Embed Size (px)
Citation preview
1
REVIEW
Copyright (c) 2011 by Michael A. Janusa, PhD. All rights reserved.
2
Measurement and Significant Figures
• To indicate the precision of a measured number (or result of calculations on measured numbers), we often use the concept of significant figures.
– Significant figures are those digits in a measured number (or result of the calculation with a measured number) that include all certain digits plus a final one having some uncertainty (first digit basically guessing).
cm
30.2246743 cm
30.22 cm
2 mL
1 mL
1.5 mL
3
• Rules for Significant Figures: – All nonzero digits are significant.
i.e. 111 1286
– Zeros between significant figures are significant.
i.e. 1001 20,006
– Zeros preceding the first nonzero digit are not significant.
i.e. 0.0002 0.00206
– Zeros to the right of the decimal after a nonzero digit are significant.
i.e. 0.00300 9.00 9.10 90.0
– Zeros at the end of a nondecimal number may or may not be significant. (Use scientific notation.)
i.e. 900 900.
111 3 SF 1286 4 SF
1001 4 SF 20,006 5 SF
0.0002 1 SF 0.00206 3 SF
0.00300 3 SF 9.00 3 SF 9.10 3 SF 90.0 3 SF
1, 2, or 3 SF 3 SF
4
• Scientific notation – is the representation of a number in the form
A. x 10n, where A is a number (sign digits only) with a single nonzero digit to the left of the decimal point and n is an integer or whole number.
900
300,000,000 (write with 3 SF)
0.0000301
843.4
0.00421
6.39 x 10-4
3.275 x 102
Note: exp or EE represents “x 10”
9 x 102 1 SF 9.0 x 102 2 SF 9.00 x 102 3 SF
3.00 x 108 3 SF
3.01x 10-5 3 SF
8.434 x 102 4 SF
4.21 x 10-3 3 SF
0.000639
327.5
5
• Number of significant figures refers to the number of digits reported for the value of a measured or calculated quantity, indicating the precision of the value. [Basically means if all quantities have X sign fig can’t report final answer with more than X sign figs: measurement or calculation dictates sign figs.]
– When multiplying and dividing measured quantities, give as many
significant figures as the least found in the measurements used. • 2.1 x 3.52 = 7.392 = 7.4
– Which gets us to rounding: left most digit to be dropped – 5 or greater add 1 to last digit to be retained, less than five leave alone – 1.2143 -- 1.21
– Multiple step calculation - Guard digit: 1.214
– When adding or subtracting measured quantities, give the same number of decimals as the least found in the measurements used.
• 84.2 (3 sign)
• +22.321 (5 sign)
• 106.521
• 106.5 (4 sign) arithmetic rules if combined ( ), x / , + -
6
3.38 – 3.012 = 0.368 = 0.37
2.4 x 10-3 + 3.56 x 10-1 =
0.0024
+0.356
0.3584 = 3.58 x 10-1
2.568 x 5.8 = 14.8944 = 3.55814 = 3.6
4.186 4.186 or 14.9 gives 3.55948
4.18 – 58.16 x (3.38 – 3.01) =
4.18 – 58.16 x (0.37) = 4.18 – 21.5192 = -17.3392 = -17
6.3 + 7.2 =
0.5256
13.5 = 25.685 = 25.7
0.5256
7
• An exact number is a number that arises when you count items or when you define a unit (conversion 12 in = 1 ft).
– For example, when you say you have nine coins in a bottle, you mean exactly nine (9.00000…. - infinite).
– When you say there are twelve inches in a foot, you mean exactly twelve.
– Note that exact numbers have no effect on significant figures in a calculation.
Measurement and Significant Figures
(cont’d)
HW 1
8
The Periodic Table • Metals, Nonmetals, and Metalloids –
generally, left of staircase metals,
staircase metalloids, right of staircase
nonmetals. This is important for
determining bond type, using proper
terminology, and making decisions.
metals
nonmetals
9
– A molecule is a definite group of atoms that are
chemically bonded together through sharing of
electrons (covalent bonding, generally nonmetal-
nonmetal including H).
