Download AtomMoleculeNaming_G1

The Components of Matter
2-1
The Components of Matter
• Elements, Compounds, and Mixtures
• The Observations That Led to an Atomic View of Matter
• Dalton’s Atomic Theory
• The Observations That Led to the Nuclear Atom Model
• The Atomic Theory Today
• Elements: A First Look at the Periodic Table
• Compounds: Introduction to Bonding
• Formulas, Names, and Masses of Compounds
• Mixtures: Classification and Separation
2-2
Element: Atoms or Molecules
Element consists of only one type of atom. It cannot be broken down
into any simpler substances by physical or chemical means. Example:
Na, Au, He, Ne
Molecule consists of two or more atoms that
are chemically bound together and thus behaves
as an independent unit. Example: F2, O3 (ozone)
S6, S8, S12, P4 (white phosphorus)
2-3
Compound vs. Mixture
Compound – a substance
composed of two or more elements
that are chemically combined.
Figure 2.1
Mixture – a group of two or more
elements and/or compounds that
are physically intermingled.
2-4
Law of Multiple Proportions
If elements A and B react to form two compounds, the
different masses of B that combine with a fixed mass of A
can be expressed as a ratio of small whole numbers.
Example: Carbon Oxides A & B
Carbon Oxide I : 57.1% oxygen and 42.9% carbon
Carbon Oxide II : 72.7% oxygen and 27.3% carbon
2-5
Assume that you have 100 g of each compound.
Carbon Oxide I
g oxygen/100 g compound 57.1
72.7
g carbon/100 g compound
27.3
72.7 =
2.66
27.3
g oxygen/g carbon
42.9
57.1
= 1.33
42.9
2.66 g O/g C in II
1.33 g O/g C in I
2-6
Carbon Oxide II
=
2
1
Dalton’s Atomic Theory
Dalton postulated that:
1.
2.
3.
4.
2-7
All matter consists of atoms; tiny indivisible particles of
an element that cannot be created or destroyed.
Atoms of one element cannot be converted into atoms
of another element.
Atoms of an element are identical in mass and other
properties and are different from the atoms of any
other element.
Compounds result from the chemical combination of a
specific ratio of atoms of different elements.
Look inside the atoms: Cathode rays
2-8
Millikan’s experiment for measuring an
electron’s charge.
2-9
Millikan’s findings were used to calculate the mass of an
electron.
determined by J.J. Thomson
and others
mass of electron =
mass
charge
x charge
= (–5.686 x 10–12 kg/C) x (–1.602 x 10–19 C)
= 9.109 x 10–31 kg = 9.109 x 10–28 g
2-10
Rutherford’s a-scattering experiment and discovery of the atomic nucleus.
2-11
General features of the atom
The atom is an electrically neutral, spherical entity composed of a
positively charged central nucleus surrounded by one or more
negatively charged electrons.
The atomic nucleus consists of protons and neutrons.
2-12
Properties of the Three Key Subatomic Particles
Charge
Name
Relative Absolute (C)*
(Symbol)
Mass
Relative
(amu)†
Absolute (g) Location in
Atom
Proton
(p+)
1+
+1.60218x10-19 1.00727
1.67262x10-24 Nucleus
Neutron
(n0)
0
0
1.67493x10-24 Nucleus
Electron
(e-)
1-
-1.60218x10-19 0.00054858 9.10939x10-28 Outside
nucleus
* The
†
2-13
1.00866
coulomb (C) is the SI unit of charge.
The atomic mass unit (amu) equals 1.66054x10-24 g.
Atomic Symbol, Number and Mass
Figure 2.8A
X = Atomic symbol of the element
A = mass number; A = Z + N
Z = atomic number
(the number of protons in the nucleus)
N = number of neutrons in the nucleus
2-14
Isotopes
Isotopes are atoms of an element with the same number
of protons, but a different number of neutrons.
Isotopes have the same atomic number, but a different
mass number.
Figure 2.8B
2-15
Getting familiar with Isotope Symbol
PROBLEM: Silicon (Si) has three naturally occurring isotopes:
28Si, 29Si, and 30Si. Determine the number of protons,
neutrons, and electrons in 29Si
29Si
2-16
has 14p+, 14e–, and 15n0 (29 – 14 = 15)
Tools of the Laboratory
Ionization through electron bombardment:
Formation of a positively charged neon particle (Ne+)
2-17
Tools of the Laboratory
The mass spectrometer: Determine the molar mass of compounds
2-18
Abundance of
Isotopes
11B
10B
• Abundance: Percentage of certain isotope among all
the isotopes.
