Download Metals

EIT Chemistry Review
S2012
Dr. Mack
Chemistry is really…
Chem
is
try
To do chemistry, you must:
Download available at:
Practice
www.csus.edu/indiv/m/mackj/
Practice
Or look for Jeffrey Mack under CSUS faculty web pages
Practice
Part 1
1
Where do we begin…
The Periodic Table
3
The Modern Periodic Table
The modern organization of the periodic table came about as a result of the work of Dimitri Mendeleyev Characteristics:
•Ordered elements by atomic mass
•Repeating pattern of properties
•Elements with similar properties in the same column
Periodic Law – when the elements are arranged in order of increasing atomic mass, certain sets of properties recur periodically.
Patterns used to predict properties of undiscovered elements
4
6
1
The Periodic Table:
Metals and Nonmetals
s
ta l
e
nm
No
Metaloids
ids
llo i ta
”
me em tors
s
“
uc
nd
co
Metals
The Periodic Table:
7
8
The modern periodic table is defined by:
Groups (families)
(columns down)
Tin is in group 4A (14) in the 5th period. Periods (rows across)
9
10
2
Today’s Periodic table
inert or “noble” gasses:
Group 1A elements are also know as the ““Alkali Metals”
Alkali Metals” as they form basic salts
Metals form Cations
Cations
non‐metals form anions
anions Metalloids can do both
11
Today’s Periodic table
12
Today’s Periodic table
Group 2A elements are also know as the ““Alkaline earth Metals”
Metals” as they are only found in the ground as metal salts (carbonates)
13
Group 2B – 8B elements are also know as the ““Transition Metals”
Transition Metals”. They may be found in the earth as pure metals or as ores (salts).
14
3
Today’s Periodic table
Today’s Periodic table
Group 7A elements are also know as the ““Halogens”
Halogens”. They form acids with hydrogen and exist as diatomic molecules. (F2, Cl2…)
Group 8A elements are also know as the ““Noble gasses”
Noble gasses”. They are inert to reaction for the most part. He is found underground!
15
16
Why do scientists use chemical symbols like C, Al and The Representation of Matter:
Fe to represent atoms? In chemistry we use chemical formulas and symbols to represent matter.
Why?
Because atoms are really small!
We are “macroscopic”: large in size on the order of 100’s of cm
10‐2 m
Atoms and molecules are “microscopic”: on the order of 10‐12 cm
108 Angstroms
17
1Angstrom = 10‐10 m
18
4
Chemical Symbols and Formulas:
Chemical Symbols and Formula:
To what we can’t see!
What we observe…
Elements:
H = hydrogen
O = oxygen
C = carbon
Molecules:
H2 = hydrogen
O2 = oxygen
some of the names are not systematic!
Chemical symbols (H2O) allow us to connect…
H2O = water
Uh‐
Uh‐Oh! this is confusing…
this is confusing…
Yes it is…
Yes it is…
Get over it and get used to it!
CH4 = methane
19
Al
B
Cr
Co
Cu
F
Fe
Au
Pb
Ag
Hg
P
K
Na
S
20
Atoms, Compounds, Molecules and Ions:
Aluminum
Boron
Chromium
Cobalt
Copper
Fluorine
Iron
Gold
Lead
Silver
Mercury
Phosphorus
Potassium
Sodium
Sulfur
An atom
atom is the smallest indivisible indistinguishable unit of a pure element. An ion
ion is an atom (or molecule) with a formal electrical charge.
An anion
anion is an atom (or molecule) that has a negative charge.
The number of electrons > number of protons
A cation
cation is an atom (or molecule) has a positive charge
The number of electrons < number of protons
21
22
5
Compounds:
Molecules and Ionic Compounds:
A compound is a distinct substance that contains two or more elements combined in a definite proportion by weight.
A molecule is the smallest uncharged individual unit of a compound formed by two or more atoms.
Atoms of the elements that constitute a compound are always present in simple whole number ratios. They are never present as fractional parts.
