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1
Q

atomic number definition

A

number of protons in the nucleus

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2
Q

mass number definition

A

total number of protons and neutrons in the atom

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3
Q

isotope

A

atoms of same element with the same number of protons, but different numbers of neutrons.

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4
Q

do isotopes have the same physical and chemical properties?

A

Isotopes have similar chemical properties because they have the same electronic
structure.
They may have slightly varying physical properties because they have
different masses.

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5
Q

relative isotopic mass definition

A

the mass of one isotope compared to one twelfth of the mass of one atom of carbon-12

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6
Q

relative atomic mass

A

the weighted mean mass of one atom compared to one twelfth of the mass of one atom of carbon-12

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7
Q

relative molecular mass

A

the average mass of a molecule
compared to one twelfth of the mass of one atom of carbon-12

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8
Q

common strong acids

A

Hydrochloric ( HCl), sulfuric (H2SO4) and nitric (HNO3) acid;
Ethanoic acid CH3COOH is a weak acid

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9
Q

most common alkalis

A

The most common alkalis are sodium hydroxide (NaOH),
potassium hydroxide (KOH) and aqueous ammonia (NH3)

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10
Q

alkali definition

A

soluble base that releases OH- ions in aqueous solutions

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11
Q

salt definition

A

formed when the H+
ion of an acid is replaced by
a metal ion or an ammonium ion

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12
Q

acid + base

A

salt and water

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13
Q

acid + carbonate

A

salt + water + co2

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14
Q

what are the observations in carbonate reactions

A
  • effervescence due to CO2 gas and solid carbonate will dissolve
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15
Q

titration method

A

*rinse equipment (burette with acid, pipette with alkali, conical
flask with distilled water) *pipette 25 cm3 of alkali into conical flask
*touch surface of alkali with pipette ( to ensure correct amount
is added)
*adds acid solution from burette
*make sure the jet space in the burette is filled with acid
*add a few drops of indicator and refer to colour change at end
point
*phenolphthalein [pink (alkali) to colourless (acid): end point pink
colour just disappears] [use if NaOH is used]
*methyl orange [yellow (alkali) to red (acid): end point orange]
[use if HCl is used]
*use a white tile underneath the flask to help observe the colour change
*add acid to alkali whilst swirling the mixture and add acid dropwise at end point *note burette reading before and after addition of acid
*repeats titration until at least 2 concordant results are
obtained- two readings within 0.1 of each other

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16
Q

how to decrease apparatus uncertainties

A

either decrease
the sensitivity uncertainty by using apparatus with a greater resolution (finer scale divisions) or you can increase the size of the measurement made.

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17
Q

Uncertainty of a measurement using a burette.

A

f the burette used in the
titration had an uncertainty for each reading of +/– 0.05 cm3
then during a titration two readings would be taken
so the uncertainty on the titre volume would be +/– 0.10 cm3.

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18
Q

how to reduce titration uncertainties

A

Replacing measuring cylinders with pipettes or burettes which have lower apparatus uncertainty will lower the % uncertainty.

To reduce the % uncertainty in a burette reading it is necessary to make the titre a larger volume. This could be done by: increasing the volume
and concentration of the substance in the conical flask or by decreasing the concentration of the substance in the burette

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19
Q

oxidation is….

A
  • the process of electron loss
  • increase in oxidation number
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20
Q

reduction is…

A
  • process of electron gain
  • decrease in oxidation number
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21
Q

oxidation number of O2 in H2O2

A

-1

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22
Q

oxidation number of hydrogen in metal hydrides

A

-1

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23
Q

shape of s and p sublevels

A
  • s are spherical
  • p are dumbbells
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24
Q

ionic bonding definition

A

the electrostatic force of attraction between oppositely charged ions formed by electron transfer.

