Transcript Slide 1

CHEM110W2
Ms Janine Kasavel
[email protected]
031-2607747
Rm: 03-041
1
Quantitative Chemistry
•matter
•units
•significant figures
•atomic structure, isotopes, periodic table
•basic nomenclature (ions, molecular &inorganic compounds)
•stoichiometry and balancing equations by inspection
•moles and Avogadro’s number
•empirical and molecular formulae
•limiting reagents
2
MATTER
“matter is anything that has mass and takes up space”
Pure Substance
- ELEMENT: can’t be decomposed into simpler
substances
- COMPOUND: composed of 2/> different elements
Mixture
-HETEROGENEOUS: visibly different composition, properties
or appearance
-HOMOGENEOUS: visibly uniform composition, properties &
appearance throughout
3
Physical properties of matter
-measured without changing the identity
or composition of the substance
Chemical properties of matter
-describe the way a substance may change
or react to form other substances
4
UNITS
Système International (SI)
MASS
LENGTH
TIME
TEMPERATURE
G
M
k
d
c
m
µ
n
p
giga
mega
kilo
deci
centi
milli
micro
nano
pico
kilogram
metre
second
Kelvin
109
106
103
10-1
10-2
10-3
10-6
10-9
10-12
kg
m
s
K
SCIENTIFIC NOTATION
5
SIGNIFICANT FIGURES
1) any figure that is not zero is significant.
2) zeroes between non-zero figures are significant.
3) exact (“counting”) numbers by definition have an ¥
number of s.f., so physical constants defined to be exact
numbers do so also.
4) leading zeroes (to the left of the first non-zero figure)
are not significant.
5) trailing zeroes (to the right of the last non-zero figure)
are significant only if the number has a d.p.
6) in measurements without a d.p., the number of s.f. is
ambiguous.
6
Using Significant Figures in Calculations
• multiplication/division
Number of s.f. in final answer is the same as the LEAST of
numbers of s.f. in each of original measurements.
• addition/subtraction
Number of d.p. in final answer is the same as the LEAST of
numbers of d.p. in each of original measurements.
7
DENSITY
ρ = mass/ volume
ρ: gcm-3
mass: g
volume: cm3
8
PRACTICE EXAMPLE
A nugget of gold with a mass of 521 g is added to 50.0 mL of water. The
water level rises to a volume of 77.0 mL. What is the density of the gold?
9
ATOMIC STRUCTURE
•
•
•
•
•
Early Atomic Theory (Dalton 1803 – 1807)
Cathode Rays & Particles (Thomson, 1897)
Electron Charge & Mass (Millikan, 1909)
Nuclear Atom (Rutherford, 1910)
Modern Atomic Structure (Rutherford, 1919)
PROTONS, NEUTRONS in the nucleus surrounded by orbiting ELECTRONS
10
Charge
Proton
Electron
Neutron
Actual/
Relative
Coulombs
1.602 x 10-19
+1
1.602 x 10-19
-1
0
0
Mass
Actual/
g
1.673 x 10-24
9.109 x 10-28
1.675 x 10-24
Relative/
u
1.00727
0.00054858
1.00866
11
Ia
1
H
PERIODIC TABLE OF THE
A
ELEMENTS
1.008
4
Li
Be
6.941
9.012
11
12
Na Mg
22.99
24.31
19
20
2
ZE
II a
3
VIII a
IV b
Vb
VI b
VII b
21
22
23
24
25
┌───VIII b───┐
Ib
II b
26
29
30
27
28
IV a
Va
VI a
VII a
5
6
7
8
9
10
B
C
N
O
F
Ne
10.81
12.01
14.01
16.00
19.00
20.18
13
14
15
16
17
18
Al
Si
P
S
Cl
Ar
26.98
28.07
30.97
32.07
35.45
39.95
31
32
33
34
35
36
4.003
A: mass number = no. protons + no. neutrons
Z: atomic number = no. protons /electrons
III b
