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Chemistry AS Level CIE > Chapter 3 Atoms, molecules and stoichiometry > Flashcards

Chapter 3 Atoms, molecules and stoichiometry Flashcards

1
Q

Atomic Mass Unit

A
  • The mass of a single atom is so small that it is impossible to weigh it directly
  • Atomic masses are therefore defined in terms of a standard atom which is called the unified atomic mass unit
  • This unified atomic mass is defined as one-twelfth of the mass of a carbon-12 isotope
  • The symbol for the unified atomic mass is u (often Da, Dalton, is used as well)
  • 1 u = 1.66 x 10-27 kg
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2
Q

relative atomic mass of an element x

A

=average mass of one atom of x/ (1/12) of the mass of one carbon-12 atom

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3
Q

Relative atomic mass, Ar

A
  • The relative atomic mass (Ar) of an element is the ratio of the average mass of the atoms of an element to the unified atomic mass unit
  • The relative atomic mass is determined by using the average mass of the isotopes of a particular element
  • The Ar has no units as it is a ratio and the units cancel each other out
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4
Q

Relative atomic mass, =

A

weighted average mass of atoms in a given sample of an element / unified atomic mass unit

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5
Q

Relative isotopic mass

A

The relative isotopic mass is the mass of a particular atom of an isotope compared to the value of the unified atomic mass unit

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6
Q

To calculate the average atomic mass of an element the percentage abundance is taken into account

A
  • Multiply the atomic mass by the percentage abundance for each isotope and add them all together
  • Divide by 100 to get average relative atomic mass
  • This is known as the weighted average of the masses of the isotopes
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7
Q

, Mr

A
  • The relative molecular mass (Mr) is the ratio of weighted average mass of a molecule of a molecular compound to the unified atomic mass unit
  • The Mr has no units
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8
Q

Mr can be found by adding up the

A
  • relative atomic masses of all atoms present in one molecule
  • When calculating the Mr the simplest formula for the compound is used, also known as the formula unit
  • —Eg. silicon dioxide has a giant covalent structure, however the simplest formula (the formula unit) is SiO2
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9
Q

Relative molecular mass=

A

weigted average mass of molecules in a given sample of a molecular compound / unified atomic masss unit

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10
Q

Relative formula mass, Mr

A

-The relative formula mass (Mr) is used for compounds containing ions
It has the same units and is calculated in the same way as the relative molecular mass

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11
Q

The Avogadro constant

A

(Na or L) is the number of particles equivalent to the relative atomic mass or molecular mass of a substance
The Avogadro constant applies to atoms, molecules, ions and electrons
The value of Na is 6.02 x 1023

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12
Q

mole (mol)

A

The mass of a substance with this number of particles

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13
Q

The mass of a substance containing the same number of fundamental units as there are atoms in exactly

A

12.00 g of 12C

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14
Q

One mole of any element is equal to the relative atomic mass of that element in grams

A
  • One mole of carbon, that is if you had 6.02 x 1023 atoms of carbon in your hand, would have a mass of 12 g
  • One mole of water would have a mass of (2 x 1 + 16) = 18 g
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15
Q

Ionic compounds

A

are formed from a metal and a nonmetal bonded together

-Ionic compounds are electrically neutral; the positive charges equal the negative charges

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16
Q

Charges on positive ions

A
  • All metals form positive ions
  • –There are some non-metal positive ions such as ammonium, NH4+, and hydrogen, H+
  • The metals in Group 1, Group 2 and Group 13 have a charge of 1+ and 2+ and 3+ respectively
  • The charge on the ions of the transition elements can vary which is why Roman numerals are often used to indicate their charge
  • Roman numerals are used in some compounds formed from transition elements to show the charge (or oxidation state) of metal ions
17
Q

Non-metal ions

A
  • The non-metals in group 15 to 17 have a negative charge and have the suffix ‘ide’
  • –Eg. nitride, chloride, bromide, iodide
  • Elements in group 17 gain 1 electron so have a 1- charge, eg. Br–
  • Elements in group 16 gain 2 electrons so have a 2- charge, eg. O2-
  • Elements in group 15 gain 3 electrons so have a 3- charge, eg. N3-

-There are also more complex negative ions, which are negative ions made up of more than one type of atom

