A. Definitions 1. Chemistry 2. Matter CHEMISTRY I. Introduction.
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Transcript of What is chemistry? - Maine-Endwell Central School District … · · 2014-08-254 Chemistry •the...
1
What is
chemistry?
Chemistry
2
What is Chemistry? CuCl2 + NH3 What happens?
CuCl2 soluble
Cu(OH)2(s) insoluble
Cu(NH3)4+2
soluble
3
Molecular
Interpretation
Cu+2 +2OH- Cu(OH)2
This course: understand & predict
Cu(OH)2 + 4NH3
Cu(NH3)4+2 + 2OH-
NH3 + H2O NH4+ + OH-
4
Chemistry
•the study of matter
•and the changes
matter undergoes
Focus:
•Underlying principles
•Molecular point-of-view
5
Classification of Matter
?? H2, cement, air, NaCl
Homo-geneous
Hetero-geneous
Mixture
Cmpd Element
Substance
Matter
“solution”
Physically
Separate
Chem. Separate.
Give examples of physical separation & chemical separation
Physical vs. Chemical Change
6
Physical change: no new substance is formed. e.g. melt, bend, cut, boil
Chemical change: new substance formed (always involves a chemical reaction).
e.g. Fe(s) + S(s) FeS(s)
7
Units of Measurement
Quantity SI Unit Others
Length m cm
Mass kg g
Time s
Temperature K oC
Metric Prefixes (know them!)
8
giga- G E9 mega- M E6 kilo- k E3 deci- d E-1 centi- c E-2 milli- m E-3 nano- n E-9 pico- p E-12
How many grams is 14 pg?
9
Derived Units
How many cm3 in 1 m3 ?
For example: volume
10
Derived Units
Density = mass/volume
What is the mass of 17.4 mL
of ethanol (D = 0.798 g/mL)?
1 g H2O @ 4oC = 1 cm3 = 1mL
11
Temperature Scales
Kelvin
0
273
373
(Size of
degree is
the same)
Celsius
Motion
stops -273
Water freezes
Water boils 100
0
12
Temperature Scales
K = oC + 273.15
Try it:
1. 72oC = ? K
2. What is room temperature in oC?
3. If T changes by 37oC, what is
the change in Kelvin?
13
Scientific Notation Avogadro’s Number:
602,200,000,000,000,000,000,000
6.022 x 1023 particles
Mass of H atom:
0.00000000000000000000000166 g
1.66 x 10-24 g
14
Using Scientific Notation
N. x 10n
Add/subtract: n must be the same
Add 7.4 x 103 and 2.1 x 102
Multiply (divide): add (subtract) n’s
Divide 8.0 x 102 by 6.2 x 104
15
Significant Figures
•All non-zeros are sig. 1.234 ? •Zeros between non-zeros are sig. 40,501 ?
How wide is this room?
16
Significant Figures
Zeros to left are NOT sig. 0.000349 ?
Zeros to right are sig. if there is a decimal point. 0.09030 ? 400 ?
17
Sig Figs in Calculations
Add/subtract: answer has same number of digits to right of decimal as the number with the fewest digits to right of decimal.
89.321 + 9.2 =
18
Sig Figs in Calculations
Multiplication/division: answer has the same number of sig figs as the number with least number of sig figs.
2.8 x 4.5039 =
19
Pure Numbers
How many sig figs in:
Five people?
2.54 cm = 1 in
20
Conversion Factors
5 apples = $1.00
5 apples $1.00
= 1
In Wegman’s, apples cost 5 for a dollar.
21
Conversion Factor
5 apples = 1 dollar
1 dollar = 100 pennies
10-2 m = 1 cm (or 1 m = 100 cm )
1 mole = 6.02 x 1023 atoms
1 g H2O = 1 mL H2O
22
Conversion Factors: Try it.
Convert radius of Ca atom (0.197 nm) to cm. (Hint: go through the base unit.)
Adult has 5.2 L of blood. What is volume in m3 ?
What is 65 miles per hour (mph) in m/s ? (1609 m = 1 mi)
23
Atoms, Molecules, Ions
Early chemists study the first known atom.