Chemical Formulas; Molecular and
Ionic Substances
• Molecular substances
– A molecular substance is a substance that is
composed of molecules, all of which are alike.
– A molecular formula gives the exact number of
atoms of elements in a molecule (i.e. C2H6O).
– Structural formulas show how the atoms are bonded
to one another in a molecule.
i.e. ethanol, CH3CH2OH
involves covalent bond – share electrons
between atoms – typically nonmetal/nonmetal
involves ionic bond – transfer electrons between
atoms – attraction between charged particles –
typically metal/nonmetal or polyatomic ions
C : C
Na+
Cl-
10
– Although many substances are molecular, others are
composed of ions (charged particles, transfer of
electrons, ionic bonding, generally metal-nonmetal).
• Ionic substances
– An ion is an electrically charged particle obtained from
an atom or chemically bonded group of atoms by
adding or removing electrons.
– Sodium chloride is a substance made up of ions.
Na Cl + 1e- -
11
– The formula of an ionic compound is written by giving
the smallest possible whole-number ratio of different
ions in the substance.
Chemical Formulas; Molecular
and Ionic Substances • Ionic substances
– The formula unit of the substance is the group of
atoms or ions explicitly symbolized by its formula.
Covalent bond (share e-) Ionic bond (transfer e-/
attraction charged particles
nm –nm m – nm and charged ions
Molecules Formula unit
Molecular substance Ionic substance
Molecular formula formula
C : O Na
+Cl
-
12
– When an atom gains extra electrons, it becomes a
negatively charged ion, called an anion (more
electrons than protons).
• Ionic substances
– An atom that loses electrons becomes a positively
charged ion, called a cation (more protons than
electrons).
– An ionic compound is a compound composed of
cations and anions.
NaCl
CaBr2
Na2SO4
CO2
ionic or molecular; formula unit or molecule; ionic or covalent bonds involved?
ionic substance; formula unit; ionic bond
ionic substance; formula unit; ionic bonds
ionic substance; formula unit; ionic and covalent bonds
molecular substance; molecule; covalent bonds
13
Ions in Aqueous Solution
• Many ionic compounds (ionic bond/m-nm) dissociate into independent ions when dissolved in water
NaCl (s) Na+(aq) + Cl-(aq)
Soluble ionic compounds dissociate 100% - referred to as strong electrolytes – breaks into charged particles
Soluble salt
14
Ions in Aqueous Solution • Most molecular (covalent bond/nm-nm)
compounds dissolve but do not dissociate into ions, exception acids.
C6H12O6 (s) C6H12O6 (aq)
These compounds are referred to as nonelectrolytes; no charged particles; soluble but no ions formed.
How would the sodium sulfate dissolve?
Na2SO4 (s) 2Na+(aq) + SO42-(aq)
15
– Most ionic compounds contain metal and nonmetal
atoms; for example, NaCl.
Chemical Substances;
Formulas and Names • Ionic compounds
– You name an ionic compound by giving the name of
the cation followed by the name of the anion.
Sodium chloride, NaCl Calcium Iodide, CaI2
Potassium Bromide, KBr
– A monatomic ion is an ion formed from a single atom.
16
– Most of the main group metals form cations with the
charge equal to their group number.
How get charge for ions?
Rules for predicting charges on monatomic ions
– The charge on a monatomic anion for a nonmetal
equals the group number minus 8.
– Most transition elements form more than one ion, each
with a different charge (exceptions Cd2+, Zn2+, Ag+).
– Other important elements with variable charge
Pb4+, Pb2+ Sn4+, Sn2+ As5+, As3+ Sb5+, Sb3+
1+ 2+ 3+ 4+
0
1- 2- 3- 4-
varies
17
– Monatomic cations are named after the element. For
example, Al3+ is called the aluminum ion.
• Rules for naming monatomic ions
– If there is more than one cation of an element (charge),
a Roman numeral in parentheses denoting the charge
on the ion is used. This often occurs with transition
elements.
