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principles Flashcards

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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
Q

what strengthens ionic bonding?

A
  • smaller ions
  • higher charges
26
Q

physical properties of ionic compounds

A

*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
Q

covalent bonding definition

A

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

28
Q

dative covalent bonds

A

the shared pair of electrons in the covalent bond come from only one of the bonding atoms.

29
Q

what type of structures are ionic

A

GIANT IONIC LATTICES

30
Q

what type of structures are covalent

A

SIMPLE MOLECULAR

31
Q

giant ionic properties

A

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
Q

simple molecular properties

A

bp.mp: low due to weak IMF between molecules
solubility: poor
conductivity: poor as no ions and electrons are localised

33
Q

electronegativity definition

A

the relative tendency of an atom in a covalent bond
in a molecule to attract electrons in a covalent bond to itself.

34
Q

factors affecting electronegativity

A

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
Q

how does a polar covalent bond form

A

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
Q

are symmetrical molecules polar

A

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
Q

what structure will compounds with very different electronegativity be

A

ionic

38
Q

what are induced dipole dipole forces

A

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
Q

what affects size of induced dipole dipole interactions

A

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
Q

explain permanent dipole dipole forces

A

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
Q

where does hydrogen bonding occur

A

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
Q

explain ice

A

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
Q

explain iodine

A

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
Q

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

A

20%

45
Q

silicon exists as three isotopes, 28Si, 29Si, and 30Si. in terms of the subatomic particles, state the difference between these isotopes of silicon

A

30Si has two more neutrons than 28Si
30Si has one more neutron than 29Si

46
Q

why is silicon in the p block

A

contains outermost electron in a p orbital

47
Q

suggest why relative atomic mass on the periodic table may be different to relative atomic mass analysed using a mass spec

A

Other isotopes could be present
Some isotopes could be absent
Different abundances of isotopes

48
Q

phosphate ion formula

A

HPO4 2-

49
Q

state how many orbitals there are in a p sub shell and how the e-‘s are arranged if this sub shell is full

A

Three atomic orbitals (in each p-sub-shell); [1 mark]
Each orbital holds two electrons; [1 mark]
With opposite spins;

50
Q

why are 2s orbitals filled before 2p orbitals

A

2p orbitals have a higher energy

51
Q

why is aq MfF2 easier to use compared with molten MgF2

A
  • high melting point
  • due to strong electrostatic attractions between oppositely charged ions in all directions
52
Q

why will bond enthalpy of a specific atom change

A

As the bond enthalpy value is a mean / average taken across a range of compounds

53
Q

why can there be 3 p sub shells but only one s sub shell

A

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
Q

why is H positioned in the middle of the periodic table and not as the first element

A

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
Q

predict whether the halogens are able to conduct electricity in any state

A
  • not able to conduct electricity
  • as no delocalised electrons
56
Q

% uncertainty equation

A

uncertainty ÷ measurement
then times by 100

57
Q

dilution equation

A

original concentration x Original volume ÷
new diluted volume

58
Q

units of the ideal gas equation

A

p = Pa
v = m3
temp = K
n = moles
R = 8.31JK-1mol-1

59
Q

number of particles equation

A

moles x avogadros constant

60
Q

atom economy

A

mr of useful reactants/ mr of all reactants x 100

DO NOT USE STOICHOMETRY

61
Q

molar gas volume equation and units

A

n = v/24
n = moles
v = dm3
24 = dm3

chem (6 decks)

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