24
Atomic Theory
Democritus vs. Dalton
500 BC 1808 AD
Pure
thought
Observation
+ Reason
Matter is not continuous!! Weird.
25
Dalton’s Atomic Theory
Elements composed of atoms
Atoms of element are identical*
Compounds: simple integers of
atoms
Chem Rxn: rearrangement, not
creation, of atoms
26
Dalton’s Theory
Law of Definite Proportions:
(mass of C and O in CO2)
Law of Multiple Proportions:
(CO and CO2)
Based on observations:
27
Subatomic Particles
The Electron
28
The Electron
Symbol: e-
Charge = -1 = -1.6 x 10-19C
Mass = 9.09 x 10-28 g
29
The Proton
Ernest
Rutherford
1910
30
The Proton
Charge = +1
Mass = 1.67 x 10–24 g
31
The Neutron
Mass ratio of He/H problem
Chadwick 1932
(nuclear chemistry)
Charge = 0
Mass = 1.67 x 10–24 g
(similar to proton)
32
Relative Proportions
Protons &
Neutrons
Electron cloud
33
Atomic Number & Mass Number
Mass number
(#p + #n)
Atomic number
(#p)
X A
Z
H 1
1 H
2
1 H
3
1
34
Practice
How many p, n, e- in:
Hg ? 200
80
-2 ion of Oxygen-16 ?
Hg-200 or
35
Periodic Table
36
Periodic Table
Group/Family vs. Period
Metals/Nonmetals/Metalloids
Key Groups
1 or 1A: Alkali Metals
2 or 2A: Alkaline Earth Metals
17 or 7A: Halogens
18 or 8A: Noble Gases
37
Atoms, Molecules, Ions
Water Ammonia Methane
38
Systematic Names: Why?
Common Name Formula
Water H2O
Lime CaO
Lye NaOH
Potash K2CO3
Laughing Gas N2O
Baking Soda NaHCO3
Muriatic Acid HCl
39
Chemical Formulas
Molecular formula: Exact
numbers of atoms of each
element in a molecule.
Cl
H
C Chloroform
CHCl3
40
Chemical Formulas Empirical: smallest whole number
ratio of atoms in a compound.
Glucose: C6H12O6 CH2O
Hydrogen Peroxide: H2O2 HO
41
Molecular Compounds
Molecular cmpds: bonded nonmetals
Many molecular compounds
are binary (only 2 elements)
ending in ---ide.
HCl = hydrogen chloride
Organic compounds have a special naming system.
42
Molecular Compounds
Name is:
prefix-atom-prefix-atom-ide mono 1
di 2
tri 3
tetra 4
penta 5
hexa 6
N2O dinitrogen monoxide
PCl3 phosphorus trichloride
43
Molecular Compounds
SiCl4
AlCl3
H2O NH3 H2S
Try it!
44
Ionic Compounds
Ionic Compounds: usually
a positive metal ion +
negative nonmetal ion.
Binary: NaCl = sodium chloride
Ternary: BaCO3 = barium carbonate
45
Ions Ion: Charged atom or group of
atoms (monatomic or
polyatomic; Cl- vs CO3-2)
chlorine atom: 17 p & 17 e-
chloride ion: 17 p & 18 e-
Cation: +positive Anion: -negative
46
Ionic Formulas
Sum of charges is zero.
e.g. Aluminum Oxide
Al2O3
“criss-cross”
2(+3) + 3(-2) = 0
Al3+ O2-
47
Ionic Formulas Ionic compounds always use
empirical formulas.
Sodium Chloride: NaCl
Polyatomic Ions (know them!)
48
Memorize the ‘ate’ Ions
49
ate: C2H3O2
- CO32- C2O4
2-
ClO3- CrO4
2- Cr2O72-
MnO4- NO3
- PO43-
SCN- SO42- S2O3
2-
Rule: if the non-oxygen atom is in an odd Group, charge is -1 (except PO4
3-)
Figure Out the ‘hydrogen ---ate’ ions
50
CO32-
HCO3-
SO42- HSO4
-
Adding a hydrogen decreases the charge by 1.