Na+ sodium ion Ca2+ calcium ion
Fe2+ iron (II) ion Fe3+ iron (III) ion
Older name: higher ox state (charge) – ic, / lower, -ous
Fe3+ ferric ion Fe2+ ferrous ion Cu2+ cupric ion
Cu+ cuprous ion Hg2+ mercuric ion Hg22+ mercurous ion
The names of the monatomic anions use the stem
name of the element followed by the suffix – ide. For
example, Br- is called the bromide ion. Br bromine
18
• The formula of an ionic compound is written by giving the smallest possible whole-number ratio of different ions in the substance.
Sodium chloride Na+ Cl-
Iron (III) sulfate Fe3+ SO42-
Chromium (III) oxide Cr3+ O2-
Calcium nitrate Ca2+ NO3-
Sodium phosphate Na+ PO43-
Strontium oxide Sr2+ O2-
NaCl
SrO
Na3PO4
Ca(NO3)2
Cr2O3
Fe2(SO4)3
19
Naming Binary Compounds
• NaF -
- lithium chloride
• MgO -
• MnBr2 -
- cobalt (III) oxide
- copper (II) chloride or cupric chloride
sodium fluoride
LiCl
magnesium oxide
manganese (II) bromide
Co2O3
CuCl2
20
– A polyatomic ion is an ion consisting of two or more
atoms chemically bonded together and carrying a net
electric charge.
– Table in book lists some common polyatomic ions.
Most are oxo anions – consists of oxygen with another
element (central element).
Chemical Substances;
Formulas and Names • Polyatomic ions
NO3- nitrate SO4
2- sulfate
NO2- nitrite SO3
2- sulfite
Most groups –ate, -ite differ by O
Mn, Br, Cl, I per- -ate, -ate, -ite, hypo- -ite
21
Ions You Should Know Polyatomic ions
• NH4+ - Ammonium
• OH- - Hydroxide
• CN- - Cyanide
• SO42- - Sulfate
• SO32- - Sulfite
• ClO4- - perchlorate
• ClO3- - chlorate
• ClO2- - chlorite
• ClO- - hypochlorite
• Hg22+ - mercury (I) or
mecurous
• S2O32- - thiosulfate
• SCN- - thiocyanate
• CNO- - cyanate
• MnO4- - permanganate
• O22- - Peroxide
• PO43- - Phosphate
• PO33- - Phosphite
• CO32- - Carbonate
• HCO3- - Bicarbonate or
Hydrogen Carbonate
• N3- - azide
• NO3- - nitrate
• NO2- - nitrite
• C2H3O2- - acetate
• Cr2O72- - dichromate
• CrO42- - chromate
• C2O42- - oxalate
• HSO4- - bisulfate or
hydrogen sulfate
• H2PO4- - dihydrogen
phosphate
22
SnSO4
sodium sulfite
Ca(ClO)2
barium hydroxide
potassium perchlorate
Cr2(SO4)3
magnesium nitride
Fe3(PO4)2
titanium (IV) nitrate
tin (II) sulfate or stannous sulfate
Na2SO3
calcium hypochlorite
Ba(OH)2
KClO4
chromium (III) sulfate
Mg3N2
iron (II) phosphate or ferrous phosphate
Ti(NO3)4
23
• molecular compounds
– Binary compounds composed of two nonmetals are
usually molecular and are named using a prefix
system (name same as ionic except must indicate
how many atoms are present using mono, di, tri,
etc.). No charges involved with molecular compounds
but we typically put more metallic compound first.
Chemical Substances;
Formulas and Names
NF3 F3N
24
– The name of the compound has the elements in the
order given in the formula.
• Binary molecular compounds
– You name the first element using the exact element
name.
– Name the second element by writing the stem name of
the element with the suffix “–ide.”
– If there is more than one atom of any given element,
you add a prefix (di, tri, tetra, penta, hexa, hepta, octa,
etc.)
