Download Atoms, Molecules and Ions

Atomic Theory &
Chemical
Nomenclature
Chapter 2
HW:16,17, 24, 30,34,35,43,57,59
10th edition: 61,66,68,71
11th edition: 65,70,72,77
Early Evidence for Atomic Theory
• Democritus – matter is made of separate, discrete
particles (~ 450 BC)
• Total mass of a reaction stays the same – Law of
Conservation of Mass (1789 by Antoine Lavoisier)
• When elements combined to make compounds, they
always occurred in specific ratios – Law of definite
proportions (1799 by Joseph Proust)
– Water always decomposes in a ratio of 8 g O2 to 1 g H2.
• Pollen particles found to exhibit apparent
random motion (1827 by Robert Brown)
– Einstein explained the phenomenon in 1905 as
being the result of microscopic collisions of
atoms (water molecules).
John Dalton’s Atomic Theory
(A New System of Chemical
Philosophy: 1808)
1. Elements are composed of small
particles called atoms.
Atomos - uncuttable
2. All atoms of a given element are
identical. The atoms of one
element are different from the
atoms of all other elements (size,
mass, properties).
3
Dalton’s Atomic Theory (1808)
3. Compounds are composed of atoms of more
than one element. In any compound, the ratio of
the numbers of atoms is fixed.
H2O, NH3, MgCl2, CO2
4. A chemical reaction involves only the
separation, combination, or rearrangement of
atoms; it does not result in their creation or
destruction.
2Fe + 4H2O → 2H2 + 2Fe(OH)2
4
Dalton’s Atomic Theory # 3
Supported Proust’s Law of Definite Proportions
Dalton’s Atomic Theory # 4
16 X
+
8Y
8 X2Y
Supported Lavoisier’s Law of Conservation of Mass
Cathode Ray Tube: Thomson; e- Charge/Mass
Magnetic field
(N/S) result
Neither or
Both fields
+
-
Electric field
(+/-) result
Cathode Ray
Tube Video
J.J. Thomson (1897), measured charge/mass of e7
e- = -1.76 x 108 C/g (1906 Nobel Prize in Physics)
Cathode Ray in a discharge tube
The ray is invisible, but the fluorescence of a zinc
sulfide coating on the glass causes it to appear green.
Cathode rays are attracted to positive charges.
They consist of negatively charged electrons
8
Millikan’s Oil Drop Experiment
Measured mass of e(1923 Nobel Prize in Physics)
e- charge = -1.60 x 10-19 C
Thomson’s charge/mass of e- = -1.76 x 108 C/g
Millikan’s Experiment
e- mass = 9.10 x 10-28 g
9
Thomson’s Model (1904): Plum-pudding model
- Knew atoms were neutral
- But they had electrons (e-)
- Theorized an evenly
distributed (+) charge “cloud”
10
Radioactivity
• Recent discovery of X-rays (1895)
• Many substances were discovered to also
give off high energy rays in late 1890’s
• Spontaneous emission of particles and/or
radiation by elements via decay breakdown
Alpha particle: +2
• Term coined by Marie Curie
Gamma radiation: 0
Beta particle: -1
(uranium compound)
Ernest Rutherford’s Gold Foil Experiment
(with Hans Geiger and student Ernest Marsden)
(Rutherford had already won 1908 Nobel Prize in Chemistry for work in radioactivity)
 particle velocity ~ 1.4 x 107 m/s
(~5% speed of light)
1 out of every 20K deflected at > 90
1. atoms positive charge is concentrated in the nucleus
2. proton (p) has opposite (+) charge of electron (-)
3. mass of p is 1840 x mass of e- (1.67 x 10-24 g)
12
atomic radius
~ 100 pm = 1 x 10-10 m
Rutherford’s Model
of the Atom
nucleus radius
~ 5 x 10-3 pm = 5 x 10-15 m
20,000X smaller
“If the atom is the Houston Astrodome,
then the nucleus is a marble on the 50yard line.”