Figure Out ‘ite’ ions
51
ClO4- perchlorate 1 more O
ClO3- know the ‘ate’
ClO2- chlorite 1 less O
ClO- hypochlorite 2 less O
Try it: name IO2-
Memorize Some Others
52
H3O+ Hg2
2+ NH4+
CN- O22- OH-
Have fun!
53
Ionic Compounds
Some metals form
more than one cation:
Fe+2 Fe+3
Classical Ferrous Ferric
Stock Iron(II) Iron(III)
FeCl3 = Iron(III) chloride
Common Metals with more than one charge (know them!)
54
Cu+ copper(I) Sn2+ tin(II)
Cu2+ copper(II) Sn4+ tin(IV)
Fe2+ iron(II) Hg22+ mercury(I)
Fe3+ iron(III) Hg2+ mercury(II)
Pb2+ lead(II) Au+ gold(I)
Pb4+ lead(IV) Au3+ gold(III)
55
Practice
NH4ClO3 calcium phosphate
PbO
cupric sulfite
iron(III) carbonate
56
Naming Acids
Acid: substance yielding
H+ (proton) in water.
HCl = hydrogen chloride
HCl(aq) = hydrochloric acid
57
Naming Acids
58
Binary Acids
Anion Example
_ide Chloride Cl-
Acid Example
hydro_ic acid hydrochloric acid
HCl
59
Oxyacids
Anion Example
_ite chlorite ClO2-
Acid Example
_ous acid chlorous acid
HClO2
60
Oxyacids
Anion Example
_ate chlorate ClO3-
Acid Example
_ic acid chloric acid
HClO3
61
Naming Acids
Anion Acid chloride HCl hydrochloric acid
hypochlorite HClO hypochlorous acid
chlorite HClO2 chlorous acid
chlorate HClO3 chloric acid
perchlorate HClO4 perchloric acid
62
Acids: Try It !!!
H2SO4
HF
HNO3
HCN
H2SO3
63
Have Some Fun
Read about naming bases,
and hydrates on your own.
Quiz on compound names and formulas.
64
Mass Relationships in Chemical Reactions
The Mole
I
RULE !!!
65
Masses (not weights)
• proton = 1.673 x 10-24 g
• neutron = 1.675 x 10-24 g
• electron = 9.109 x 10-28 g
66
Atomic Mass Unit
1 amu 1/12 mass of C-12 atom
Mass of C-12 atom 12.00… amu
Mass of H 1 amu mass 1 p
Mass of O 16 amu
67
Average Atomic Mass
Periodic Table lists
weighted-average
atomic mass of the
naturally occurring
isotopes.
Determining Mass and Abundance of Isotopes.
68
Mass Spectrometry or Mass Spectroscopy
ion detector
vacuum pump
accel. field
ionizing e- beam
injection
Mass Spectrometry
69
1. Substance injected into vacuum tube. 2. Particles are ionized by e- beam. 3. Ionized particles are accelerated in elec. field. 4. Particles are deflected by magnetic field
that can be varied in strength. 5. Deflection depends on particle mass-to-
charge ratio, particle speed, and magnetic field strength.
6. Particles enter ion detector.
Strontium Example
70
Relative peak heights: • Sr-84 0.68 • Sr-86 12.00 • Sr-87 8.47 • Sr-88 100.00 Total 121.15
Strontium Weighted Average
71
.68 121.15
x 84.0 amu = 0.47 amu
12.00 121.15
x 86.0 amu = 8.52 amu
8.47 121.15
x 87.0 amu = 6.08 amu
100.00 121.15
x 88.0 amu = 72.64 amu
87.71 amu
Fractional abundance
72
Average Atomic Mass
e.g. Chlorine has 2 isotopes:
75.77% Cl-35 @34.969 amu
24.23% Cl-37 @36.966 amu
(.7577)(34.969 amu) +
(.2423)(36.966 amu) = 35.453 amu
73
Average Atomic Mass TRY IT: Of every 100 atoms of
Cu, 69.09 are Cu-63 (62.93 amu)
and the rest are Cu-65 (64.93
amu). Calculate the mass of Cu
as given on the Periodic Table.
