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foundations in chemistry Flashcards

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1
Q

what is relative atomic mass (Ar)?

A

Ar is the weighted mean mass of an atom relative to 1/12th of the mass of an atom of carbon-12

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2
Q

what does Ar take into account?

A

the percentage abundance of each isotope and the relative isotopic mass of each isotope.

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3
Q

what is relative isotopic mass?

A

relative isotopic mass is the mass of one isotope relative to 1/12th of the mass of one atom of carbon-12

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4
Q

how would you work out the percentage abundance of an isotope to work out relative atomic mass?

A

using a mass spectrometer

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5
Q

how much roughly is 1u (atomic mass unit)

A

mass of one proton

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6
Q

what is the relative mass of one electron

A

1/1836

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7
Q

what does the principal quantum number, n refer to?

A

the shell number/energy level number

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8
Q

what is the maximum number of electrons which can be held in shell:
1
2
3
4

A

n=1–>2 electrons
n=2–>8 electrons
n=3–>18 electrons
n=4–> 32 electrons

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9
Q

electron shells themselves are split into sub-shells. what are these labelled?

A

s, p, d and f

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10
Q

which sub shells are in each electron shell?

A

n=1 shell has the s sub shell
n=2 shell has the s and p sub shell
n=3 shell has the s p and d sub shell
n=4 shell has the s p d and f sub shell

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11
Q

what is the maximum number of electrons each sub shell can hold?

A

s sub shell can hold 2 electrons max.
p sub shell can hold 6 electrons max.
d. sub shell can hold 10 electrons max.
f. sub shell can hold 14 electrons max

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12
Q

the s, p, d, and f sub shells are themselves divided into further atomic orbitals. How many orbitals does each sub shell contain?

A

s sub shell- 1 s-orbital
p sub shell- 3 p-orbitals
d sub shell- 5 d-orbitals
f sub shell- 7 f-orbitals

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13
Q

how is the direction of the electron spin represented in an atomic orbital?

A

a box with arrows. the box represents the atomic orbital. arrows for the electrons. clockwise spin- upwards arrow
anticlockwise spin- downwards arrow

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14
Q

what is an atomic orbital?

A

a region of space in an atom that can hold up to 2 electrons with opposite spins

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15
Q

what are the 4 types of orbitals?

A

s, p, d and f

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16
Q

what is the rule for metals in term of bonding?

A

metals always lose electrons to form positive ions/cations

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17
Q

what is a polyatomic ion?

A

An ion that contains more than one element bonded together

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18
Q

list the 8 polyatomic ions we need to know and their charges

A

ammonium- NH4 +
Hydroxide-OH -
nitrate- NO3 -
carbonate-CO3 2-
sulfate- SO4 2-
zinc- zn 2+
silver- Ag+
phosphate- PO4 3-

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19
Q

what is amount of substance?

A

what we use to measure the number of particles

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20
Q

what is 1 mole?

A

amount of substance that contains 6.02x10^23 particles

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21
Q

what is relative formula mass/molar mass and its units?

A

sum of the atomic masses, g/mol -1

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22
Q

how do you work out no. of molecules?

A

1) work out moles of entire compound
2) multiply moles by avogadros constant

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23
Q

how do you work out no. of atoms of one element in a compound? (3)

A

1) work out moles of entire compound
2) if the required element has no little number next to it then it remains the same no. of moles/ if it has a number e.g a 3 then multiply the mol. by the number
3) multiply by Avogadro’s constant

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24
Q

how do you work out total no. of atoms in a compound?

A

1) work out total no. of moles for each individual element by working out mol of total compound then comparing it to the small number next to the element
2) add these together
3) multiply by avogadros constant

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25
Q

how do you work out empirical formula?

A

1) For each element, divide the mass/percentage given in the question by the Ar
2) choose the smallest number
3) divide each answer from part 1 by the smallest number
4) this gives you the ratio the elements come in
5) if you get a number with .5 ,multiply the whole compound by 2

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26
Q

what is relative molecular mass (Mr) ?