Examples:
AB
Never:
A2B
Ionic compounds are made of positively and negatively charged ions. A molecule
molecule can exist as an entity on its own.
AB2
An ionic compound
ionic compound is represented by a formula unit that describes the simplest ratio of cations to anions.
A½B
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24
Compounds fall into one of two classes:
Ionic or Inorganic Compounds:
Inorganic Salts
Molecules
metal cation
non‐metal
+
+
non‐metal or
polyatomic anion
non‐metal
(no cations or anions)
Metals and Nonmetals
1
IA
1
18
VIIIA
2
IIA
13
IIIA
2
3
4
5
6
3
IIIB
4
IVB
5
VB
6
VIB
7
VIIB
8
9
VIIIB
10
Metals
Cations
11
IB
12
IIB
14
IVA
15
VA
16
VIA
17
VIIA
s
ta l
e
nm ns
No Anio
7
The two use different formalisms for naming…
25
26
6
Chemical Nomenclature:
Molecular Compounds:
Nonmetals and Nonmetals
1
IA
1
18
VIIIA
2
IIA
13
IIIA
14
IVA
15
VA
16
VIA
17
VIIA
Go to the lab page on my website and download the worksheet!
2
3
IIIB
3
4
IVB
5
VB
6
VIB
7
VIIB
8
9
VIIIB
10
11
IB
s
ta l
e
nm
No
12
IIB
4
5
6
7
27
Ion Charges:
28
Binary Compounds: Metal & non‐Metal
+1
+3
‐3 ‐2
Metal of fixed oxidation (charge) state combined with a non‐metal.
‐1
+2
non‐metal takes on “ide” suffix
Examples:
Variable
29
Cation
Anion
Formula
Name
K+
Cl−
KCl
Ca2+
O2‐
CaO
Calcium Oxide
Na+
S2‐
Na2S
Sodium sulfide
Al3+
S2‐
Al2S3
Aluminum sulfide
Potassium chloride
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7
Metals of variable charge (transition) with a non‐metal
Examples:
Some common polyatomic ions:
modify transition metal name with roman numeral
Cation
Anion
Pb2+
Cl−
Formula
Name
PbCl2
lead (II) chloride
NH4+
H3O+
CO32–
HCO32–
NO2–
NO3–
SO42–
SO32–
PO43–
C2H3O2–
pronounced: lead ‐
lead ‐ two ‐
two ‐ chloride
Pb4+
Cl−
PbCl4
lead (IV) chloride
Fe3+
O2−
Fe2O3
Iron (III) oxide
ammonium
hydronium
carbonate
hydrogen carbonate or bicarbonate
nitrite
nitrate
sulfate
sulfite
phosphate
acetate
31
Ternary Compounds: Those with three different elements
metal of fixed charge with a complex ion
Cation
Anion
Formula
33
Metal of variable charge transition with a complex ion
Cation
Anion
Formula
Name
Name
K+
OH−
KOH
Potassium hydroxide
Ca2+
OH−
Ca(OH)2
Calcium hydroxide
Na+
SO 24−
Na2SO4
Sodium sulfate
Al3+
SO 24−
Al2(SO4)3
Aluminum sulfate
34
Fe3+
NO3−
Fe(NO3)3
Iron (III) nitrate
Fe2+
NO −2
Fe(NO2)2
Iron (II) nitrite
35
8
Examples:
Formula Name:
Non‐
Non‐metal with a non‐
metal with a non‐metal
When non‐metals combine, they form molecules.
They may do so in multiple forms: CO
CO2
Because of this we need to specify the number of each atom by way of a prefix.
BCl3
boron trichloride
SO3
sulfur trioxide
NO
nitrogen monoxide
we don’t write:
1 = mono
2 = di
5 = penta
3 = tri
6 = hexa
4 = tetra
N2O4
7 = hepta
nitrogen monooxide
or mononitrogen monoxide
dinitrogen tetraoxide
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37
Writing formulas for acids and Bases
CO2
carbon dioxide
carbon monoxide
CO
P2O5
diphosphorous pentaoxide
nitrogen trihydride
NH3
•An acid
acid is a substance that produces H+ when dissolved in water.