74
Molecular Mass
The sum of atomic masses in amu
Molecular mass of HNO3:
H: 1 x 1.008 amu = 1.008 amu
N: 1 x 14.01 amu = 14.01 amu
O: 3 x 16.00 amu = 48.00 amu
= 63.02 amu
TRY IT: What is mass of N2O4?
75
The Mole
Mole:
number of C atoms in
exactly 12 g of C-12.
“chemist’s dozen”
76
Avogadro’s Number
1 mole =
6.022 x 1023 particles
77
Avogadro’s Number
602,204,500,000,000,000,000,000 # of grains of sand in the world!
# of stars in the universe!
78
The Mole
1 mole C atoms = 6.02 x 1023 atoms
1 mole CO2 molecules = 6.02 x 1023
molecules
1 mole pencils = 6.02 x 1023 pencils
1 mole Br- ions = 6.02 x 1023 ions
79
Molar Mass
Molar Mass
mass (in g) of one mole
e.g. 1 mol C atoms
= 6.02 x 1023 atoms C
= 12.0 g C
80
(6 x 12.0 g) + (12 x 1.0 g) + (6 x 16.0 g)
= 180.0 g
Mass of one mole of glucose,
C6H12O6
Molar Mass
81
Molar Mass: Other Terms
Gram atomic mass:
gam of Fe is 55.9 g
Gram molecular mass:
gmm of CO2 is 44.0 g
Gram formula mass:
gfm of NaCl is 58.4 g
82
Molar Mass
Mass (g) # Particles
Molar mass
g/mol
6.02 x 1023
part./mol Moles
Conversion Factors
83
Mole Practice
What is mass in g of one C-12 atom?
How many grams in one amu? (Hint: use carbon as an example.)
What is the molar mass of C6H8O6 ?
How many H atoms in 25.6 g urea?
(NH2)2CO
84
Percent Composition
% by mass of each element in a cmpd
= n x molar mass of element
molar mass of cmpd X 100
where n = element subscript
CO2
85
Percent Composition
What is % of O in CO2?
= n x molar mass of element
molar mass of cmpd X 100 %O
= 2 x 16.00
44.01 X 100
= 72.71%
86
Percent Composition
What is the mass of Cu (in kg)
in 3.71 x 103 kg of
chalcopyrite (CuFeS2)?
TRY IT !!
87
Percent Composition
If % composition is known,
a compound’s empirical formula
can be determined.
88
Empirical Formula
What is empirical formula of compound:
24.75% K 34.77% Mn 40.51% O?
Divide by smallest number KMnO4
assume 100 g
24.75 g 34.77 g 40.51g
.6330 mol .6330 mol 2.532 mol
go to moles
89
Empirical Formula
Sometimes it is not so easy!
What is emp. formula of Vitamin C?
C = 40.92%
H = 4.58%
O = 54.50%
Try it.
90
Empirical Formulas of C-H-O Compounds
Heat
O2
H2O CO2 absorber absorber
Sample
91
Example
If combustion of an 11.5 g sample
of ethanol produced 22.0 g of CO2
and 13.5 g of H2O, what is the
empirical formula of ethanol?
Try It
92
If combustion of a 21.4 g sample
of an unknown C-H-O compound
produced 29.5 g of CO2 and 24.1 g
of H2O, what is the empirical
formula of the compound?
93
Molecular Formulas
If the empirical formula and
approximate molar mass are
known, molecular formula can
be determined.
94
Molecular Formulas
Problem: A compound contains
only 1.52 g N and 3.47 g O and
has molar mass 95. What is it’s
molecular formula?
95
Molecular Formulas
N: 1.52 g 0.108 mol
O: 3.47 g 0.217 mol NO2
NO2 has empirical molar mass of
46.02 g compared to molar mass 95 g.
Molar mass
Emp.mol.mass = 95 g
46.02 g N2O4
96
Molecular Formulas
Problem: 200 g of a compound
contains 77.34 g C, 32.44 g H and
90.22 g N and has molar mass .
What is it’s molecular formula?
TRY IT !!
97
Chemical Equations
Chemical Reaction:
Atoms & molecules
react to form
new substances.
A chemical reaction is represented
by a chemical equation.