A

Ar units added up

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27
Q

how do you work out molecular formula when given empirical formula and Mr?

A

1) add up the Ar values of the given empirical formula
2) divide the Mr by this value
3) multiply each element in the empirical formula by the number you get from part 2

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28
Q

how do orbitals fill? (2)

A

1) filled in a specific order to produce the lowest energy arrangement possible
2) Orbitals with the same energy will fill with one electron each first and when all orbitals are occupied then they will pair up

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29
Q

in ionic bonding, how are electrons transferred?

A

from metal atoms to non metal atoms

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30
Q

how are ions formed? (2)

A

1) when metals and non metals react

2) the overall energy change for the reaction is favourable

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31
Q

what happens to the metal and non metal atoms after ionic bonding?

A

both form a stable electronic structure like that of the nearest noble gas

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32
Q

why are ionic compounds typically solid at room temperature and pressure?

A

there is insufficient energy to overcome the strong electrostatic forces of attraction between the oppositely charged ions in the giant ionic lattice

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33
Q

what are the structures of ionic compounds?

A

giant ionic lattice consisting of regular repeating positive and negative ions in all three dimensions

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34
Q

what is ionic bonding?

A

electrostatic attraction between
positive and negative ions

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35
Q

why do ionic compounds have high melting points?

A

high temperatures needed to provide the large quantity of energy required to overcome the strong electrostatic attraction between ions

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36
Q

what does ionic attraction depend on?

A

1) size of ions
2) ionic charge (greater charge, higher melting point due to stronger attraction between ions)

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37
Q

how do ionic compounds dissolve in polar solvents such as water? (2)

A

1) ionic lattice is broken down
2) water molecules attract and surround the ions

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38
Q

why are compounds with larger charges less likely to dissolve?

A

ionic attraction may be too strong for water to be able to break down the lattice structure

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39
Q

what does the solubility of an ionic compound depend on?

A

1) relative strengths of attraction within the giant ionic lattice
(solubility decreases as charge increases)
2) attractions between ions and water molecules

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40
Q

describe ionic compounds in their solid state (3)

A

1) ions in fixed position in lattice
2) no mobile charge carriers
3) non-conductor of electricity

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41
Q

describe ionic compounds in their liquid/molten state (3)

A

1) lattice broken down
2) ions free to move as mobile charge carriers
3) conductor of electricity

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42
Q

describe the structure of metals (3)

A

1) giant lattice structure
2) metal cations in a sea of delocalised valence electrons
3) electrons free to move, so account for the flow of electricity in metals

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43
Q

describe how the sea of delocalised valence electrons and metal cations interact and how this affects their physical properties (2)

A

1) whole structure is held together by the electrostatic attractions between them
2) causes relatively high melting and boiling points

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44
Q

what is covalent bonding?

A

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

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45
Q

describe the attraction in covalent bonds

A

attraction is localised, solely between the shared pair of electrons and the nuclei of bonded atoms

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46
Q

how do you approach “water of crystallisation questions”? (4)

A

1) use mass= mr x mol equation
2) find mol of water lost/ water produced
3) find mol of salt
4) divide mol of water by mol of crystal

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47
Q

What is the “water of crystallisation”?

A

Water within the structure of crystals

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48
Q

what two things are assumed when working out the water of crystallisation?

A

1) the salt does not decompose when heated (lose mass)
2) all water has been lost

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49
Q

how do you convert from cm3 to dm3 to m3

A

cm3 to dm3 divide by 1000
dm3 to m3 divide by 1000

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50
Q

what is the rule for gases at RTP?

A

1 mole of gas has a volume of 24dm3 at RTP

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51
Q

what are you indirectly measuring when calculating volume of gas?

A

the number of gas molecules

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52
Q

when do we use the ideal gas equation?

