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2014/07/23
Atomic Structure
CHM 172 Chapter 2 part 1
Atoms, Molecules, Ions
End of Chapter Questions:
7, 9, 17, 19, 22,27, 43, 49 51, 60
• Electronic Charge: e– = − 1.602 x 10−19 C Use
multiples of this charge, i.e. 1 e– = − 1
• p+ = + 1.602 x 10−19 C
1 p+ = + 1
• Chemical properties of elements and
molecules depend largely on the electrons of
atoms involved.
• 1 u = 1/12th of mass of C – 12 atom = 1.66054 x 10−24 g
• Can use symbols for atoms:
represents the element.
Q1: What is the mass number of an iron atom with
30 neutrons? Give the symbol.
Q2. What is the composition of an atom of
phosphorus with 16 neutrons? Give mass number
and symbol.
• If the atom has an actual mass of 30.9738 u, what
is its mass in grams?
• What is the mass of this atom relative to carbon–
12?
Isotopes …..
Rel Atomic mass = (% abundance Isotope 1)(mass Isotope1)
100
+ (% abundance Isotope 2)(mass Isotope 2)
100
Q3. Cℓ: 35Cℓ mass = 34.96885 u, abundance =
75.77%
37Cℓ mass = 36.96590 u, abundance = 24.23%
Relative Atomic mass?
A
ZX,
where X
Isotope abundance
% abundance = number of atoms of given isotope x 100
Total number of atoms of all isotopes
• e.g. Silver (Ag) has 2 isotopes – one with 60 no (%
abundance = 51.839%) and the other 62 no.
• What are their mass numbers?
• What is the abundance of the isotope with 62 no.
Q4: Silicon consists of 3 stable isotopes, 28Si
(92.23%) with mass 27.97693 u; 29Si (4.68%) with
mass 28.97649 u: and 30Si (3.09%) with mass
29.97377 u. Calculate the relative atomic mass of
Si.
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Fractional Abundance
Fractional abundance =
number of atoms of given isotope =
Total number of atoms of all isotopes
% abundance
100
eg. Antimony (Sb) has 2 stable isotopes, 121Sb, mass
120.904 u; and 123Sb, mass 122.904u. What are
their relative abundances?
• From the periodic table – get weighted relative
atomic mass, i.e. 121.760 u
• Rel. atomic mass = [(fractional abundance,
isotope 1)(mass isotope 1)] + [(fractional
abundance, isotope 2)(mass isotope 2)
Fractional abundances involving 3 isotopes:
eg. Q6. The element silicon has 3 naturally
occurring isotopes, 28Si (27.976927 u); 29Si
(28.976495 u) and 30Si (29.973771 u).
• The collective fractional abundance of 28Si and
29Si is 0.9690.
• Determine the % abundance of all 3 isotopes.
Mass Defect
• Isotopic masses are generally less than the sum
of masses of the subatomic particles composing
that atom.
• This “missing mass” is = the energy (i.e binding
energy) that holds the nuclear particles
(nucleons) together = difference between the
mass of the nucleus and the sum of the proton
and neutron masses (∆m).
• Can be expressed using Einstein’s equation:
E = mc2
• Binding Energy:
Eb = (∆m)c2
• Nuclear stabilities of different elements can be
compared using binding energy per mole of
nucleons.
Q5: Gallium (Ga) consists of 2 naturally occurring
isotopes with masses of 68.9257u and 70.9249u.
How many protons and neutrons are in the nucleus
of each isotope? Write complete atomic symbols
for each isotope.
Calculate the % abundance of each isotope. [69Ga =
60.12%; 71Ga = 39.88%]
Q7 : Neon consists of 3 isotopes: 20Ne
(19.992435 u); 21Ne (20.993843 u); 22Ne
(21.991383). The percentage abundance of 21Ne
is 0.27%, determine the percentage abundance
of the other 2 isotopes.
Molecules, Compounds, Formulae
• COMPOUNDS are a combination of 2 or
more elements in definite ratios by mass.
• MOLECULES are the smallest unit of a
compound that retains the characteristics of the
compound.
• Molecular formula summarises the composition
of the substance – does not show how atoms
come together to form the molecule.
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MOLECULAR FORMULAS
• Formula for glycine is C2H5NO2
• Can also write glycine
formula as
MOLECULAR FORMULAS
• The physical and chemical properties of a
molecular compound are often closely related to
its structure.