98
Evidence of Reactions
• Macroscopic evidence of chemical
reactions:
Heat Light
Gas Solid (precipitate)
Color change
• Chemical change involves creation of
new substances with new properties,
usually accompanied by a noticeable
energy change.
99
Chemical Equations
H2 + O2 H2O
But mass must be conserved!
Hydrogen burns to form water.
100
Chemical Equations
Balanced: 2H2 + O2 2H2O
Also show state:
2H2 (g) + O2 (g) 2H2O (l)
2H2 O2 2H2O
+
101
Balancing Equations
Identify reactants & products
Write correct formulas first
Change only coefficients
2H2 + 1O2 2H2O
Double check
102
Balancing Equations
Potassium chlorate decomposes
upon heating to yield oxygen
and potassium chloride:
KClO3 KCl + O2 3 2 2
103
Balancing Equations
Complete combustion of
Ethane (C2H6) yields carbon
dioxide and water.
Try it.
104
Types of Reactions
• Combination
• Decomposition
• Single Replacement
• Double Replacement
• Combustion
105
Combination
2 Br2 + 2 NaBr Na
106
Decomposition
CaCO3 CaO + CO2
107
Single Replacement
Cu AgNO3 2 Cu(NO3)2 2 + Ag
Cu + ZnCl2 No Reaction
+
108
Double Replacement
BaCl2 + H2SO4 2 BaSO4 + HCl
3 CrCl3(aq) + NaOH(aq)
3 Cr(OH)3(s) + NaCl(aq)
Soluble Ionic Compounds
109
Ionic compounds that are soluble in
water actually exist as separated ions.
CrCl3(aq) is actually:
Cr+3(aq) + 3Cl-(aq)
Net Ionic Equations
110
CrCl3(aq) + 3NaOH(aq)
Written as a net ionic equation is:
Cr(OH)3(s) + 3NaCl(aq)
Cr3+(aq) + 3 OH-(aq) Cr(OH)3(s)
Cl- and Na+ are “spectator ions”
111
Combustion
C2H5OH + 3 O2 2 CO2 + 3 H2O
112
Stoichiometry
The quantitative study of reactants
and products in a chemical reaction.
Interpret coefficients in chemical
equation as molecules or moles.
2H2 + O2 2H2O
2 mol H2 + 1 mol O2 2 mol H2O
113
Stoichiometry
2H2 + O2 2H2O
2 mol H2 1 mol O2 2 mol H2O
means stoichiometrically
equivalent to
114
Stoichiometry
2Li + 2H2O 2LiOH + H2
How many moles of H2 are
obtained from complete
reaction of 6.23 moles of Li ?
6.23 mol Li 1 mol H2 = 3.12 mol H2 x 2 mol Li
115
Stoichiometry
2Li + 2H2O 2LiOH + H2
How many g H2 from 81 g Li ?
81 g Li 1 mol Li 6.9 g Li
1 mol H2 2 mol Li
2.0 g H2 1 mol H2
= 12 g H2
x x
x
116
Stoichiometry
You try it! A key reaction in smog formation is:
2NO + O2 2NO2 1. How many g NO2 formed by
complete reaction of 1.44 g NO ?
2. How many molecules of NO2
formed from 0.00052 mg of NO ?
117
Limiting Reagents
How many cheese sandwiches
can be made from:
25 slices of cheese
12 tomatoes
7 heads of lettuce
4 slices of bread
Bread is the “limiting reagent”
118
Burn S in F2: S + 3F2 SF6
What if 4 mol S added to 20 mol F2 ?
Since 1 mol S 3 mol F2
Then 4 mol S 12 mol F2
Thus S is limiting and F2 is excess.
Limiting Reagents
119
Limiting Reagents
Urea is made by reacting
ammonia and carbon dioxide:
2NH3 + CO2 (NH2)2CO + H2O
If 641 g NH3 reacts with 1140 g CO2 •Which chemical is limiting ?
•How much urea is formed ?
•How many g of excess reagent is left ?