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25
what strengthens ionic bonding?
- smaller ions - higher charges
26
physical properties of ionic compounds
*High melting points - There are strong electrostatic attractive forces between the oppositely charged ions in the lattice *Non conductor of electricity when solid- The ions are held together tightly in the lattice and can not move so no charge is conducted *Good conductor of electricity when in solution or molten – The ions are free to move when in solution and molten. Charge can be carried * They are usually soluble in aqueous solvents.
27
covalent bonding definition
the strong electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atoms
28
dative covalent bonds
the shared pair of electrons in the covalent bond come from only one of the bonding atoms.
29
what type of structures are ionic
GIANT IONIC LATTICES
30
what type of structures are covalent
SIMPLE MOLECULAR
31
giant ionic properties
bp/mp: high due to giant lattice of ions with strong electrostatic forces between oppositely charged ions. solubility: good conductivity when solid: poor conductivity when molten: good as ions can move
32
simple molecular properties
bp.mp: low due to weak IMF between molecules solubility: poor conductivity: poor as no ions and electrons are localised
33
electronegativity definition
the relative tendency of an atom in a covalent bond in a molecule to attract electrons in a covalent bond to itself.
34
factors affecting electronegativity
Electronegativity increases across a period as the number of protons increases and the atomic radius decreases because the electrons in the same shell are pulled in more. It decreases down a group because the distance between the nucleus and the outer electrons increases and the shielding of inner shell electrons increases A compound containing elements of similar electronegativity and hence a small electronegativity difference will be purely covalent
35
how does a polar covalent bond form
when the elements in the bond have different electronegativities. When a bond is a polar covalent bond it has an unequal distribution of electrons in the bond and produces a charge separation, (dipole)
36
are symmetrical molecules polar
A symmetric molecule (all bonds identical and no lone pairs) will not be polar even if individual bonds within the molecular are polar. Symmetric molecules The individual dipoles on the bonds ‘cancel out’ due to the symmetrical shape of the molecule. There is no net dipole moment: the molecule is non-polar.
37
what structure will compounds with very different electronegativity be
ionic
38
what are induced dipole dipole forces
Induced dipole–dipole interactions are also called London forces. They occur between all simple covalent molecules and the separate atoms in noble gases. In any molecule the electrons are moving constantly and randomly. As this happens the electron density can fluctuate and parts of the molecule become more or less negative i.e. small temporary or transient dipoles form. These temporary dipoles can cause dipoles to form in neighbouring molecules. These are called induced dipoles. The induced dipole is always the opposite sign to the original one. THEY DO NOT OCCUR IN IONIC SUBSTANCES
39
what affects size of induced dipole dipole interactions
The more electrons there are in the molecule the higher the chance that temporary dipoles will form. This makes the induced dipole–dipole interactions stronger between the molecules and so boiling points will be greater.
40
explain permanent dipole dipole forces
occurs between polar molecules *It is stronger than induced dipole–dipole interactions and so the compounds have higher boiling points *Polar molecules have a permanent dipole. (commonly compounds with C-Cl, C-F, C-Br H-Cl, C=O bonds) *Polar molecules are asymmetrical and have a bond where there is a significant difference in electronegativity between the atoms.
41
where does hydrogen bonding occur
in compounds that have a hydrogen atom attached to one of the three most electronegative atoms of nitrogen, oxygen and fluorine, which must have an available lone pair of electrons
42
explain ice
Water can form two hydrogen bonds per molecule, because oxygen is very electronegative, and it has two lone pairs of electrons The molecules are held further apart than in liquid water and this explains the lower density of ice
43
explain iodine
There are covalent bonds between the iodine atoms in the I2 molecule. The crystals contain a regular arrangement of I2 molecules held together by weak induced dipole– dipole interactions intermolecular forces.
44
a sample of neon contains the isotopes 20Ne and 21Ne. the relative atomic mass is 20.2. what is the percentage of 21Ne atoms in the sample
20%
45
silicon exists as three isotopes, 28Si, 29Si, and 30Si. in terms of the subatomic particles, state the difference between these isotopes of silicon
30Si has two more neutrons than 28Si 30Si has one more neutron than 29Si
46
why is silicon in the p block
contains outermost electron in a p orbital
47
suggest why relative atomic mass on the periodic table may be different to relative atomic mass analysed using a mass spec
Other isotopes could be present Some isotopes could be absent Different abundances of isotopes
48
phosphate ion formula
HPO4 2-
49
state how many orbitals there are in a p sub shell and how the e-'s are arranged if this sub shell is full
Three atomic orbitals (in each p-sub-shell); [1 mark] Each orbital holds two electrons; [1 mark] With opposite spins;
50
why are 2s orbitals filled before 2p orbitals
2p orbitals have a higher energy
51
why is aq MfF2 easier to use compared with molten MgF2
- high melting point - due to strong electrostatic attractions between oppositely charged ions in all directions
52
why will bond enthalpy of a specific atom change
As the bond enthalpy value is a mean / average taken across a range of compounds
53
why can there be 3 p sub shells but only one s sub shell
S-orbitals are spherical so multiple sub-shells are not possible P-orbitals are propeller-shaped and can be positioned along the x, y or z axis/ in three different directions Three p-orbitals would not overlap significantly
54
why is H positioned in the middle of the periodic table and not as the first element
Hydrogen has very different physical properties to Group 1 metals Hydrogen has very different chemical properties to Group 1 metals Despite having an outer shell configuration ending in s1
55
predict whether the halogens are able to conduct electricity in any state
- not able to conduct electricity - as no delocalised electrons
56
% uncertainty equation
uncertainty ÷ measurement then times by 100
57
dilution equation
original concentration x Original volume ÷ new diluted volume
58
units of the ideal gas equation
p = Pa v = m3 temp = K n = moles R = 8.31JK-1mol-1
59
number of particles equation
moles x avogadros constant
60
atom economy
mr of useful reactants/ mr of all reactants x 100 DO NOT USE STOICHOMETRY
61
molar gas volume equation and units
n = v/24 n = moles v = dm3 24 = dm3

chem (6 decks)

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