•
e.g. O2 – essential for life, odourless
•
O3 – toxic, sharp, pungent smell
• Structural formula shows how atoms are
connected within molecule.
e.g. H2O
H2O2
CH4
to show to atom ordering
• Can also write glycine formula in the form of a
structural formula
• Molecular formula summarises the composition
of the substance – does not show how atoms
come together to form the molecule.
IONS AND IONIC COMPOUNDS
NAMES of IONS
• Consist of ions (i.e. atoms or groups of atoms that
bear a positive or negative electric charge).
• CaIons: atom loses one/more electron(s) →
more protons than electrons →
• Anions: atom gains one/more electron(s) → more
electrons than protons →
• Monatomic ions: single atoms that gain/lose
electrons.
• Metals generally lose electrons during reactions
• Non-metals usually gain electrons during
reactions
• Cations: metal + ‘cation’, e.g. Mg2+ =
• Transition metals: more than one cation, e.g.
Fe2+ = iron(II) cation; Fe3+ = iron (III) cation.
• - Ammonium NH4+
• Anions
• Monatomic – add -ide to stem of the name
of the non-metal.
• S2- sulphide N3- nitride O2-, oxide
POLYATOMIC IONS
NAMES of POLYATOMIC IONS
• Oxoanions - with greatest number of oxygen atoms → suffix
• Polyatomic ions are made up of 2 or more
atoms with a collective charge:
• e.g. Nitrate, NO3- : 31 p+, 32 e- for the group of
atoms.
•
Ammonium, NH4+: 11 p+, 10 e• See Table 2.4 Formulae and names of
common polyatomic ions.
• Also …centre pages of Study Guide
-ate
– smaller number of oxygen atoms → -ite
NO3- (nitrate);
NO2- (nitrite)
2SO4 (sulphate); SO32- (sulphite)
• More than 2 members: largest number of oxygen –
prefix persuffix -ate
• Lowest number of oxygen atoms – prefix hyposuffix
-ite
e.g. CℓO4- (perchlorate);
CℓO3- (Chlorate);
CℓO2 (chlorite);
CℓO- (hypochlorite)
If it contains hydrogen: hydrogen ‘oxoanion’
e.g.
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Formulae of Ionic Compounds
Names of Ionic Compounds
• Electrically neutral, i.e. no net electric charge
→ charges balanced by the number of posiIve
and negative ions.
e.g. aluminium oxide – consists of Aℓ3+ and O2ions
• lowest mulIple of the charges is 6 → i.e. 2
Aℓ3+ and 3 O2- → Aℓ2O3
• By convention: symbol of cation given first,
then anion.
Ca2+
+
CℓCa2+
+
CO322+
Ca
+
PO43-
– built from names of positive and negative ions:
Molecular Compounds: Formulae
and Names
• Combinations of 2 non-metals = binary
compounds.
• Hydrogen forms binary compounds with all
non-metals (except noble gases)
e.g. HF hydrogen fluoride
H2S hydrogen sulphide
• Formula usually written with elements in
order of increasing group number: H2O;
HF
– exceptions CH4; NH3
Names and Formulae of Acids
• Acids = hydrogen containing compounds
• → yield hydrogen ions (H+) in water
→ H written first
• Considered as an anion connected to enough H+
- ions to neutralise / balance the anion charge.
→ Name related to the name of the anion.
• Anions whose names end in –ide become –ic
acid with the prefix hydro - …..
e.g. Cℓ- (chloride) →
CN- (cyanide) →
• e.g.
MgCO3
magnesium carbonate
Mg2+ and CO32•
Fe2(SO4)3 iron(III) sulphate
Fe3+ and SO42• Lithium nitrate
• Sodium perchlorate
Q8: Arrange the following compounds in order
of increasing lattice energy: NaF, CsI, CaO
Molecular Compounds: Formulae
and Names
• Number of atoms of a given type are
designated by the prefix, di-, tri-, tetrae.g. NF3:
nitrogen trifluoride
N2O4:
dinitrogen tetraoxide
PCℓ5:
phosphorus pentachloride
Review and Check Ex. 2.7 and Ex 2.8
Q9. Name the following:
N2F4; (NH4)2S; V2O3
Names and Formulae of Acids
• Anions whose names end in –ate or –ite
become –ic or –ous acid
e.g. CℓO4- (perchlorate) → HCℓO4
CℓO2- (chlorite) → HCℓO2
SO42- (sulphate) →
H2SO4
2SO3 (sulphite) → H2SO3
Q10. Name the following:
HNO3 ; H2CO3 ; HBr
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