13
James Chadwick’s Experiment (Neutron) (1932)
(1935 Nobel Prize in Physics)
Rutherford’s proposal accounted for only the mass of protons
H atoms: 1 p+; He atoms: 2 p+
mass He/mass H should then be = 2/1
Actual measured ratio = 4/1
Chadwick subjected 9Be sheet to alpha radiation
 + 9Be
1n
+ 12C + energy
neutron (n) is neutral (charge = 0)
n mass ~ p mass = 1.67 x 10-24 g
mass p ≈ mass n ≈ 1840 x mass e-
15
Atomic Number, Mass Number, and Isotopes
Atomic number (Z) = number of protons in nucleus
Mass number (A) = number of protons + number of neutrons
= atomic number (Z) + number of neutrons
Mass Number
A
ZX
Atomic Number
1
1H
235
92
2
1H
U
(D)
238
92
*Position may change
Element Symbol
3
1H
U
(T)
Atomic # is always
a whole number
Ions and Isotopes
• Ion – Charged particle (atoms), various e– Element with unequal numbers of electrons and protons
– Atoms can gain/lose e- to form ions (never P+’s)
Mg → Mg+2 + 2eN + 3e- → N-3
• Isotopes – chemically similar atoms, various N0
– Same # of P+, unequal numbers of neutrons
– Atoms can vary neutrons to form isotopes (not P+’s)
The Isotopes of Hydrogen
99.985%
0.015%
Trace;
radioactive
Natural Abundance
18
Practice: How many P+, N0, e-?
Ne
P+
10
N0
20-10 = 10
e10
Ca
20
20
40-20 = 20
20
* 30
65-30 = 35
30 – 2e- = 28
20
10
40
65
14
+2
Zn
-3
N
*Carbon-14
7
6
14 - 7 = 7
14 - 6 = 8
*the atom is neutral,
# e- = # P+
7 + 3e- = 10
6
the ion is positive
= it has lost ethe ion is negative
= it has gained e-
*If we know the element we can find atomic # on Periodic Table
*if only one number is listed, it is always referring to the atomic mass
Practice: How many P+, N0, e-?
P+
N0
e-
P
15
15
14
15
59
Ni+3
28
31
25
I-
53
70
54
8
9
8
34
46
36
29
123
Oxygen-17
Se-2
80
Group
Period
Noble Gas
Halogen
Transition Metals
Alkali Metal
Alkaline Earth Metal
The Modern Periodic Table
21
Chemistry In Action
Natural abundance of elements in Earth’s crust:
Natural abundance of elements in human body:
22
A molecule is an aggregate of 2 or more atoms in
a definite arrangement held together by
chemical forces.
Composed of only non-metals
H2
H2O
NH3
CH4
NaCl ?
*Note: If it contains a metal it is not a molecule
Ionic compound
A molecule is an aggregate of 2 or more atoms in a
definite arrangement held together by chemical forces.
The 7 diatomic elements shown in blue are always found bonded to
each other (or another element).
*Know these
diatomic elements
A polyatomic molecule contains more than 2 atoms:
H2O, NH3, CH4, P4, P8
Allotropes – multiple
forms of an element
Representation: Formulas and Models
25
Example
2.2
Write the molecular formula of methanol, an organic solvent and
antifreeze, from its ball-and-stick model, shown below.
There are four H atoms, one C
atom, and one O atom. Therefore,
the molecular formula is CH4O.
However, the standard way of writing the molecular formula
for methanol is CH3OH because it shows how the atoms are
joined in the molecule.
A molecular formula shows the exact
number of atoms of each element in the
smallest unit of a substance.
An empirical formula shows the simplest wholenumber ratio of the atoms in a substance.
molecular
empirical
N2H4
NH2
C6H14N2O2
O3
H2O
C3H7NO
O
H2O
Empirical data: knowledge
based on observation or
experience rather than
theory or pure logic.
These ratios were first
experimentally observed
from decomposition of
substances
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Example
2.3 Empirical formulas
Write the empirical formulas for the following molecules:
(a) acetylene (C2H2), which is used in welding torches
There are two carbon atoms and two hydrogen atoms in acetylene.
Dividing the subscripts by 2, we obtain the empirical formula CH.
(b) glucose (C6H12O6), a substance known as blood sugar
In glucose there are 6 carbon atoms, 12 hydrogen atoms, and 6 oxygen atoms.
Dividing the subscripts by 6, we obtain the empirical formula CH2O.
(c) nitrous oxide (N2O), used as an anesthetic gas (“laughing
gas”) and as an aerosol propellant for whipped creams.