He
III a
K
Ca
Sc
Ti
V
Cr Mn Fe
Co
Ni
Cu Zn Ga Ge As
Se
Br
Kr
39.10
40.08
44.96
47.88
50.94
52.00
54.94
55.85
58.93
58.69
63.55
65.39
69.72
72.61
74.92
78.96
79.90
83.80
37
38
39
40
41
42
43
44
45
46
47
48
49
50
51
52
53
54
Rb
Sr
Y
Zr
Nb Mo Tc Ru Rh Pd Ag Cd
In
Sn
Sb
Te
I
Xe
85.47
87.62
88.91
91.22
92.91
95.94
*98.91
101.1
102.9
106.4
107.9
112.4
114.8
118.7
121.8
127.6
126.9
131.3
55
56
57
72
73
74
75
76
77
78
79
80
81
82
83
84
85
86
Cs
Ba
La
Hf Ta
W
Re Os
Ir
Pt
Au Hg
Tl
Pb
Bi
Po
At
Rn
132.9
137.3
138.9
178.5
181.0
183.8
186.2
190.2
192.2
195.1
197.0
200.6
204.4
207.2
209.0
*209.0
*210.0
*222.0
87
88
89
Fr
Ra
*223.0
*226.0
59
60
61
62
63
64
65
66
67
68
69
70
71
*
Ac
**
*227.0
58
*Lanthanides
**Actinides
Ce
Pr
Nd Pm Sm Eu Gd Tb Dy Ho
Er Tm Yb Lu
140.1
140.9
144.2
*146.9
150.4
152.0
157.3
158.9
162.5
164.9
167.3
168.9
173.0
175.0
90
91
92
93
94
95
96
97
98
99
100
101
102
103
Th
Pa
U
*232.0
*231.0
*238.0
Np Pu Am Cm Bk
Cf
Es Fm Md No
*237.1
*251.1
*252.1
*244.1
*243.1
*247.1
*247.1
*257.1
*258.1
*259.1
Lr
*260.1
12
Isotopes
Atoms of the same element with different mass numbers due to:
different numbers of neutrons
13
Average atomic mass
AAM: average atomic mass
IM: isotopic mass
14
PRACTICE EXAMPLE
Naturally occurring Mg has three isotopes:
24Mg (78.90 %) 23.9850 u
25Mg (10.00 % )24.9858 u
26Mg (11.10 %) 25.9826 u
AAM=?
15
16
IONS
• If electrons are added to or removed from a neutral
atom, an ion is formed.
• When an atom or molecule loses electrons it becomes
positively charged  CATION (E+)
11 e11 p+
Na atom
10 e11 p+
Na+ ion
17
L.Pillay 2010
When an atom or molecule gains electrons it becomes
negatively charged  ANION (E-).
17
18 e-
e-
17 p+
17 p+
Cl atom
Cl- ion
• Generally, metal atoms tend to lose electrons
(forms cations) and non-metal atoms gain electrons
(forms anions).
18
L.Pillay 2010
COMMON
CATIONS
CHARGE FORMULA NAME
+1
H+
Hydrogen ion
+
Li
Lithium ion
+
Na
Sodium ion
K+
Potassium ion
+
Rb
Rubidium ion
+
Cs
Cesium ion
+
Ag
Silver ion
NH4+
Ammonium ion
+
Cu
Copper(I) or cuprous ion
2+
+2
Mg
Magnesium ion
2+
Ca
Calcium ion
2+
Sr
Strontium ion
2+
Ba
Barium ion
2+
Zn
Zinc ion
2+
Cd
Cadmium ion
2+
Co
Cobalt(II) or cobaltous ion
2+
Cu
Copper(II) or cupric ion
2+
Fe
Iron(II) or ferrous ion
2+
Mn
Manganese(II) or manganous ion
2+
Hg
Mercury(II) or mercuric ion
2+
Hg2
Mecury(I) or mercurous ion
2+
Ni
Nickel(II) or nickelous ion
Pb2+
Lead(II) or plumbous ion
2+
Sn
Tin(II) or stannous ion
3+
+3
Al
Aluminium ion
3+
Cr
Chromium(III) or chromic ion
Fe3+
Iron(III) or ferric ion
19
COMMON
ANIONS
CHARGE FORMULA NAME
-1
H
Hydride ion
F
Fluoride ion
Cl
Chloride ion
Br
Bromide ion
I
Iodide ion
CN
Cyanide ion
OH
Hydroxide ion
CH3COO A cetate ion
ClO3
Chlorate ion
ClO4
Perchlorate ion
NO3
Nitrate ion
MnO4
Permanganate ion
2-2
O
Oxide ion
2O2
Peroxide ion
2S
Sulfide ion
2CO3
Carbonate ion
2CrO4
Chromate ion
2Cr 2O7
Dichromate ion
2+
SO4
Sulphate ion
3-3
N
Nitride ion
3PO4
Phosphate ion
20
Ionic compounds
• Composed of nonmetal and metal
• Cations and anions attract each other to form a neutral compound
NAMES:
• Name of metal (cation) written first
• If metal has more than one common charge , write the charge in