•Certain gaseous molecules become acids when dissolved in water.
•A base
base produces OH− when dissolved in water.
•Bases often are Group I and Group II hydroxide salts.
(ammonia)
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9
H+ and S2‐
Type I Acids: Acids derived from –
–ide anions.
it takes 2 H+ to cancel one S2‐
The names for these acids follows the formula:
“hydro”
+
the root of the ide anion +
ic “acid”
Anion:
Acid:
chloride
HCl
hydrochloric acid
fluoride
HF
hydrofluoric acid
H2S
Name:
hydro sulfuric acid
40
Examples: Oxy Acids: Those derived from –ate anions.
Anion:
Acid:
Name:
NO3−
HNO3
nitric acid
−
(chlorate) ClO3
HClO3
chloric acid
(sulfate)
SO24 −
H2SO4
sulfuric acid
(acetate)
C2 H 3O2−
HC2H3O2
acetic acid
vinegar
(nitrate)
41
Oxy Acids: Those derived from –ite anions.
root name of the anion with –ous replacing the ‐ite
Anion:
42
Acid:
Name:
(nitrite)
NO2−
HNO2
nitrous acid
(chlorite)
ClO2−
HClO2
chlorous acid
(sulfite)
SO32 −
H2SO3
sulfurous acid
43
10
Common Names:
Bases:
Ba(OH)2
barium hydroxide
sodium hydroxide
NaOH
NH4OH
ammonium hydroxide
H2O
water
ammonia
NH3
CH4
methane
NO
nitric oxide
N2O
nitrous oxide
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45
Modern Atomic Theory:
Practice:
•Atoms are made of protons, neutrons and electrons.
N2SO4
sodium sulfate
barium carbonate
BaCO3
FeO
Iron (II) oxide zinc phosphide
Zn3P2
The charge on the proton is +1
NiBr2
nickel (II) bromide
There is no charge on the neutron
•The nucleus of the atom carries most of the mass.
• It consists of the protons and neutrons surrounded by a cloud of electrons.
The charge on the electron is –1
The Atomic Number
Atomic Number or number of protons in the nucleus defines an element.
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11
The Composition of an Atom: Atomic Number, Z
An element’s identity is defined by the number of protons in the nucleus: Z
The atom is mostly
empty space
Atomic number
13
•protons and neutrons in the nucleus.
Al
•the number of electrons is equal to the number of protons.
Atom symbol
26.981
Atomic weight
•electrons in space around the nucleus.
49
Ch. 2
Isotopes, Atomic Numbers, and Mass Numbers
50
Example: Selenium‐75
Elemental Isotopic Symbols: For a given element “XX”, an isotope is written by:
Atomic number (Z) = number of protons in the nucleus.
75 protons + neutrons
Mass number (A) = total number of nucleons in the nucleus (i.e., protons and neutrons).
75
34 Se
34 protons
One nucleon has a mass of 1 amu
(Atomic Mass Unit
Atomic Mass Unit) a.k.a “Dalton
Dalton” or u
A
Z
X
Isotopes have the same Z but different total number of nucleons (A).
51
protons = 34
electrons = 34
neutrons = 75‐34 = 41
52
12
Ch. 2
The average weighted atomic mass is determined by the following mathematical expression:
Average mass of a Cl atom
mass of a Cl‐35 atom
m Cl (u) = m Cl‐35 ×
fraction that are Cl‐35
mass of a Cl‐37 atom
abundance + m Cl‐37 ×
of Cl‐35
Avogadro’s Number and the Mole
Ch. 2
The concept of a mole is defined so that we may equate the amount of matter (mass) to the number of particles (mole).
fraction that are Cl‐37
abundance of Cl‐37
The Standard is based upon the C‐12 isotope.