A

when an experiment has been carried out at a temperature or pressure different to RTP

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53
Q

what is an ‘ideal gas?’ (4)

A

1) random motion
2) elastic collisions
3) negligible size
4) no intermolecular forces

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54
Q

what is the ideal gas equation? (including units)

A

pV= nRT
p= pressure (Pa)
V= volume (m3)
n= moles (mol)
R= gas constant (8.314)
T= temperature (K, add 273 to degrees C)

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55
Q

what is a dative covalent bond?

A

when both electrons come from the same atom

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56
Q

how are dative covalent bonds represented in displayed formulae?

A

shown by an arrow pointing away from the atom that donates the pair of electrons

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57
Q

what are the physical properties of covalently bonded simple molecules? (4)

A

1) relatively small amounts of energy are needed to separate one molecule from another
2) relatively low melting and boiling points
3) do not conduct electricity (no charged particles)
4) do not dissolve readily in water

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58
Q

what is average bond enthalpy?

A

a measurement of covalent bond strength

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59
Q

what shape and bond angles will be in a molecule/ion with 2 bonding pairs of electrons?

A

shape: linear
bond angles: 180

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60
Q

what shape and bond angles will be in a molecule/ion with 3 bonding pairs of electrons?

A

shape: trigonal planar
bond angles: 120

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61
Q

what shape and bond angles will be in a molecule/ion with 4 bonding pairs of electrons? give an example

A

shape: tetrahedral
bond angles: 109.5
example: CH4

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62
Q

what shape and bond angles will be in a molecule/ion with 5 bonding pairs of electrons?

A

shape: triganol bipyramid
bond angles: 90 and 120

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63
Q

what shape and bond angles will be in a molecule/ion with 6 bonding pairs of electrons?

A

shape: octahedral
bond angles: 90

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64
Q

what shape and bond angles will be in a molecule/ion with 3 bonding pairs of electrons and 1 lone pair of electrons?
give an example

A

shape: pyramidal
bond angles: 107
example: NH3

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65
Q

what shape and bond angles will be in a molecule/ion with 2 bonding pairs of electrons and 2 lone pairs of electrons?
give an example and explain the reason for the bond angle

A

shape: non linear (V-shaped/bent)
bond angles: 104.5 (lone pairs of electron on the oxygen cause a further repulsion of 2.5 from the pyramidal shape)
example: H2O

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66
Q

what is the relative repulsive strength of bonded pairs compared to lone pairs of electrons?

A

lone pairs repel more than bond pairs as a result of being closer to the nucleus of the atom

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67
Q

how does electron pair repulsion explain the shapes of molecules and ions? (3)

A

1) in simple covalent molecules, two types of electrons exist, bonding and lone pairs of electrons
2) when a molecule forms the outer shell electrons, the bonding and lone pairs repel each other
3) the shape the molecule takes is that which minimises the repulsion between different pairs of electrons

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68
Q

what should you do when doing a reacting masses question?

A

always show working for mass and mol

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69
Q

why may we not achieve a yield of 100%? (3)

A

1) side reactions may have taken place
2) purification may have lead to the loss of some product
3) reaction may have not gone to completion

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70
Q

how would you approach a percentage yield question where you’re given two masses? (4)

A

1) write out balanced symbol equation to find out mole ratio
2) look at the masses, DONT use the one which is the actual yield
3) create a table of the elements which the masses are related to and treat it like a reacting masses question
4) use the mass you get from step 3 and use percentage yield equation by dividing the actual yield (given in Q) by the answer and multiplying by 100

71
Q

how do you approach a limiting reagents question which asks which is limiting reactant and how much moles the product has?