roman numerals in brackets
• Name of nonmetal (anion) written next with –ide ending
FORMULAE:
• compounds are electrically neutral, the formula of a compound can
easily be constructed simply by:
-writing value of cation charge as subscript on anion
-writing value of anion charge as subscript on cation
21
PRACTICE EXAMPLE
NaCl
K2SO4
Ba(OH)2
cobalt(II) nitrate
silver sulfide
ferric chloride
22
Oxyanion
ClO4ClO3ClO2ClO-
perchlorate ion (one more O atom than chlorate)
chlorate ion (one more O atom than chlorite)
chlorite ion (one more O atom than hypochlorite)
hypochlorite ion
Acids
- acids containing anions whose names end in -ide are
named by changing the -ide ending to -ic, adding the
prefix hydro- to this anion name, and then following with
the word acid
- acids containing anions whose names end in -ate/-ite are
named by changing the -ate ending to -ic or the -ite ending
to -ous and then adding the word acid
23
PRACTICE EXAMPLE
Anion
Corresponding acid
ClS2ClO4ClO3ClO2ClO-
24
MOLECULAR COMPOUNDS
• Generally composed only of nonmetals
• Diatomic species includes O2 N2, F2, Br2, I2
NAMING:
• name of element furthest left on periodic table generally written first
• both elements in same group on periodic table, element with higher
Z written first
• name of 2nd element given the ending –ide
• Greek prefixes used to indicate number of atoms of each element
Greek prefixes:
mono, di, tri, tetra, penta, hexa, hepta, octa, nona, deca
1 2 3 4
5
6
7
8
9
10
25
PRACTICE EXAMPLE
SO2
PCl5
N2O3
NF3
P4S10
silicon tetrabromide
26
STOICHIOMETRY
“quantities of substances consumed and produced in chemical reactions”
• Atoms are neither created or destroyed in a chemical
reaction.
• A chemical equation must have equal numbers of atoms of
each element on each side of the arrow.
• The molecular composition of certain ions must remain the
same on each side of the arrow.
27
PRACTICE EXAMPLE
C2H6 + O2 → CO2 + H2O
Al + HCl → AlCl3 + H2
28
MOLE & AVOGADRO’S NUMBER
Number of atoms/molecules/ions represented as mole amounts
Avogadro’s number: NA = 6.022 X 1023
1 mol 12C atoms = 6.022 X 1023 12C atoms
1 mol H2O molecules = 6.022 X 1023 H2O molecules
1 mol NO3 - ions = 6.022 X 1023 NO3- ions
29
Molar mass
“Mass in grams of one mole of a substance”
• Related to mole amount of a substance by the equation:
m
n
MM
n: number of moles (in mol)
m: mass (in grams)
MM: molar mass (in grams per mole)
30
PRACTICE EXAMPLE
How many oxygen atoms are in 1.50 mol of sodium carbonate?
31
EMPIRICAL AND MOLECULAR
FORMULA
“Ratio of atoms of each element in a compound”
Mass %
elements
Assume
100g
sample
Grams of
each
element
Use molar
mass
Moles of
each
element
Empirical
formula
Calculate
mole ratio
32
PRACTICE EXAMPLE
Determine the empirical formula of a compound with 10.4% C, 27.8% S and 61.8% Cl.
33
QUIZ 2
• Wed 17 August during the tut
• Stoichiometry (slide 28) until end quantitative
chemistry
TUTS
• Check chem foyer for new tut group lists after
1pm on Monday 15/08
34
PRACTICE EXAMPLE
Eucalyptol has an empirical formula of C10H18O. The experimentally determined
molecular mass of this substance is 152 u. What is its molecular formula?