18
Q

Balancing Equations: Ionic equations

A
  • In aqueous solutions ionic compounds dissociate into their ions
  • Many chemical reactions in aqueous solutions involve ionic compounds, however only some of the ions in solution take part in the reactions
  • The ions that do not take part in the reaction are called spectator ions
  • An ionic equation shows only the ions or other particles taking part in a reaction, without showing the spectator ions
19
Q

Empirical & Molecular Formulae

A
  • The molecular formula is the formula that shows the number and type of each atom in a molecule\
  • The empirical formula is the simplest whole number ratio of the elements present in one molecule or formula unit of the compound
20
Q

Empirical & Molecular Formulae: organic molecules, simple inorganic molecules and ionic compounds

A
  • Organic molecules often have different empirical and molecular formulae
  • Simple inorganic molecules however have often similar empirical and molecular formulae
  • Ionic compounds always have similar empirical and molecular formulae
21
Q

Calculating molecular formula

A
  1. calculate relative formula mass of empirical formula
  2. divide relative fromula by mass of X by relative formula mass of empirical formula
  3. multiply each number of elements by empirical number
22
Q

Empirical formula from mass

A
  1. note the mass of each element
  2. divide the masses by atomic masses
  3. divide by the lowest figure to obtain the ration
23
Q

Empirical formula from %

A
  1. note the % by mass of each atom
  2. divide the % by atomic masses
  3. Divide by the lowest figure to obtain the ration
24
Q

Water of crystallisation

A

is when some compounds can form crystals which have water as part of their structure

25
Q

hydrated compound

A

A compound that contains water of crystallisation

26
Q

The water of crystallisation is separated from the main formula by a

A

dot when writing the chemical formula of hydrated compounds

Eg. hydrated copper(II) sulfate is CuSO4∙5H2O

27
Q

anhydrous compound

A

A compound which doesn’t contain water of crystallisation

28
Q

A compound can be hydrated to different degrees

A

Eg. cobalt(II) chloride can be hydrated by six or two water molecules
CoCl2 ∙6H2O or CoCl2 ∙2H2O

29
Q

The conversion of anhydrous compounds to hydrated compounds is

A

reversible by heating the hydrated salt
Anhydrous to hydrated salt:
CuSO4 + 5H2O → CuSO4∙5H2O

Hydrated to anhydrous salt (by heating):
CuSO4∙5H2O → CuSO4 + 5H2O

30
Q

mole calculation

A

no. of moles = mass of substance in grams / molar mass (g mol^-1)

31
Q

calculating Reacting masses

A
  • the chemical equation is required
  • This equation shows the ratio of moles of all the reactants and products, also called the stoichiometry, of the equation
  • To find the mass of products formed in a reaction the following pieces of information is needed:
  • —The mass of the reactants
  • —The molar mass of the reactants
  • —The balanced equation
32
Q

why are reacting mass calculations needed

A

The masses of reactants are useful to determine how much of the reactants exactly react with each other to prevent waste

33
Q

Percentage yield

A
  • In a lot of reactions, not all reactants react to form products which can be due to several factors:
  • —Other reactions take place simultaneously
  • —The reaction does not go to completion
  • —Reactants or products are lost to the atmosphere
34
Q

percentage yield calculation

A

=actual yield/ predicted yield (theoretical yield)

x100

35
Q

Excess & limiting reagents

A

excess=of one or more of the reactants (excess reagent)

limiting reagent= is not in excess

36
Q

determining which reactant is limiting

A
  • The number of moles of the reactants should be calculated

- The ratio of the reactants shown in the equation should be taken into account

37
Q

Volumes of gases calculation

A

volume of gas (dm3) = amount of gas (mol) x 24

-At room temperature (20 degrees Celsius) and pressure (1 atm) one mole of any gas has a volume of 24.0 dm3

38
Q

concentrations of solutions

A

-The concentration of a solution is the amount of solute dissolved in a solvent to make 1 dm3 of solution

  • A concentrated solution is a solution that has a high concentration of solute
  • A dilute solution is a solution with a low concentration of solute
39
Q

concentration of solutions calculation

A

number of moles (mol) = concentration (mol dm-3) x volume (dm3)

  • When carrying out calculations involve concentrations in mol dm-3 the following points need to be considered:
  • –Change mass in grams to moles
  • –Change cm3 to dm3

Chemistry AS Level CIE (21 decks)

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