Because the subscripts in N2O are already the smallest possible whole numbers,
the empirical formula for nitrous oxide is the same as its molecular formula.
A monatomic ion contains only one atom:
Na+, Cl-, Ca2+, O2-, N3-, Fe2+, Fe3+
Some elements can have multiple charge states (oxidation
states), especially transition metals. Their charges can also
be represented as Roman Numerals (Iron II vs Iron III)
A polyatomic ion contains more than one atom:
They are molecules that have a net charge
Hydroxide: OH-1
Carbonate: CO3-2
Sulfate: SO4-2
Phosphate: PO4-3
Nitrate: NO3-1
Chlorate: ClO3-1
Ammonium: NH4+
Memorize these
common seven
Nitrate (NO3-1)
+1
Common Ions Shown on the Periodic Table
“Memorize these”
+2
Transition metals can have
Various charge states
+3
0
-3 -2 -1
Charges only if found in an ionic compound,
But NOT if neutral or found in a molecule
Group 4A doesn’t often form ionic compounds
30
An ion is an atom, or group of atoms, that has a
net positive or negative charge.
cation – ion with a positive (+) charge
If a neutral atom loses one or more electrons
it becomes a cation (oxidation).
Na
11 protons
11 electrons
Na+
11 protons
10 electrons
anion – ion with a negative (-) charge
If a neutral atom gains one or more electrons
it becomes an anion (reduction).
Cl
17 protons
17 electrons
Cl-
17 protons
18 electrons
Ionic compounds form from alternating
cations(+) and anions(-) (Metal+ and nonmetal-)
The sum of the charges on the cation(s) and anion(s)
in each formula unit must equal zero.
MgO, CaCl2, CoF2, K2S
The formula is the same
Mg2O2
as the empirical formula.
Is NOT A MOLECULE, no fixed
quantity of atoms
32
Ionic Compounds form a Crystal Lattice
• In NaCl, there are 6 Cl- interacting with 1 Na+.
– Each Cl- is interacting with 6 Na+.
• Therefore, they are in a 1:1 ratio.
• Na1Cl1 gives the element ratio, but is not a molecule
NaCl is the formula unit
Cl-
Cl-
ClNa+
Cl-
Cl-
Cl-
Crystal Lattice
Molecule or Ionic Compound?
• C6H12O6 • Molecule
*Remember
• NaCl
• Ionic Compound
Molecules are made of
only non-metals
• CO2
• Molecule
Ionic compounds have a
metal and non-metal
• O2
• Molecule
• LiOH
• Ionic Compound
• Fe
• Neither, just an atom
• Fe(NO3)2 • Ionic Compound
“Compound”
• Term only refers to a chemical that consists of
multiple types of elements.
• Ionic compounds are always compounds:
(Metal/Non-metal) NaCl, FeBr2, Al(NO3)3
• Molecules are normally compounds, but not always.
• Compound: H2O, SF6, CO2, C6H12O6
• Non-compound: O2, S8, Br2, P4
Polyatomic ions
• Covalently bonded molecules that possess a net
charge
• They can then participate in ionic bonds, but act
as single units
–
–
–
–
–
–
–
Ammonium: NH4+
Carbonate: CO3-2
Phosphate: PO4-3
Sulfate: SO4-2
Hydroxide: OHNitrate: NO3Chlorate: ClO3-1
Most common elements
in Ionic Compounds
The most reactive metals (green) and the
most reactive nonmetals (blue) combine to
form ionic compounds.
2Na (s) + Cl2 (g) → 2NaCl (s) + Energy!
Example
2.4
Write the formula of Calcium carbonate (limestone),
containing the Ca2+ and CO32− ions.
Solution To satisfy electrical
neutrality, the following
relationship must hold:
(+2)x + (−2)y = 0
The simplest ratio is 1:1 giving the formula unit: CaCO3
Check The subscripts are reduced to the smallest whole- number ratio
of the atoms because the chemical formula of an ionic compound is
usually its empirical formula.