Example: 10B 19.9% vs. 11B 80.1% B-10 B-11 B-11 B-11
B-11
Because of different mass among isotopes, atomic mass
is the weighted average among all isotopes.
• Average mass = ATOMIC WEIGHT
• Atomic mass = S abundance  mass number
=abundance1 x #mass1 + abundance2 x #mass2 + …
2-19
Calculating the Atomic Mass of an Element
PROBLEM: Silver (Ag, Z = 47) has two naturally occurring isotopes,
107Ag and 109Ag. From the mass spectrometric data
provided, calculate the atomic mass of Ag.
Isotope
107Ag
109Ag
Mass (amu)
Abundance (%)
106.90509
108.90476
51.84
48.16
= 107.87amu
2-20
Practice: Isotopes & Atomic Weight
Assuming the mass of isotope equals the mass number in
amu:
• 6Li = 7.5% abundant and 7Li =
– Atomic weight of Li = ______________
•
28Si
= 92.23%, 29Si = 4.67%, 30Si = 3.10%
– Atomic weight of Si = ______________
2-21
6.925
28.11
The modern periodic table.
2-22
The formation of an ionic compound.
Transferring electrons from the atoms of
one element to those of another results in
an ionic compound.
2-23
© McGraw-Hill Education/Steven Frisch, Photograher
What affects the strength of ionic bonding?
#Charge and Size(Distance)
2-24
#Charge on the Ions:
The relationship between ions formed and the nearest noble gas
2-25
Elements that occur as molecules
2-26
Chemical Formulas
• A chemical formula consists of
– element symbols with numerical subscripts.
• The chemical formula indicates the
– type and number of each atom present in the
smallest unit of a substance.
– Example:
One unit of (NH4)3PO4 = 3 N atoms + 12 H atoms +
1 P atom + 4 O atoms
2-27
Binary Molecular Compounds:
Two Nonmetals (such as CO2)
1. Name first element in formula first
–
use the full name of the element
2. Name the second element in the formula with an -ide
–
as if it were an anion, however, remember these
compounds do not contain ions!
3. Use a prefix in front of each name to indicate the
number of atoms
a) Never use the prefix mono- on the first element
28
2-28
Subscript - Prefixes
• 1 = mono-;
– not used on first
nonmetal
• 2 = di• 3 = tri• 4 = tetra-
•
•
•
•
•
5 = penta6 = hexa7 = hepta8 = octadrop last “a” if
name begins with
vowel
29
2-29
Practice:
Naming Molecular Compounds
•
•
•
•
•
•
•
CO
ClO3
SO2
P2O5
N2O4
IF7
SF6
30
2-30
Ionic Compounds
• Made of Cation (+) and Anion (-)
• Name: Cation Anion
example: NaCl Sodium Chloride
– Cation:
• Type I metal
• Type II metal
• Polyatomic ion: ammonium NH4+
– Anion:
• Nonmetal: Chloride Cl-, Oxide O2• Polyatomic ion: SO42- , OH- , NO331
2-31
Metal Cations: Type I
Type I (Groups IA, IIA, AZA)
– only have one possible charge
• Groups IA, IIA, Ag+, Zn2+, Al3+
– Charge by position on the
Periodic Table
• IA = +1, IIA = +2,
• Ag+ (IB), Zn2+(IIB) Al3+(IIIA)
How do you know a
metal cation is Type II?
its not Type I !!!
32
2-32
Metal Cations: Type II
Type II: Metal ions that are
other than Type I
Common Examples: Fe2+/3+,
Cu+/2+, Cr3+/6+, Mn2+/4+,
Mn2+/4+, Pb2+/4+, Sn2+/4+, etc )
– have more than one possible
charge
– determine charge by charge on
anion
How do you know a
metal cation is Type II?
its not Type I !!!
