2. Atoms, molecules and stoichiometry Flashcards
define the unified atomic mass unit
one twelfth of the mass of a carbon-12 atom
define relative atomic mass, Ar, relative isotopic mass, relative molecular mass, Mr, and relative formula
mass in terms of the unified atomic mass unit
Ar-
Mr-
isotopic mass-
relative formula mass-
define and use the term mole in terms of the Avogadro constant
common ions
NO3–, CO32–, SO42–, OH–, NH4+, Zn2+, Ag+,HCO3–, PO4 3–
define and use the terms empirical and molecular formula
empirical-
molecular-
understand and use the terms anhydrous, hydrated and water of crystallisation
perform calculations including use of the mole concept, involving:
(a) reacting masses (from formulas and equations) including percentage yield calculations
(b) volumes of gases (e.g. in the burning of hydrocarbons)
(c) volumes and concentrations of solutions
(d) limiting reagent and excess reagent
moles= mass/ Ar/Mr
moles= volume*conc
moles= volume/24
% yield= actual yield/theoretical yield
% composition=mass/total mass*100
limiting-excess= ratio of moles and lower moles is limiting.
define electronegativity
power of an atom to attract electrons to itself
explain the factors influencing the electronegativities of the elements
nuclear charge-
atomic radius-
shielding by inner shells and sub-shells-
explain the trends in electronegativity across a period and down a group
use the differences in Pauling electronegativity values to predict the formation of ionic and covalent bonds
< 1- covalent
1-2 - polar covalent
> 2-ionic
ionic bonding with examples
electrostatic attraction between oppositely charged ions (positively charged
cations and negatively charged anions)
sodium chloride, magnesium oxide and calcium fluoride
define metallic bonding
define covalent bonding
electrostatic attraction between positive metal ions and delocalised electrons.
electrostatic attraction between the nuclei of two atoms and a shared pair of
electrons.
describe covalent bonding in molecules including:
* hydrogen, H2
* oxygen, O2
* nitrogen, N2
* chlorine, Cl 2
* hydrogen chloride, HCl
* carbon dioxide, CO2
* ammonia, NH3
* methane, CH4
* ethane, C2H6
* ethene, C2H4
- hydrogen, H2- 1 pair
- oxygen, O2- 2 pairs
- nitrogen, N2- 3 pairs
- chlorine, Cl 2- 1 pair
- hydrogen chloride, HCl- 1 p
- carbon dioxide, CO2- 2-2 p
- ammonia, NH3- 1,1,1
- methane, CH4- 1,1,1,1
- ethane, C2H6 1 bw all
- ethene, C2H4 2 bw C, 1 bw C-H
period 3 property of octet
period 3 can expand their octet including in the compounds sulfur dioxide, SO2, phosphorus pentachloride, PCl 5 , and sulfur hexafluoride, SF6
describe coordinate (dative covalent) bonding
reaction between ammonia and hydrogen chloride gases to form the ammonium ion, NH4+, and in the Al 2Cl 6 molecule
describe covalent bonds in terms of orbital overlap giving σ and π bonds:
- σ bonds are formed by direct overlap of orbitals between the bonding atoms
- π bonds are formed by the sideways overlap of adjacent p orbitals above and below the σ bond
use the concept of hybridisation to describe sp, sp² and sp³ orbitals
describe how the σ and π bonds form in molecules including H₂, C₂H₆, C₂H₄, HCN and N₂
define bond energy and bond length.
use bond energy values and the concept of bond length to compare the reactivity of covalent molecules
the energy required to break one mole of a particular covalent bond in the gaseous state.
the internuclear distance of two covalently bonded atoms
state and explain the shapes of, and bond angles in, molecules by using VSEPR theory, including as
simple examples:
- BF3 (trigonal planar, 120°)
- CO2 (linear, 180°)
- CH4 (tetrahedral, 109.5°)
- NH3 (pyramidal, 107°)
- H2O (non-linear, 104.5°)
- SF6 (octahedral, 90°)
- PF5 (trigonal bipyramidal, 120° and 90°)
trigonal planar- 3B 0L
linear- 2B, 0L
tetrahedral- 4B, 0L
pyramidal- 3B, 1L
non-linear- 2B, 2L
trigonal bipyramidal- 5B, 0L
octahedral- 6B, 0L
describe hydrogen bonding
molecules containing N–H and O–H groups.
ammonia and water- examples
use the concept of hydrogen bonding to explain the anomalous properties of H₂O
- its relatively high melting and boiling points
- its relatively high surface tension
- the density of the solid ice compared with the liquid water
use the concept of electronegativity to explain bond polarity and dipole moments of molecules
describe van der Waals’ forces- generic term to describe all intermolecular forces.
describe the types of van der Waals’ forces:
the intermolecular forces between molecular entities other than those due to bond formation.
- instantaneous dipole–induced dipole (id-id) forces, also called London dispersion forces
- permanent dipole–permanent dipole (pd-pd) forces, including hydrogen bonding
describe hydrogen bonding and state order of strength of bonds
hydrogen bonding is a special case of permanent dipole–permanent dipole forces between molecules where hydrogen is bonded to a
highly electronegative atom. FONCl.
in general, ionic, covalent and metallic bonding are stronger than intermolecular forces.