Chemical Substances;
Formulas and Names
25
• Binary molecular compounds
– N2O3
– SF4
– chlorine dioxide
– sulfur hexafluoride
– Cl2O7
– HCl (g)
Name this compound but think about bonding:
MgCl2
Older names: water - H2O, ammonia – NH3,
hydrogen sulfide – H2S, nitric oxide – NO, hydrazine – N2H4
dinitrogen trioxide
sulfur tetrafluoride
ClO2
SF6
dichlorine heptoxide
hydrogen chloride
magnesium chloride; ionic bond no prefix
26
– Acids are traditionally defined as compounds with a
potential H+ as the cation.
• Acids
– Binary acids consist of a hydrogen ion and any single
anion. For example, HCl (aq) is hydrochloric acid.
– An oxoacid is an acid containing hydrogen, oxygen,
and another element. An example is HNO3, nitric acid.
Chemical Substances;
Formulas and Names
27
oxoacids
Anion prefix/suffix acid prefix/suffic
per- -ate ion per- -ic acid
-ate ion -ic acid
-ite ion -ous acid
hypo- -ite ion hypo- -ous acid
NO3- nitrate ion HNO3 nitric acid
NO2- nitrite ion HNO2 nitrous acid
ClO4- perchlorate ion HClO4 perchloric acid
HW 2
SO42- sulfate ion H2SO4 sulfuric acid
PO43- phosphate ion H3PO4 phosphoric acid
28
Molecular Weight and Formula
Weight, Molar Mass • The molecular weight of a substance is the sum
of the atomic weights of all the atoms in a molecule of the substance. – For, example, a molecule of H2O contains 2 hydrogen
atoms (at 1.01 amu each) and 1 oxygen atom (16.00 amu), giving a molecular weight of 18.02 amu.
– Molecular wt – mass one molecule
or
– do for 1 mole of substance called molar mass:
18.02 g H2O/mol H2O
29
Working with Solutions Molar Concentration
• When we dissolve a substance in a
liquid, we call the substance the solute
(being dissolved) and the liquid the
solvent (doing the dissolving).
– The general term concentration refers to the
quantity of solute in a standard quantity of
solution. There are many concentration terms but
we will concentrate on one.
30
• Molar concentration, or molarity (M),
is defined as the moles of solute
dissolved in one liter (cubic decimeter)
of solution.
solution of liters
solute of moles(M) Molarity
Working with Solutions Molar Concentration
solute + solvent volume
31
• The molarity of a solution and its volume are
inversely proportional. Therefore, adding water
makes the solution less concentrated. Most of time
will be using a stock solution and diluting to new
concentration. Basically using
ddcc VCVC
– So, as water is added, increasing the final volume, Vf,
the final molarity, Mf, decreases. Thing to realize
here is that M x V = mols: want new concentration of
substance take mols and divide by total volume
Working with Solutions Molar Concentration
32
• Mixture example – A solution is prepared by mixing 12.9 mL of 0.245 M HCl and 56.7 mL of
0.847 M HCl, then add 630.4 mL of water. Assuming the liquid volumes
are additive, calculate the molarity of HCl in the resulting solution.
HW 3 HClMHClmmol
HClmmol
0731.0mL 00.07
91.51
mL 630.4 mL 56.7 mL 12.9
91.51 HCl M
HClmmolHClmmolHClmmol 91.51025.48161.3
HCl) L
mol (0.847 x ) mL (56.7 HCl)
L
mol (0.245 x mL) (12.9 HCl of mmols
mol x mL =
L
____ = mol
tot mL L
mmol
Na2SO4
3.161 mmol
0.00452 M HCl
33
Solubility Rules for Ionic Compounds (Dissociates 100%)
1.) All compounds containing alkali metal cations and the ammonium ion are soluble.
2.) All compounds containing NO3-, ClO4
-, ClO3
-, and C2H3O2
- anions are
soluble.
3.) All chlorides, bromides, and iodides are soluble except those containing Ag+, Pb2+, or Hg2
2+.
4.) All sulfates are soluble except those containing Hg22+, Pb2+, Ba2+, Sr2+,
or Ca2+. Ag2SO4 is slightly soluble.
5.) All hydroxides are insoluble except compounds of the alkali metals and Ca2+, Sr2+, and Ba2+ are slightly soluble.