A

1) calculate number of moles of each substance if needed
2) write out symbol equation for mol ratio if needed
3) work out limiting reactant by diving mol by mol ratio number (smallest one is limiting reagent)
4) the number of moles in the products will be the same as the mol of limiting reactant

72
Q

how do you approach a limiting reactants question where it is asking for the mass of a product from two reactants? (3)

A

1) work out moles of each reactant
2) smallest one is the moles of product formed
3) multiply mol by mr of product to find mass

73
Q

which groups and periods in the periodic table can expand the octet rule? (2)

A

1) groups 5, 6, 7, and 0
2) periods beyond 3

74
Q

what happens if a molecule contains lone pairs in terms of shapes of molecules? (2)

A

1) since lone pairs are more repulsive than bonding pairs, they distort the basic shape we expect of the molecule
2) reduces bond angle between bonding pairs by 2.5 degrees

75
Q

what is electronegativity?

A

the ability of an atom to attract the bonding electrons in a covalent bond

76
Q

how is electronegativity measured?

A

The Pauling scale is used to compare the electronegativity of the atoms of different elements

77
Q

how can you measure electronegativity on the periodic table? (2)

A

1) electronegativity increases towards Flourine in the periodic table (most electronegative element is fluorine with a Pauling value of 4)
2) electronegativity increases up and across the periodic table

78
Q

what is electronegativity like in group 1?

A

least electronegative atoms

79
Q

which elements are the most electronegative?

A

non metals (nitrogen, oxygen, fluorine) and chlorine

80
Q

what determines a non polar bond/pure covalent bond? (2)

A

1) bonded atoms are the same
2) bonded atoms have the same or similar elctronegativity

81
Q

how can ideas about electronegativity be used to predict chemical bond type? (3)

A

1) if the electronegativity difference is large, one bonded atom will have a much greater attraction for the electron pair than the other bonded atoms
2) the more electronegative atom will have gained control of the electrons
3) therefore the bond will now be ionic rather than covalent

82
Q

what happens in a non-polar/pure covalent bond?
what molecules are examples of this?

A

the bonded electron pair is shared equally between the bonded atoms
e.g H2, O2, N2 and Cl2

83
Q

what happens in polar-bonds?

A

the bonded electron pair is shared unequally between the bonded atoms

84
Q

what determines a polar bond? what does this result in? (2)

A

1) when the bonded atoms are different
2) when the bonded atoms have different electronegativity values (one atom will have a greater attraction for the bonded pair of electrons which will cause a dipole)

85
Q

what is a dipole?

A

the separation of opposite charges

86
Q

what is a permanent dipole? where are they found? (2)

A

1) a dipole in a polar covalent bond which does not change (called a permanent dipole to distinguish it from an induced dipole)
2) found within molecules containing covalently bonded atoms with different electronegativities

87
Q

for molecules with more than 2 atoms there may be 2 or more polar bonds, what does this mean in terms of dipoles?

A

depending on the shape of the molecule:
1) Dipoles may reinforce one another to produce a larger dipole over the whole molecule
2) dipoles may cancel out if the dipoles act in opposite directions

88
Q

what is a polar molecule? give an example

A

A​ polar molecule​ requires ​polar bonds​ with dipoles that do not cancel due to their direction
E.g. H2O and CO2 both have polar bonds but only H2O has an overall dipole so H2O is polar but CO2 is not

89
Q

describe the solubility of polar and non polar molecules (3)

A

1) polar solutes tend to dissolve best in polar solvents
2) non-polar solutes tend to dissolve best in non-polar solvents.
3) a polar solute in a non-polar solvent (or vice versa), tend to be insoluble or only soluble to a miniscule degree.

90
Q

what are the 3 main categories of intermolecular forces?

A

1) induced dipole-dipole interactions (London forces/ dispersion forces)
2) permanent dipole-dipole interactions
3) hydrogen bonding (special type of permanent dipole - permanent dipole forces)

91
Q

what are intermolecular forces?

A

the weak interactions between different molecules

92
Q

what are induced dipole-dipole interactions/ London forces/ dispersion forces?

A

weak intermolecular forces that exist between all molecules as a result of the presence of electrons in the molecule and the formation of temporary dipoles

93
Q

what are permanent dipole-dipole interactions?