33
2-33
Monatomic Nonmetal Anion
n(A )
• How much is the charge? the position on the
Periodic Table
• Name of the anion: change ending on the
element name to –ide
4A = -4
5A = -3
6A = -2
7A = -1
C4- = carbide
N3- = nitride
O2- = oxide
F- = fluoride
Si4- = silicide
P3- = phosphide
S2- = sulfide
Cl- = chloride
34
2-34
Type I Binary Ionic Compounds
Binary: only two kinds of ions in one UNIT
Example: MgO, CaCl2
• Metal listed first in formula & name
1. Metal Cation  Nonmetal Anion
2. Cation name <= Metal name: Magnesium,
Calcium
3. Nonmetal anion <= Nonmetal name ends with
–ide: Oxide, Chloride
35
2-35
Type II Binary Ionic Compounds
Metal listed first in formula & name
1. Metal cation  Nonmetal anion
2. metal cation  Metal(Roman Numeral): to indicate
its charge. Iron(II), Copper(I)
–
–
determine charge from anion charge
Common Type II cations in Table 5.5
3. Nonmetal anion  Nonmetal name ended with –ide:
Chloride, Oxide
Example: Iron(II) chloride, Copper(I) oxide
36
2-36
How to find the charge on Type II
metal ions?
• Example: Name Compound Fe2O3
Since the sum of all charges equals zero, the charge on
iron ions are unknown and oxide ion each has –2
charge, then we have
2 x Fe + 3 x (-2) = 0
Fe = +3, each iron ion has a charge of +3
Name: iron(III) oxide
Key: knowing the charge on ANIONs!
37
2-37
Practice: Naming Ionic compounds
•
•
•
•
•
•
•
HgF2
CuI2
CaCl2
Fe2S3
SnCl4
Mg3N2
Ag2S
38
2-38
•
•
•
•
•
•
•
Answer key: names of ionic
compounds
HgF2 = Mercury(II) fluoride
CuI2 = copper(II) iodide
CaCl2 = calcium chloride
Fe2S3 = Iron(III) sulfide
SnBr4 = tin(IV) bromide
Mg3N2 = magnesium nitride
Ag2S = silver sulfide
39
2-39
Polyatomic Anions: -ATE ions
CO32-
NO3-
carbonate
nitrate
SiO32-
PO43-
SO42-
ClO3-
silicate
phosphate
sulfate
chlorate
AsO43-
SeO42-
BrO3-
arsenate
selenate
bromate
IO3iodate
2-40
40
Periodic Pattern of Polyatomic Ions
-ate groups
IIIA
3BO3
IVA
VA
VIA
VIIA
2CO3
NO3
2SiO3
3PO4
2SO4
ClO3
3AsO4
2SeO4
BrO3
2TeO 4
IO3
41
2-41
Patterns for Polyatomic Ions
1. elements in the same Group form similar
polyatomic ions
–
same number of O’s and same charge
ClO3- = chlorate (-1 charge)
BrO3- = bromate (-1 charge)
2. if the polyatomic ion starts with H, the name adds
hydrogen- prefix before name and add 1 to the
charge
CO32- = carbonate \ HCO3- = hydrogen carbonate
42
2-42
Patterns for Polyatomic Ions
-ate ion
– chlorate = ClO3-
• -ate ion + 1 O  same charge, per- prefix
– perchlorate = ClO4-
• -ate ion – 1 O  same charge, -ite suffix
– chlorite = ClO2-
• -ate ion – 2 O  same charge, hypo- prefix,
-ite suffix
– hypochlorite = ClO43
2-43
Polyatomic Anions:
-ite, hypo- -ite, (-ate), per- -ate
ClOhypochlorite
NO2-
PO33-
SO32-
ClO2-
nitrite
phosphite
sulfite
chlorite
NO3-
PO43-
SO42-
ClO3-
nitrate
phosphate
sulfate
chlorate
ClO4perchlorate44
2-44
Polyatomic Ions to Remember
Name
Formula
Name
Formula
acetate
C2H3O2–
hypochlorite
ClO–
carbonate
CO32–
chlorite
ClO2–
chlorate
ClO3–
perchlorate
ClO4–
sulfate
SO42–
HSO4–
hydrogen carbonate
(aka Bicarbonate)
HCO3
hydroxide
OH–
nitrate
NO3–
nitrite
NO2–
Hydrogen sulfate
(aka Bisulfate)
permanganate
MnO4–
sulfite
SO32–
chromate
CrO42–
Hydrogen sulfite
(aka Bisulfite)
HSO3–
cyanide
CN–
dichromate
Cr2O7
ammonium
NH4+
–
2–
45
2-45
Practice: Naming Ionic compounds
•
•
•
•
•
•
•
Hg2SO4
CuClO3
Zn(NO3)2
FeCO3
Sn(SO3)2
CoPO4
Al(ClO4)3
46
2-46
Naming Acids
All names have acid at end