6.) All other compounds containing PO43-, S2-, CO3
2-, CrO42-, SO3
2- and most other anions are insoluble except those that also contain alkali metals or NH4
+.
Generally, compound dissolves > 0.10 M - soluble (aq)
< 0.01 M - insoluble (s)
in between - slightly soluble
(this class we will assume slightly soluble as soluble)
Hg2Cl2 (s) insoluble
KI (aq) soluble
Pb(NO3)2 (aq) soluble
34
Strong Acids (Ionizes 100%)
HCl, HBr, HI, HClO4, HNO3, H2SO4
Strong Bases (Dissociates 100%)
NaOH, KOH, LiOH, Ba(OH)2, Ca(OH)2,
Sr(OH)2
35
• A molecular/formula unit equation is one in which the reactants and products are written as if they were molecules/formula units, even though they may actually exist in solution as ions.
Calcium hydroxide + sodium carbonate
M.E.
Ca(OH)2
Ions in Aqueous Solution Molecular and Ionic Equations
+ Na2CO3 CaCO3 + NaOH 2 (aq)
strong base strong base soluble salt insoluble salt
(aq) (s) (aq)
s solid
l liquid
aq aqueous (acid/bases and soluble salts dissolve in water)
g gases
36
• An total ionic equation, however, represents strong electrolytes as
separate independent ions. This is a more accurate representation of the
way electrolytes behave in solution.
– A complete ionic equation is a chemical equation in which strong
electrolytes (such as soluble ionic compounds, strong acids/bases) are
written as separate ions in solution. (note: g, l, insoluble salts (s), weak
acid/bases do not break up into ions)
M.E.
Ca(OH)2 (aq) + Na2CO3 (aq) CaCO3 (s) + 2 NaOH (aq)
Total ionic
Ions in Aqueous Solution
Molecular and Ionic Equations
Ca2+ (aq) + 2OH- (aq)
strong base soluble salt insoluble salt strong base
+ 2Na+ (aq) + CO32-
(aq) CaCO3 (s) + 2Na+ (aq) + 2OH- (aq)
37
Net ionic equations.
– A net ionic equation is a chemical equation from
which the spectator ions have been removed.
– A spectator ion is an ion in an ionic equation that
does not take part in the reaction. M.E.
Ca(OH)2 (aq) + Na2CO3 (aq) CaCO3 (s) + 2 NaOH (aq)
Total Ionic Ca2+ (aq) + 2OH- (aq) + 2Na+ (aq) + CO3
2- (aq) CaCO3 (s) + 2Na+ (aq) + 2OH- (aq)
Net
Ca2+ (aq) + CO32-
(aq) CaCO3 (s)
38
Types of Chemical Reactions
• Oxidation-Reduction Reactions (Redox rxn)
– Oxidation-reduction reactions involve the
transfer of electrons from one species to another.
– Oxidation is defined as the loss of electrons.
– Reduction is defined as the gain of electrons.
– Oxidation and reduction always occur
simultaneously.
39
27.1 Reduction and Oxidation
Redox reactions – transfer of e-
reduction – oxidation reactions
Reduction – gain of e- / gain of H / lost of O
Fe3+ + 1e- Fe2+ (lower ox state)
note: must balance atoms and charges
40
Oxidation - loss of e- / loss of H / gain of O
Fe2+ Fe3+ + 1e- (higher ox state)
H2O + BrO3- BrO4
- + 2H+ + 2e-
(Br oxidized: charge 5+ 7+)
2H+ + 2e- H2 (H reduced: charge 1+ 0)
Oxidizing agent is species that undergoes reduction.
Reducing agent is species that undergoes oxidation.
Note: need both for reaction to happen; can’t have
something being reduced unless something else is being
oxidized.
Br + 3(-2) = -1
Br = -1 +6 = +5
Br + 4(-2) = -1
Br = -1 +8 = +7
41
27.3 Balancing Redox Reactions
- Must know charges (oxidation numbers) of species
including polyatomic ions
- Must know strong/weak acids and bases
- Must know the solubility rules
Oxidation Numbers – hypothetical charge assigned to the
atom in order to track electrons; determined by rules.