A

the attractive forces between two neighbouring molecules with a permanent dipole

94
Q

how do induced dipole-dipole interactions occur? (5)

A

1) movement of electrons produces a changing dipole in a molecule
2) an instantaneous dipole will exist at any instant but its position is constantly shifting
3) this instantaneous dipole will induce a dipole to neighbouring molecules
4) this induced dipole induces further dipoles on neighbouring molecules
5) these molecules attract each other with London forces

95
Q

what is the duration of induced dipoles?

A

only temporary, in the next instant of time the induced dipoles may disappear as the movement of electrons keeps changing

96
Q

what is the strength of induced dipoles?

A

weak and easily broken

97
Q

why does every atom/molecule experience London forces?

A

because London forces are caused by random electron movement even if they experience other intermolecular forces as well

98
Q

what does the strength of London forces depend on?

A

the number of electrons

99
Q

The more electrons in each molecule…(in London forces) (3)

A

1) the larger the instantaneous and induced dipoles
2) the greater the London forces
3) The stronger the attractive forces between molecules

100
Q

how does the number of electrons in a molecule affect its boiling point? (3)

A

1) atoms with a greater number of electrons have a higher boiling point
2) this is because they experience stronger London forces
3) more energy is needed to overcome these forces, increasing boiling point

101
Q

what does van Der Waals’ forces apply to?

A

both permanent and induced dipole-dipole interactions

102
Q

what are the properties of simple molecular substances? explain (2)

A

1) low melting and boiling points
2) in a simple molecular lattice, the weak intermolecular forces can be broken by the energy present at low temperatures

103
Q

what do simple molecules form in their solid state?
how does this happen?
(2)

A

1) a simple molecular lattice which has covalently bonded molecules that are attracted by intermolecular forces
2) all simple molecular substances can be solidified into simple molecular lattices by reducing the temperature

104
Q

why does hydrogen chloride have a higher boiling point than flourine? (3)

A

1) hydrogen chloride molecules are polar and have both London forces and permanent dipole-dipole forces interactions between molecules
2) fluorine molecules are non polar and only have London forces acting between molecules
3) extra energy is needed to break the additional permanent dipole-dipole interactions between the hydrogen chloride molecules

105
Q

what is a simple molecular substance? (2)
give examples

A

1) substance made up of simple molecules covalently bonded together, containing a definite number of atoms with a definite molecular formula
2) at RT all simple molecular substances may exist as solids, liquids or gases
e.g
Ne, H2, H2O, CO2

106
Q

what occurs in simple molecular lattices? (2)

A

1) the molecules are held in place by weak intermolecular forces
2) the atoms within each molecule are bonded together strongly by covalent bonds

107
Q

what happens when a simple molecular lattice is broken apart during melting? (2)

A

1) only the weak intermolecular forces break
2) the covalent bonds are strong and do not break

108
Q

which two categories do covalent substances with simple molecular structures fall into, in terms of solubility? (2)

A

polar and non polar (non polar is easier to predict solubility of)

109
Q

what is the solubility like in non polar simple molecular substances? (4)

A

1) tend to be soluble in non polar solvents
2) when a non polar simple molecular compound is added to a non polar solvent such as hexane, intermolecular forces form between the molecules and the solvent
3) these interactions weaken the intermolecular forces in the simple molecular lattice
4) intermolecular forces break and the compound dissolves

110
Q

how do non polar simple molecular substances react in polar solvents? why? (3)

A

1) simple molecular substances tend to be insoluble in polar solvents
2) there is little interaction between the molecules in the simple molecular lattice and the polar solvent molecules
3) the intermolecular bonding within the polar solvent is too strong to be broken

111
Q

why is the solubility of polar simple molecular substances hard to predict? give an example

A

the solubility depends on the strength of the dipole
e.g
hydrogen chloride gas is extremely soluble in water due its polar H-Cl bond, forming hydrochloric acid whereas sugar is also soluble in water due to many O-H bonds but not as much.