Formulas of Ionic Compounds: Must equal 0
2 x +3 = +6
Al3+
3 x -2 = -6
Al2O3
1 x +2 = +2
Mg+2 & N-3
O21 x -2 = -2
CaS
Ca2+
2 x +1 = +2
Na+
S21 x -2 = -2
Na2CO3
“Crisscross and reduce”
CO3239
Ionic bond formation practice
1st determine ionized form
•
•
•
•
•
•
•
Mg and Br
Al and OHRb and N
Sr and PO4-3
Co (II) and O
Ca and CO3-2
NH4+ and O
2nd Cross charges
and reduce
Mg+2 Br-1
Al+3 OHRb+ N-3
Sr+2 PO4-3
MgBr2
*Al(OH)3
Rb3N
*Sr3(PO4)2
Co+2 O-2
CoO
Ca+2 CO3-2
CaCO3
NH4+ O-2
*(NH4)2O
*Polyatomic ions are in parentheses if needing multiple charges
Chemical Nomenclature
Naming Ionic Compounds
• Ionic Compounds
– Often a metal + nonmetal (or metal + polyatomic ion)
– Name metal first (left unchanged)
– Then Anion (nonmetal), but add “-ide” to element name
BaCl2
barium chloride
K2O
potassium oxide
Mg(OH)2
magnesium hydroxide
KNO3
potassium nitrate
If anion is a polyatomic ion, simply name it
42
• Transition metals in ionic compounds
– Must indicate charge on metal using Roman numerals
FeCl2
2 Cl- -2 so Fe is +2
Iron (II) chloride
“Ferrous chloride”
FeCl3
3 Cl- -3 so Fe is +3
Iron (III) chloride
“Ferric chloride”
Cr2S3
3 S-2 -6 so Cr is +3 (6/2)
Chromium (III) sulfide
Tungsten Carbide (WC) drill bits
*Miner Carbide head lamps (CaC2)
3x stiffer than steel
Reacts with H2O produce Acetylene
+ charge
- charge
45
Example
2.5 Formulas to Names
Name the following compounds:
(a) Cu(NO3)2
Copper (II) nitrate
(b) K3PO4
Potassium phosphate.
(c) (NH4)2CO3
Ammonium carbonate
(d) Ca3P2
Calcium Phosphide
(e) Cr(OH)3
Chromium (III) Hydroxide
(f) NiS
Nickel (II) Sulfide
Example
2.5 Solution
(a)The nitrate ion (NO3-1) bears one negative charge, so the copper ion
must have two positive charges. Copper (II) nitrate.
(b)The cation is K+ and the anion is PO4-3 (phosphate). Because
potassium only forms one type of ion (K+), there is no need to use
potassium (I) in the name. The compound is potassium phosphate.
(c) The cation is NH4+1 (ammonium ion) and the anion is CO3-2 . The
compound is ammonium carbonate.
(d) Calcium Phosphide, Ca is not a transition metal so we do not need
to specify the charge with Roman numerals.
(e) Chromium (III) Hydroxide
Example
2.6 Name to Formulas
Write chemical formulas for the following compounds:
(a) chromium (III) nitrate
Cr(NO3)3
(b) cesium sulfide
Cs2S
(c) calcium phosphate
Ca3(PO4)2
(d) lead (II) nitrate
(e) Calcium Carbonate
Pb(NO3)2
CaCO3
Example
2.6 Solutions
(a) We can cross the +3 charge of Chromium (III) with the -1
charge of Nitrate to yield Cr(NO3)3 using parentheses to
signify we need 3 x NO3-1 to 1 x Cr+3
(b) Each sulfide ion bears two negative charges, and each cesium
ion bears one positive charge. Therefore, the formula is Cs2S.
(c) Each calcium ion (Ca2+) bears two positive charges, and each
phosphate ion (PO4-3) bears three negative charges.
3(+2) + 2(−3) = 0
The formula is Ca3(PO4)2
(d) Pb(NO3)2
(e) Crossing the charges would give Ca2(CO3)2, but this reduces
to CaCO3
• Molecular compounds
− Nonmetals or nonmetals + metalloids
− Some have common names (small)
− H2O, NH3, CH4 , PH3, O3
− Element furthest to the left in a
period and closest to the bottom of a
group on periodic table is placed first
in formula
− If more than one compound can be
formed from the same elements, use
prefixes to indicate number of each
kind of atom
− Last element name ends in -ide
50
Molecules can arrange the same
elements in various arrangements.
