what is a dative covalent bond
(may be called coordinate too)

A

-covalent bond in which the shared pair of electrons has been provided by one of the bonded atoms only
- the shared pair of electrons comes from an original lone pair which has been donated
- if an ion, put brackets and charge outside like ionic diagram, e.g. NH4+ ion
- if displayed, put arrow showing where electrons are going to

117
Q

what is average bond enthalpy

A
  • measure of strength of covalent bond
118
Q

what is electron pair repulsion theory

A

1) bonding pair regions repel each other as far as possible
2) lone pairs of electrons repel more than bonding pairs, reducing the angle by 2.5 degrees each time ((as slightly closer to nuclei))

  • means that repulsion is minimised, so held in definite shape
  • bonding pair REGION = accounts for multiple bonds too, as they are treated same as a single pair
119
Q

how do you show the 3D shape of a molecule on paper

A

1) solid line= bond in plane of paper
2) solid wedge= bond coming out of plane of paper
3) dotted wedge= bond going into plane of paper

120
Q

2 bonding pairs - shape of molecule

A

-linear
-180 degrees bond angle

121
Q

3 bonding pairs- shape of molecule

A

-trigonal planar
-120 degree bond angle

122
Q

4 bonding pair - shape of molecule

A

-tetrahedral
-109.5 degrees bond angle

123
Q

5 bonding pairs- shape of molecule

A

-octahedral
-90 degree bond angle

124
Q

3 bonding pairs
1 lone pair
shape of molecule?

A

-pyramidal
-107 degree bond angle
-lone pair = top line = 2 dots

125
Q

2 bonding pairs
2 lone pairs
shape of molecule

A
  • non-linear
  • 104.5 degree bond angle
  • top and left lone pairs = 4 dots
126
Q

what is electronegativity

A

the relative ability of an atom to attract the bonded pair of electrons in a covalent bond

127
Q

what scale determines electronegativity

A

Pauling electronegativity values

128
Q

what does electronegativity depend on, and where is it highest?

A

1) nuclear charge (increases across periodic table)
2) atomic radius (decreases across periodic table)

  • highest at top left of periodic table
129
Q

most electronegative and least electronegative ions

A
  • Fluorine = 4.0 (most)
  • nitrogen, oxygen, chlorine

-group 1 metals (least)

130
Q

how do you calculate with electronegativities the type of bond present

A
  • covalent = EN difference of 0
  • polar covalent = EN difference of 0-1.8
  • ionic = EN difference of over 1.8
131
Q

what are non polar bonds

A
  • where the bonded pair of electrons is shared equally between bonded atoms
    -e.g. diatomic molecules or hydrocarbons
132
Q

what are polar bonds

A

where the bonded pair of electrons are unevenly distributed between the atoms

133
Q

what is a dipole

A
  • the separation of opposite charges in a polar bond, where one side is slightly negative and one side is slightly positive
134
Q

how do you show a dipole

A
  • arrow with line facing least to most electronegative (from δ+ to δ-)
135
Q

explain why a molecule may have polar permanent dipoles, but not be polar overall

A
  • if the dipoles cancel out
  • by acting in opposite directions
  • so overall, no dipole
  • e.g. CO2
  • RATHER THAN H2O, which would be polar as the dipoles reinforce each other (don’t oppose, contributing to a larger overall dipole)
136
Q

what are intermolecular forces

A

weak interactions between dipoles of different molecules

137
Q

difference in covalent bonds and intermolecular forces

A

covalent - identity and chemical reactions of a molecule
intermolecular - physical properties such as melting and boiling points

138
Q

what are induced dipole dipole interactions

A
  • london forces
  • intermolecular forces which occur between ANY and ALL molecules, whether polar or not, acting in between induced dipoles
139
Q

how do london dispersion forces arise

A
  • electrons are constantly moving
  • causes an uneven distribution of charge
  • resulting in an instantaneous dipole
  • dipole induces another dipole in neighbouring particle and so on
  • dipoles attract each other
140
Q

what is the nature of instantaneous dipoles

A
  • it is constantly changing
  • at any instant, will exist, but position is constantly shifting
141
Q

how can you increase the strength of london dispersion forces

A
  • increase the number of electrons
  • larger instantaneous and induced dipoles
  • stronger london dispersion forces (so stronger attractive forces between molecules)
  • more energy needed to overcome the IMF
142
Q

what are permanent dipole-dipole interactions

A
  • difference in electronegativity of bonded atoms causes a small charge difference across the bond, resulting in a permanent dipole
  • permanent dipole-dipole interactions act between the permanent dipoles in polar molecules (dotted line in between δ+ and δ-)
  • REMEMBER, they still have london forces too
143
Q

3 intermolecular forces in order of strength, most to least

A
  • hydrogen bonding
  • permanent dipole dipole interactions
  • induced dipole dipole interactions
144
Q

what are simple molecular substances

A
  • made of simple molecules
  • small units containing a definitive number of atoms with a definitive molecular formula
145
Q

explain forces present in simple molecular substances

A
  • molecules held inlattice with WEAK intermolecular forces
  • atoms in each molecule bonded with STRONG covalent bonds
146
Q

melting point of simple molecular substances

A
  • low melting or boiling point
  • as only weak intermolecular forces holding them together in lattice break
  • not the strong covalent bonds which stay
147
Q

electrical conductivity of simple molecular substances

A
  • do not conduct electricity
  • as no mobile charge carriers present to carry charge
148
Q

solubility of polar + polar (vice versa) molecular substances

A

1) new interactions form between the solute and solvent
-for IONIC, between the δ+ and δ- ends of molecule and the positive and negative ions present, surrounding them
- for POLAR, the molecule and solvent can attract each other
-for NON-POLAR, e.g. simple molecular substances, IMF form between the molecules and solvent
2) the existing IMF weaken and break if new interactions strong enough
3) compound dissolves

149
Q

solubility of polar + non polar substances

A
  • new interactions which form between solute and solvent are weak and insignificant
  • insufficient to break existing IMF (e.g. within the polar solvent)
  • solute does not dissolve
150
Q

what is hydrogen bonding

A
  • type of permanent dipole-dipole
  • between lone pair of electrons on electronegative atom
  • and hydrogen atom attached to electronegative atom
151
Q

what are the atoms hydrogen bonding can occur with

A

N
O
F
(or OH, NH2 groups)

152
Q

how do you draw hydrogen bonding

A
  • with dashed lines (straight, as shape around H is linear)
  • between lone pair of electrons on the EN atom in one molecule
  • and hydrogen on another molecule
  • always include slightly positive and negative signs
153
Q

how can you increase the hydrogen bonding

A
  • use atom with more electron lone pairs
  • so more hydrogen bonds can be made
154
Q

water’s density

A
  • ice (solid) is less solid than water and floats
  • hydrogen bonding forms 4 bonds (as 2 lone pairs on O2 and 2 Hs attached)
  • holding water molecules apart in an open lattice structure
  • so is spread out as solid as molecules are further apart (than as a liquid)
155
Q

water’s melting and boiling point

A
  • has strong hydrogen bonding as well as given London forces
  • strongest IMF, so higher amounts of energy needed to break it down
  • so high melting and boiling points

Module 2 Chemistry (1 decks)

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