112
Q

what is the solubility like in polar simple molecular substances? (3)

A

1) polar simple molecular substances may dissolve in polar solvents
2) the polar solute molecules and the polar solvent molecules can attract each other
3) similar process to dissolving an ionic compound

113
Q

give an example of a polar simple molecular substance dissolving in a polar solvent (3)

A

1) sugar dissolves in water (polar solvent)
2) sugar is a polar covalent compound with many polar O-H bonds
3) these O-H bonds attract and bond with polar water molecules

114
Q

which compounds can dissolve in both polar and non-polar solvents?

A

compounds such as ethanol which contain both polar (O-H) and non polar (carbon chain) parts in their structure

115
Q

some biological molecules have hydrophobic and hydrophilic parts, what does this mean? (2)

A

1) hydrophilic part will be polar and contain electronegative atoms (usually oxygen) that can interact with water
2) hydrophobic parts will be non polar and be compromised of a carbon chain

116
Q

what are simple molecular structures like in terms of electrical conductivity? (3)

A

1) no mobile charged particles in simple molecular structures
2) there is nothing to complete an electrical circuit
3) simple molecular substances are non-conductors of electricity

117
Q

what is a hydrogen bond?

A

an attraction between a lone pair of electrons on an electronegative atom in one molecule and a hydrogen atom in another molecule attached to an electronegative atom

118
Q

how do you draw a hydrogen bond?

A

the bond must run from the hydrogen atom directly to the lone pairs of electrons

119
Q

how does hydrogen bonding affect the properties of water? (5)

A

1) gives water some anomalous properties that support the existence of life on earth
2) hydrogen bonds hold water molecules apart in an open lattice structure
3) the water molecules in ice are further apart than in water
4) solid ice is less dense than liquid water and so floats on water.
5) water is one of the few substances in which the solid is less dense than water

120
Q

why does water have relatively high melting and boiling points? (3)

A

1) water has both London forces and Hydrogen bonds acting between molecules
2) large quantity of energy needed to break hydrogen bonds
3) water has much higher melting and boiling points than would be expected from just London forces

121
Q

what happens when ice melts?

A

the rigid arrangement of hydrogen bonds between ice break.

122
Q

what happens when water boils?

A

hydrogen bonds break completely

123
Q

what is the equation for atom economy?

A

AE= mr of desired product divided by the sum of the Mrs of the reactants multiplied by 100

124
Q

what is atom economy?

A

a measure of how well atoms of reactants have been utilised

125
Q

what is an acid?

A

acids release H+ ions in aqueous solution

126
Q

what are some examples of acids? (5)

A

HCl
HNO3
H2SO4
H3PO4
CH3COOH

127
Q

what is a strong acid? give an example

A

acid that fully dissociates in water
e.g
HCl –> H+ + Cl-

128
Q

what is a weak acid? give an example

A

acid that partially dissociates in water
e.g
CH3COOH –> CHCOO- + H+
<–

129
Q

what is a base? (2)

A

1) something that neutralises an acid to form a salt
2) readily accepts H+ ions from an acid

130
Q

what are some examples of bases? (4)

A

metal oxides
metal hydroxides
metal carbonates
ammonia

131
Q

what is an alkali?

A

a base that dissolves in water and release OH– ions in aqueous solution

132
Q

what is a neutralisation? give an example (3)

A

1) an acid and a base reacting together to form a salt
2) the reaction of H+ and OH– to form H2O
3) in the salt, the H+ ions of the acid are replaced by a metal or ammonium ion
e.g
HCl –> NaCl (sodium/metal replaces hydrochloric acid)

133
Q

what are 3 examples of neutralisation reactions?

A

1) acid + metal oxide –> salt + water
2) acid + metal hydroxide –> salt + water
3) acid + metal carbonate –> salt + water + carbon dioxide

134
Q

what are spectator ions?

A

ions that are present during the reaction but are unchanged by the reaction

135
Q

how do you write ionic equations? (4)

A

1) make sure symbol equation has state symbols
2) write out individual ions involved in reaction (aqueous state symbol elements)
3) cancel out spectator ions
4) left over ions/ elements are what you use in ionic equation

136
Q

what is a titration?

