Chapter 3 Atoms, molecules and stoichiometry Flashcards
Atomic Mass Unit
- 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
relative atomic mass of an element x
=average mass of one atom of x/ (1/12) of the mass of one carbon-12 atom
Relative atomic mass, Ar
- 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
Relative atomic mass, =
weighted average mass of atoms in a given sample of an element / unified atomic mass unit
Relative isotopic mass
The relative isotopic mass is the mass of a particular atom of an isotope compared to the value of the unified atomic mass unit
To calculate the average atomic mass of an element the percentage abundance is taken into account
- 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
, Mr
- 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
Mr can be found by adding up the
- 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
Relative molecular mass=
weigted average mass of molecules in a given sample of a molecular compound / unified atomic masss unit
Relative formula mass, Mr
-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
The Avogadro constant
(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
mole (mol)
The mass of a substance with this number of particles
The mass of a substance containing the same number of fundamental units as there are atoms in exactly
12.00 g of 12C
One mole of any element is equal to the relative atomic mass of that element in grams
- 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
Ionic compounds
are formed from a metal and a nonmetal bonded together
-Ionic compounds are electrically neutral; the positive charges equal the negative charges
Charges on positive ions
- 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
Non-metal ions
- 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
Balancing Equations: Ionic equations
- 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
Empirical & Molecular Formulae
- 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
Empirical & Molecular Formulae: organic molecules, simple inorganic molecules and ionic compounds
- 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
Calculating molecular formula
- calculate relative formula mass of empirical formula
- divide relative fromula by mass of X by relative formula mass of empirical formula
- multiply each number of elements by empirical number
Empirical formula from mass
- note the mass of each element
- divide the masses by atomic masses
- divide by the lowest figure to obtain the ration
Empirical formula from %
- note the % by mass of each atom
- divide the % by atomic masses
- Divide by the lowest figure to obtain the ration
Water of crystallisation
is when some compounds can form crystals which have water as part of their structure
hydrated compound
A compound that contains water of crystallisation
The water of crystallisation is separated from the main formula by a
dot when writing the chemical formula of hydrated compounds
Eg. hydrated copper(II) sulfate is CuSO4∙5H2O
anhydrous compound
A compound which doesn’t contain water of crystallisation
A compound can be hydrated to different degrees
Eg. cobalt(II) chloride can be hydrated by six or two water molecules
CoCl2 ∙6H2O or CoCl2 ∙2H2O
The conversion of anhydrous compounds to hydrated compounds is
reversible by heating the hydrated salt
Anhydrous to hydrated salt:
CuSO4 + 5H2O → CuSO4∙5H2O
Hydrated to anhydrous salt (by heating):
CuSO4∙5H2O → CuSO4 + 5H2O
mole calculation
no. of moles = mass of substance in grams / molar mass (g mol^-1)
calculating Reacting masses
- 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
why are reacting mass calculations needed
The masses of reactants are useful to determine how much of the reactants exactly react with each other to prevent waste
Percentage yield
- 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
percentage yield calculation
=actual yield/ predicted yield (theoretical yield)
x100
Excess & limiting reagents
excess=of one or more of the reactants (excess reagent)
limiting reagent= is not in excess
determining which reactant is limiting
- The number of moles of the reactants should be calculated
- The ratio of the reactants shown in the equation should be taken into account
Volumes of gases calculation
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
concentrations of solutions
-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
concentration of solutions calculation
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