Necessitates the usage of prefixes to distinguish
NO: Nitrogen Monoxide
Dinitrogen trioxide: N2O3
Dinitrogen pentaoxide
N2O5
Nitrogen dioxide
NO2
Dinitrogen monoxide
N2O
Molecular Compounds
P4O10
tetraphosphorus decoxide
NF3
nitrogen trifluoride
SF6
sulfur hexafluoride
N2Cl4
dinitrogen tetrachloride
NO2
nitrogen dioxide
*The prefix mono- is not normally used if on first element
52
53
Example
2.8 Molecule Name to Formula
Write chemical formulas for the following
molecular compounds:
(a) carbon disulfide
(a) Because there are two sulfur
atoms and one carbon atom present,
the formula is CS2.
(b) There are two silicon atoms and
(b) disilicon hexabromide six bromine atoms present, so the
formula is Si2Br6.
(c) B2O3
(c) Diboron Trioxide
55
Hydrates are ionic compounds that have a specific
number of water molecules attached to them.
BaCl2•2H2O
barium chloride dihydrate
LiCl•H2O
lithium chloride monohydrate
MgSO4•7H2O
magnesium sulfate heptahydrate
Sr(NO3)2•4H2O
strontium nitrate tetrahydrate
CuSO4•5H2O
Copper (II) sulfate
pentahydrate
Copper (II) sulfate
“anhydrous”
CuSO4
56
An acid can be defined as a substance that yields
hydrogen ions (H+) when dissolved in water.
For example: HCl gas and HCl in water
• Pure substance, hydrogen chloride
• Dissolved in water (H3O+ and Cl−),
hydrochloric acid
57
Simple acids:
Hydrogen bonded to single element (often a Halogen)
Specifically does not have oxygen present
58
An oxoacid is an acid that contains
hydrogen, oxygen, and another element.
HNO3
nitric acid
H2CO3
carbonic acid
H3PO4
phosphoric acid
59
Naming Oxoacids and Oxoanions
60
Prefix and Suffix change as number of oxygen
atoms change (oxidation state)
61
The rules for naming oxoanions, anions of oxoacids
1. When all the H ions are removed from the “-ic” acid,
the anion’s name ends with “-ate.”
Nitric Acid (HNO3) vs Nitrate (NO3-1)
2. When all the H ions are removed from the “-ous” acid,
the anion’s name ends with “-ite.”
Nitrous Acid (HNO2) vs Nitrite (NO2-1)
3. The names of anions must indicate the number of H
ions present.
–H PO - dihydrogen phosphate
2
4
–HPO4 2- hydrogen phosphate
–PO43- phosphate
62
Example
2.9
Name the following oxoacid and oxoanion:
(a)HNO3 and NO3-1
(b)H2CO3 and CO3-2
(c)H3PO3 and PO3-3
(d)HClO4 and ClO4-1
Nitric Acid and Nitrate
Carbonic Acid and Carbonate
Our reference acid is phosphoric acid (H3PO4).
Because H3PO3 has one less O atom, it is called
phosphorous acid. PO3-3 has one less O atom than
phosphate so it is called phosphite.
Our reference acid is Chloric Acid (HClO3).
Because HClO4 has one more O atom, it is called
perchloric acid. ClO4-1 is called perchlorate.
A base can be defined as a substance that yields
hydroxide ions (OH-) when dissolved in water.
As an ionic compound, it follows the same nomenclature,
the anion is always hydroxide
NaOH
KOH
Ba(OH)2
sodium hydroxide
potassium hydroxide
barium hydroxide
64
Organic Chemistry
Carbon Molecules
If the molecule has Carbon
and Hydrogen it follows
organic nomenclature instead
Single bonds: end with -ane
Double bonds: end with -ene
Triple bonds: end with -yne
Ethane
Ethene
Ethyne
“Acetylene”
Organic chemicals are often named by their
Functional Groups that have similar atomic
arrangements for similar properties.
H
H
Alcohol
C
OH
H
methanol
H
H
Amine
C
NH2
H
methylamine
Carboxylic Acid
H O
H
C
C
OH
H
acetic acid
66
Organic Structure Representation
• Octonol: C8H18O or
C8H17OH
• CH3CH2CH2CH2CH2CH2CH2CH2OH
• CH3(CH2)7OH
• Condensed
• Space-filled