A

a technique used to accurately measure the volume of one solution that reacts with another solution

137
Q

what are the 3 main uses of a titration?

A

1) find the concentration of a solution
2) identifying an unknown chemical
3) finding the purity of a substance

138
Q

what is a standard solution?

A

a solution of a known concentration

139
Q

what techniques and procedures would you use to prepare a standard solution? (10)

A

1) measure out the mass of the solid on a weighing boat
2) place in a beaker
3) swill out weighing boat residue into beaker
4) dissolve in deionised water and stir until fully dissolved
5) use a funnel to pour the beaker contents into a volumetric flask
6) swill out any residue in beaker into volumetric flask using deionised water
7) swill residue on funnel with deionised water
8) fill volumetric flask with deionised water until you get close to the line
9) use a pipette to fill until the meniscus is on the bottom of the line
10) put lid on and swirl around/ tip upside down

140
Q

what is a volumetric flask?

A

a flask which is manufactured to measure a certain volume within a certain tolerance

141
Q

what techniques and procedures would you use to carry out an acid-base titration? (8)

A

1) pour acid/alkali into burette with a beaker and fill to 0.00cm3
2) open tap and pour out some acid/alkali into the same beaker get rid of any air stored in the tap
3) measure a known volume (usually 20 or 25 cm3) from the standard solution with a volumetric pipette and place it into a conical flask
4) place a white tile underneath
5) place a few drops of indicator into the conical flask
6) take a rough titration
7) take accurate titrations
8) keep going until you obtain 2 concordant results (rough titration not included)

142
Q

how do you take a rough titration? (2)

A

1) the tap of the burette is opened to allow the solution inside to flow into a known volume of the solution in the conical flask
2) The amount of solution from the burette required to reach the end point is recorded.

143
Q

what are concordant results?

A

results within 0.1cm3 of each other which exclude rough titrations

144
Q

how do you take accurate titrations? (4)

A

1) after rough titration is taken, refill burette to 0.00cm3, refill conical flask with standard solution etc
2) run the burette portion by portion until about 5cm3 within the rough titration
3) as you near the end point, add slowly drop by drop and continue swirling until end point is reached
4) if needed use deionised water to swill drops on the end of the burette

145
Q

what is the set pattern for analysing results?

A

1) work out mol for solution you have conc. + vol. for
2) use mol ratio in balanced symbol equation to work out mol of other solution
3) work out the unknown information about the solute in the other solution (this can involve multiple steps)

146
Q

how do you approach an ‘identifying carbonates’ question? (6)

A

1) find mol of substance that has known conc. + vol.
2) find mol of other substance using mol ratio
3) to find mol of original substance find out the multiplier between substance in original container vs the container its been put into
4) multiply mol in step 2 by multiplier in step 3
5) find the Mr of the unknown carbonate by dividing the mr by the mol
6) minus known Mrs
7) identify carbonate using Mr

147
Q

what are the rules for assigning and calculating oxidation number for atoms in elements?

A

oxidation number is always zero for elements

148
Q

what are the rules for assigning and calculating oxidation number compounds and ions? (3)

A

1) each atom in a compound has an oxidation number
2) an oxidation number has a sign which is placed before the number
3) the oxidation number of an ion of an element is numerically the same as the ionic charge

149
Q

describe redox in terms of oxidation number (2)

A

1) reduction is a decrease in oxidation number
2) oxidation is an increase in oxidation number.

150
Q

which shapes of molecules are symmetrical? (5)

A

linear, trigonal planar, tetrahedral, octahedral and trigonal ​bipyramidal

151
Q

what are the solid structures of simple molecular lattices? give an example

A

covalently bonded molecules attracted by intermolecular forces (e.g I2)

152
Q

what is the effect of structure and bonding on the physical properties of covalent compounds with simple molecular lattice structures? (3)

A

1) low melting and boiling points due to weak intermolecular forces
2) solubility of these compounds are determined if they’re polar or non polar
3) non conductors of electricity due to there being no charged particles present

153
Q

what are the anomalous properties of H2O as a result of hydrogen bonding?