120
2NH3 + CO2 (NH2)2CO + H2O
641 g 1140 g
Start with NH3 & convert to urea:
641 g NH3 1 mol NH3 17.0 g NH3 2 mol NH3
1 mol urea
1 mol NH3
60.1 g urea = 1130 g urea
x x
x
121
2NH3 + CO2 (NH2)2CO + H2O
641 g 1140 g
Start with CO2 & convert to urea:
1140 g CO2 1 mol CO2 44.0 g CO2 1 mol CO2
1 mol urea
1 mol urea 60.1 g urea = 1560 g urea
x x
x
1130 g
122
2NH3 + CO2 (NH2)2CO + H2O
641 g 1140 g
1130 g
1560 g
Since NH3 yields less urea, it is
the limiting reagent, and 1130 g
urea will be produced.
CO2 is the excess reagent.
123
To calculate the excess amount of
CO2, start with limiting reagent, NH3.
641 g NH3 1 mol NH3 17.0 g NH3 2 mol NH3
1 mol CO2
1 mol CO2 44.0 g CO2
= 829 g CO2
2NH3 + CO2 (NH2)2CO + H2O
1140 g – 829 g = 311g excess CO2
x x
x
124
Limiting Reagents
You try one!
Al reacts with iron(III) oxide highly
exothermically. If 124 g Al reacts with
601 g iron(III) oxide:
•Write a balanced equation.
•How much aluminum oxide is formed ?
•How much excess reagent is left ?
125
Reaction Yield
from limiting reagent calculation
from experiment
% Yield = Actual Yield
Theoretical Yield x 100%
126
Reaction Yield Ammonia is produced by the Haber
process by reacting nitrogen gas and
hydrogen gas:
N2 + 3H2 2NH3
What is the % yield if 254 g NH3
are produced by reacting 246 g of N2 with 62 g of H2 ?
127
N2 + 3H2 2NH3 246 g 62 g 254 g
1. Determine limiting reagent
246 g N2 28.0 g N2
1 mol N2
1 mol N2
3 mol H2
2.02 g H2
1 mol H2 = 53.2 g H2
Thus N2 is the limiting reagent.
x x
x
128
2. Calculate theoretical yield of NH3
246 g N2 28.0 g N2
1 mol N2 1 mol N2
2 mol NH3
17.0 g NH3
1 mol NH3 = 299 g NH3
limiting
N2 + 3H2 2NH3 246 g 62 g 254 g
x x
x
129
3. Calculate % yield of NH3
vs.
299 g theo.
= 254 g
299 g x 100%
= 85.0 % yield
N2 + 3H2 2NH3 246 g 62 g 254 g
% Yield = Actual Yield
Theoretical Yield x 100%
130
TRY IT !!
Methanol is made by reacting hydrogen
gas and carbon monoxide gas.
Write balanced equation.
Calculate the % yield if 3.57 x 104 g of
CH3OH is produced by reacting 68.5 kg
carbon monoxide and 8.60 kg hydrogen.
Warm-up
132
What is 65 miles per hour in m/s?
Warm-up
133
Read the volume in the graduated cylinder.
Warm-up
134
The density of aluminum is 2.7g/mL.
What is its density in kg/m3?
Warm-up
135
Name or give formula:
Cl2O lithium nitride KMnO4 strontium phosphite NH3 AuSCN Hg2(NO3)2 Cu2O
Warm-up
136
Name or give formula:
HCN(aq) SnCl4
carbonic acid N2O4
magnesium nitride AlI3
Warm-up
137
What is the mass in grams of one water molecule?
Warm-up
138
83 g of a compound contains 36 g phosphorous, the remainder being oxygen. Its molar mass is 280 g/mol. What is its molecular formula?
Warm-up
139
For the reaction:
CH4(g) + 2O2(g) CO2(g) + 2H2O(l)
How many mL of water is produced by burning 17g CH4 in excess O2?
Warm-up
140
N2 is prepared by passing ammonia over solid copper(II)oxide. If 18.1 g NH3 reacts with 90.4 g copper(II)oxide, how many grams of N2 are formed?
2NH3(g) + 3CuO(s)
N2(g) +3Cu(s) + 3H2O(g)
Lab 2: Empirical Formula CuI or CuI2?
141
Beaker + clean Cu strip 56.1065 g
Beaker + CuxIy product 56.1245 g
Beaker after removing CuxIy 56.0978 g
• How do you determine mass iodine? • How do you determine mass copper?