The mass of one 12C‐atom is 1.99265 × 10‐23 g.
35.45 u = 34.96885u × 0.7553 + 36.96590 × 0.2447
The atomic mass of 12C is defined as exactly 12 u.
(4 sig. fig)
This is the value that is reported on the periodic table.
Note that: 0.7553 + 0.2447 = 1.0000 (100%)
Therefore: 1u = (the mass of one 12C atom ÷12) = 1.66054 × 10‐24g = 1.66054 × 10‐27 kg
53
Ch. 2
Avagagro’s Number
54
Ch. 2
Molar Masses
Since we can equate mass (how much matter) with moles (how many particles) we now have a conversion factor
that relates the two.
Since one mole of 12C has a mass of 12g (exactly), 12g of 12C contains 6.022142 ×
1023
C‐12 atoms.
But carbon exists as 3 isotopes: C‐12, C‐13 &C‐14
mols
The average atomic mass of carbon is 12.011 u. molar mass (g/mol) = grams
The Molar Mass of a substance is the amount of matter that contains one‐mole or 6.022 × 1023 particles.
From this we conclude that 12.011g of carbon contains 6.022142 × 1023 C‐atoms
×
aka: Avogadro's number (N
Avogadro's number (NA)
Is this a valid assumption?
The atomic masses on the Periodic Table also represent the molar masses of each element in grams per mole (g/mol)
Yes, since NA is so large, the statistics hold.
55
56
13
So if you have 12.011g of carbon…
you have 6.022×1023 carbon atoms!
Ch. 2
Grams, Moles and Molar Mass:
The molar mass of an atom is a conversion factor that relates mass (grams) to moles
moles and vice versa.
So if you have 39.95g of argon…
you have 6.022×1023 argon atoms!
(how much matter)
(number of atoms)
grams ÷ molar mass = moles
if you have a mole of dollar bills…
g ×
you are Bill Gates…
you have 6.022×1023 bucks!
mol
= mol
g
Avogadro's number (NA) relates moles numbers of individual particles:
and if you have 6.022×1023 avocados…
you have…
1 mol ×
a “guacamole”
6.022 × 1023 particles
= 6.022 × 1023 particles
mol
58
Question: How many magnesium atoms are there in 150.0 g of Question: magnesium?
Ch. 2
Ch. 2
The molar mass of an ionic compound is the molar mass of one formula unit of the compound.
Solution: Use the molar mass of Mg from the periodic table.
150.0g Mg ×
59
For the compound Al2(SO4)3 the molar mass would be:
1 mol Mg
6.022 x 1023 Mg atoms
×
24.305g Mg
1 mol Mg
2 × (26.98 g/mol) = 3.717 × 1024 Mg atoms
(4 sig figs) big number as expected!
2 Al’s
+ 3 × (32.07 g/mol) 3 S’s
+ 3 × 4 × (16.00g/mol) 3 × 4 O’s
342.17 g/mol
60
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14
How many oxygen atoms are there in 25.1g of chromium (III) acetate?
(a) How many grams of nitrogen will be needed to produce 0.384 mols of ammonia? step 1: write the correct chemical formula…
Cr3+ & C2H3O2−
step 2: calculate the molar mass…
3 2(g)
N2(g) + H
Cr(C2H3O2)3
229.13 g/mol
0.384 mols NH3 ×
step 3: use dimensional analysis to solve the problem…
3×2 = 6
25.1g Cr(C 2 H3O 2 )3 × 1mol Cr(C 2 H3O 2 )3
229.13g Cr(C2 H 3O 2 )3
6.022 × 1023 O − atoms
×
=
1mol O
6mol O
×
1mol Cr(C 2 H 3O 2 )3
3.96 ×
1023
1mol N 2 28.02g N 2
×
2mol NH 3 1mol N 2
= 5.38g N2
(b) How many moles of hydrogen are needed to combine with 5.84g of nitrogen?