A

1) relatively high melting and boiling points due to hydrogen bonds as well as London forces which require a large quantity of energy to break
2) ice is less dense than water due to hydrogen bonding. In ice water molecules are held further apart by hydrogen bonds in an open lattice structure.

154
Q

what is hydrogen bonding?

A

intermolecular bonding between lone pair of electrons on molecules containing N, O or F and the H atom of -NH, -OH or -HF.

155
Q

what are the oxidation numbers we need to know? (5)

A

1) O is -2
2) H is +1
3) group 7 elements are -1
4) group 1 elements are +1
5) group 2 elements are +2

156
Q

what are the exceptions when it comes to oxidation number?

A

1) H in metal hydrides is -1
2) O in peroxides is -1
3) O bonded to F is +2

157
Q

how do you work out oxidation number? (3)

A

1) sum of oxidation numbers = total charge
2) work out oxidation numbers of elements you know
3) write an equation which adds up to total charge

158
Q

in redox reactions, the total change in oxidation number is…

A

balanced

159
Q

what are oxidation numbers used for? (3)

A

1) to tell if oxidation or reduction has taken place
2) to work out what has been oxidised/ reduced
3) to construct half equations and balance redox equations

160
Q

what are the oxidation rules? (6)

A

1) the oxidation number of an uncombined element is zero
2) many atoms or ions have fixed oxidation numbers in compounds (look back to oxidation numbers for O, H etc)
3) oxidation number of an element in a monoatomic ion is always the same as charge
4) the sum of the oxidation number in a compound is zero
5) the sum of oxidation number in ions is equal to the charge on the ion
6) in a compound or ion, the more electronegative element is given the oxidation state

161
Q

what are oxidation numbers like for metals? (3)

A

1) have positive values in compounds
2) oxidation number is usually the same as group number
3) oxidation number goes no higher than group number e.g Sn can be +2 or +4 and Mn can be +2, +4, +6, or +7

162
Q

what are oxidation numbers like for non metals? (2)

A

1) mostly negative based on their usual ion
2) can have values up to their group number e.g Cl can be +1, +3, +5 or +7

163
Q

what are Roman numerals used for in terms of oxidation number?

A

used to show the oxidation states of transition metals which can have more than one oxidation state

164
Q

what reaction is classified as a redox reaction? why? (2)

A

1) Metal + acid → salt + hydrogen
2) during this reaction, there are changes in oxidation number

165
Q

how do you explain if a reaction is redox? (3)

A

1) write balanced symbol equation
2) deduce changes in oxidation number
3) explain which species is reduced/ oxidised

166
Q

what is an oxidising agent?

A

something which gets reduced and accepts electrons

167
Q

what is a reducing agent?

A

something which gets oxidised and donates electrons

168
Q

how can you identify the oxidising agent and reducing agent in a reaction? (3)

A

1) Deduce the oxidation numbers on both sides of the reaction
2) Identify which species has been oxidised and which has been reduced by looking at the oxidation numbers
3) Identify the oxidising and reducing agent

169
Q

how do you write half equations? (4)

A

1) deduce what has been oxidised and reduced
2) for the oxidation half equation, write the electron after the arrow
3) for the reduction half equation, write the electron before the arrow
4) when writing charge, only use the symbol and represent the number as the big number before the element

170
Q

what is a mole?

A

amount of substance that has the same number of particles as there are atoms in 12g of 12C

171
Q

what happens when a metal reacts with water?

A

a metal hydroxide and hydrogen are formed

172
Q

why do isotopes of an element still have the same chemical properties?

A

because they still have the same electron configuration

173
Q

how do you work out moles from the ideal gas equation?

A

n= pV/RT

chemistry a level (6 decks)

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