35
Ia
1
H
VIII a
PERIODIC TABLE OF THE
ELEMENTS
II a
1.008
2
III a
IV a
Va
VI a
VII a
He
4.003
3
4
5
6
7
8
9
10
Li
Be
B
C
N
O
F
Ne
6.941
9.012
10.81
12.01
14.01
16.00
19.00
20.18
11
12
13
14
15
16
17
18
Al
Si
P
S
Cl
Ar
26.98
28.07
30.97
32.07
35.45
39.95
31
32
33
34
35
36
Na Mg
22.99
24.31
19
20
III b
IV b
Vb
VI b
VII b
21
22
23
24
25
┌───VIII b───┐
Ib
II b
26
29
30
27
28
K
Ca
Sc
Ti
V
Cr Mn Fe
Co
Ni
Cu Zn Ga Ge As
Se
Br
Kr
39.10
40.08
44.96
47.88
50.94
52.00
54.94
55.85
58.93
58.69
63.55
65.39
69.72
72.61
74.92
78.96
79.90
83.80
37
38
39
40
41
42
43
44
45
46
47
48
49
50
51
52
53
54
Rb
Sr
Y
Zr
Nb Mo Tc Ru Rh Pd Ag Cd
In
Sn
Sb
Te
I
Xe
85.47
87.62
88.91
91.22
92.91
95.94
*98.91
101.1
102.9
106.4
107.9
112.4
114.8
118.7
121.8
127.6
126.9
131.3
55
56
57
72
73
74
75
76
77
78
79
80
81
82
83
84
85
86
Cs
Ba
La
Hf Ta
W
Re Os
Ir
Pt
Au Hg
Tl
Pb
Bi
Po
At
Rn
132.9
137.3
138.9
178.5
181.0
183.8
186.2
190.2
192.2
195.1
197.0
200.6
204.4
207.2
209.0
*209.0
*210.0
*222.0
87
88
89
Fr
Ra
*223.0
*226.0
59
60
61
62
63
64
65
66
67
68
69
70
71
*
Ac
**
*227.0
58
*Lanthanides
**Actinides
Ce
Pr
Nd Pm Sm Eu Gd Tb Dy Ho
Er Tm Yb Lu
140.1
140.9
144.2
*146.9
150.4
152.0
157.3
158.9
162.5
164.9
167.3
168.9
173.0
175.0
90
91
92
93
94
95
96
97
98
99
100
101
102
103
Th
Pa
U
*232.0
*231.0
*238.0
Np Pu Am Cm Bk
Cf
Es Fm Md No
*237.1
*251.1
*252.1
*244.1
*243.1
*247.1
*247.1
*257.1
*258.1
*259.1
Lr
*260.1
36
LIMITING REACTANTS
• In chemical reactions one reactant (the
LIMITING reactant (LR)) is used up first
• The reaction stops once LR used up leaving
the excess reactants as leftovers
• The LR limits the amount of product formed
37
1. Balance equation
2. Determine the molar masses
3. Calculate the number of moles for each
reactant and divide by stoichiometric
coefficients
4. Smallest mole amount =LR
38
PRACTICE EXAMPLE
Determine (i) which reactant is the limiting reactant and (ii) how much excess
reactant is leftover in the following reaction:
3NH4NO3 + Na3PO4 → (NH4)3PO4 + 3NaNO3
30.0 g
50.0 g
39
Theoretical yield
• Quantity of product that is calculated to form, based on LR
• Actual yield is the mass of product obtained when the
experiment is carried out
• Percent yield relates actual yield to theoretical yield
% Yield = (actual yield/theoretical yield) x 100%
40
PRACTICE EXAMPLE
Determine the theoretical yield and the % yield of NaNO3 if 15.0 g of sodium nitrate
is formed when the reaction is carried out.
3NH4NO3 + Na3PO4 → (NH4)3PO4 + 3NaNO3
30.0 g
50.0 g
41
QUANTITATIVE CHEMISTRY REVISION
CHAPTER 1 and 2
42