• Binary Acids (HnX) = hydro prefix + stem of the
name of the nonmetal + ic suffix
Example: HCl (Hydrochloric acid)
• Oxyacids (HnXOm : H2CO3 , H2SO4)
– if polyatomic ion ends in –ate = name of polyatomic
ion with –ic suffix :
H2SO4 (Sulfuric acid); H2CO3 (Carbonic acid);
HNO3 (Nitric acid);
H3PO4 (Phosphoric acid)
– if polyatomic ion ends in –ite = name of polyatomic ion
with –ous suffix
47
2-47
Naming Binary Acids – HF
1. First of all, it is binary acid HX
2. Identify the anion
F  F-, fluoride because Group 7A
2. Name the anion with an –ic suffix
F- = fluoride  fluoric
3. Add a hydro- prefix to the anion name
hydrofluoric
4. Add the word acid to the end
hydrofluoric acid
48
2-48
Naming Oxyacids: H2SO4
1. Identify the anion
SO4 = SO42- = sulfate
2. If the anion has –ate suffix, change it to –ic. If the anion
has –ite suffix, change it to -ous
SO42- = sulfate  sulfuric
3. Write the name of the anion followed by the word acid
sulfuric acid
(kind of an exception, to make it sound nicer!)
49
2-49
•
•
•
•
•
•
•
Practice: Naming Acids
first: what is the anion?
HNO33
HNO
HClO33
HClO
HBr
HBr
H22CO
H
CO33
H22SO
H
SO33
H33PO
H
PO44
HClO44
HClO
nitrate
chlorate
bromide
carbonate
sulfite
phosphate
perchlorate
50
2-50
Review: Naming Compounds
•
•
•
•
•
•
•
•
CuSO3
AgClO
N2O5
H2S
FeI2
Sn(NO3)4
Ba3(PO4)2
(NH4)2S
1. Common exceptions?
H2O, NH3, CH4, C12H22O11
2. Identify as Molecular or
Ionic?
3. Identify
•
•
Binary molecular
Type I or II metal ion
51
2-51
Review: Naming Compounds
•
•
•
•
•
•
•
•
CuSO3
AgClO
N2O5
H2S
FeI2
Sn(NO3)4
Ba3(PO4)2
(NH4)2S
copper(II) sulfite
silver hypochlorite
dinitrogen pentoxide
hydrosulfuric acid
iron(II) iodide
tin(IV) nitrate
barium phosphate
ammonium sulfide
52
2-52
Write Chemical Formula using the
charge of known ions
• Example: Compound between Ca2+ and PO43, the number of ions of each needs to be 3 and
2, so that the combined charge
= 3 x (+2) + 2 x (-3) = 0
Therefore the formula for the compound is
Ca3(PO4)2
53
2-53
Write Chemical Formula using
the charge of known ions
“CrissCross-Simplify”:
• The charge of an ion turns into the subscript (the
number) of the counterpart ion
Pb4+ O2-
 Pb2O4
• Since the subscripts in an ionic compound represents
the RATIO among the ions, the subscripts need to be
simplified when there is common denominator
Pb2O4  PbO2
54
2-54
Writing formulae
•
•
•
•
•
•
•
•
copper(II) sulfate
aluminum perchlorate
hydroiodic acid
iron(III) bromide
Diphosphorus pentoxide
lead(IV) nitride
zinc carbonate
helium gas
55
2-55
Tools of the Laboratory
Basic Separation Techniques
Filtration: Separates components of a mixture based upon
differences in particle size. Filtration usually involves separating
a precipitate from solution.
Crystallization: Separation is based upon differences in solubility
of the components in a mixture.
Distillation: Separation is based upon differences in volatility.
Extraction: Separation is based upon differences in solubility in
different solvents (major material).
Chromatography: Separation is based upon differences in
solubility in a solvent versus a stationary phase.
2-56
Tools of the Laboratory
Distillation: Separation through Difference in Boiling Temperatures
2-57
Tools of the Laboratory
Procedure for column chromatography: Separation through Adsorption
2-58
Tools of the Laboratory
Principle of gas-liquid chromatography (GLC).
2-59