42
Rules to balance redox
1.) Convert to net ionic form if equation is originally in molecular form
(eliminate spectator ions).
2.) Write half reactions.
3.) Balance atoms using H+ / OH- / H2O as needed:
– acidic: H+ / H2O put water on side that needs O or H (comes from
solvent)
– basic: OH- / H2O put water on side that needs H but if there is no H
involved then put OH- on the side that needs the O in a 2:1 ratio
2OH- / H2O balance O with OH, double OH, add 1/2 water to
other side.
4.) Balance charges for half rxn using e-.
5.) Balance transfer/accept number of electron in whole reaction.
6.) Convert equation back to molecular form if necessary (re-apply
spectator ions).
Zn(s) + AgNO3(aq) Zn(NO3)2(aq) + Ag(s)
Total ionic:
Net ionic:
Zn(s) + Ag+(aq) + NO3-(aq) Zn2+(aq) + 2NO3
-(aq) + Ag(s)
Zn(s) + Ag+(aq) Zn2+(aq) + Ag(s)
43
Net: Zn(s) + Ag+(aq) Zn2+
(aq) + Ag(s)
Oxidation:
Reduction:
Balanced net:
Balanced eq:
Zn(s) Zn2+(aq) + 2e-
Ag+(aq) Ag(s) 1e- +
Zn(s) + 2 Ag+(aq) Zn2+(aq) + 2 Ag(s)
[ ] 2
Zn(s)
44
+ 2 AgNO3(aq) Zn(NO3)2(aq) + 2 Ag(s)
H+
Net: MnO4-(aq) + Fe2+
(aq) Mn2+(aq) + Fe3+
(aq)
Ox:
Red:
Balanced net:
Fe2+(aq) Fe3+(aq) + 1e- [ ] 5
MnO4-(aq) Mn2+(aq) + H2O(l) 4 8 H+(aq) + 5e- +
8 H+(aq) + MnO4-(aq) + 5 Fe2+(aq) Mn2+(aq) + 5 Fe3+(aq) + 4 H2O(l)
45
KMnO4(aq) + NaNO2(aq) + HCl(aq) NaNO3(aq) + MnCl2(aq) + KCl(aq) + H2O(l)
Net:
Ox:
Red:
Balanced net:
Balanced eq:
MnO4-(aq) Mn2+(aq) + NO2
-(aq) NO3
-(aq) + + H+(aq) + H2O(l)
NO2-(aq) NO3
-(aq)
MnO4-(aq) Mn2+(aq) + 4 H2O(l) 8 H+(aq) +
H2O(l) + + 2 H+(aq)
5 e- +
+ 2 e- [ ] 5
[ ] 2
2 MnO4-(aq) + 5 NO2
-(aq) + 16 H+(aq) + 5 H2O(l) 2Mn2+(aq) + 8 H2O(l) + 5 NO3-(aq) +10 H+(aq)
2 MnO4-(aq) + 5 NO2
-(aq) + 6 H+(aq) 2Mn2+(aq) + 3 H2O(l) + 5 NO3-(aq)
2 KMnO4(aq) + KCl 2 46
+ 5 NaNO2(aq) + 6 HCl(aq) 2MnCl2(aq) + 3 H2O(l) + 5 NaNO3(aq)
Net: OH-
CrI3 (s) + Cl2 (g) CrO42-
(aq) + IO4-(aq) + Cl-(aq)
Ox:
Red:
Balanced net:
CrI3(s) CrO42-(aq) + IO4
-(aq)
Cl2(g) Cl-(aq) 2
3 32 OH-(aq) + + 16 H2O(l)
2 e- +
+ 27 e- [ ] 2
[ ] 27
64 OH-(aq) + 2 CrI3(s) + 27 Cl2(g) 2 CrO42-(aq) + 6 IO4
-(aq) + 54 Cl-(aq) + 32 H2O(l)
HW 4
47