5.84g N 2 ×
O‐atoms 3 sf
3 sf
2 NH3(g)
1mol N 2
3mol H 2
×
= 0.625mols H 2
28.02 gN 2
1mol N 2
Ch. 2
63
How many grams of nitrogen are needed to completely
react with 0.525 g of hydrogen in the formation of
ammonia?
64
How many grams of nitrogen are needed to completely
react with 0.525 g of hydrogen in the formation of
ammonia?
To answer this question, one must go through moles.
3 2(g)
N2(g) + H
3 2(g)
N2(g) + H
2 NH3(g)
0.525gH 2×
The equation relates moles, not mass.
2 NH3(g)
1 mol H 2 1mol N 2
28.02g N 2
×
×
= 2.43g N 2
2.02g H 2 3mol H 2
1mol N 2
How many grams of ammonia form?
g H2
mols H2
mols N2
g N2
0.525gH 2×
65
1 mol H 2 2 mol NH 3 17.04g NH 3
×
×
= 2.95g NH 3
1mol NH 3
2.02g H 2
3mol H 2
66
15
How many grams of O2 will result form the decomposition of 15.6g of potassium chlorate?
Step 1. How many grams of O2 will result form the decomposition of 15.6g of potassium chlorate?
Write the chemical equation
2 KClO3(s)
3 O2(g)
2 KClO3(s)
3 O2(g)
+ KCl(s)
2
+ KCl(s)
2
Step 2: Balance
15.6g KClO3×
1mol KClO3
3mol O2
32.00g O2
×
×
122.55g KClO3
2mol KClO3 1mol O2
Step 3: Set up the solution
g KClO3
mols KClO3
mols O2
= 61.1g O2
g O2
67
Chemical Reactions
Ch. 3
A chemical reaction
chemical reaction can be described by a chemical equation
chemical equation.
All chemical equations have three parts:
Reactants
Æ
68
How does one know when a reaction has occurred?
Generally, what you end up with looks nothing like what you started with.
Before
Ch. 3
After
Products
When a chemical changes occurs:
Atoms, Compounds or Molecules
Æ
New Compounds, Molecules or atoms
69
70
16
Indicators of a Chemical Reaction:
Ch. 3
Gas Evolution
Indicators of a Chemical Reaction:
Temperature change
71
Indicators of a Chemical Reaction:
Ch. 3
Ch. 3
Color Change
72
Indicators of a Chemical Reaction:
Ch. 3
Precipitate Formation
73
74
17
The chemical equation for the formation of water can be visualized as two hydrogen molecules reacting with one oxygen molecule to form two water molecules:
Quantitative Information About Chemical Reactions Mass must be conserved in a chemical reaction.
Total mass of reactants
2H2 + O2 → 2H2O
Total mass of products
=
Chemical equations must therefore be balanced for mass.
Numbers of atoms on the reactant side
Numbers of atoms on the product side
=
11.2g
+ 88.8g → 100.0g
mass is conserved!
76
When writing chemical reactions one starts with:
Balancing Chemical Reactions: Example
C2H6
+
2 C’s & 6 H’s
→
O2
2 O’s
balance last
CO2
+
Reactants
H2O
1 C & 2 O’s
+
N2(g) + 3H2(g)
→
O2
CO2
Ch. 3
products
2 H’s & 1 O
balance H first
2 2H6
___C
77
3 2O
6
___ H
+
2NH3(g)
Some reactions can also run in reverse:
balance C next
2C2H6 +
O2
→
4 2
___ CO
+
6H2O
2NH3(g)
balance O
2C2H6 +
4 C’s
7 2
____ O
12 H’s 14 O’s
→
4CO2 +
4 C’s
N2(g) + 3H2(g)
Under these conditions, the reaction can be written:
6H2O
3H2 (g) + N2 (g) R 2NH3 (g)
12 H’s 14 O’s
Dynamic Equilibrium!
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18