Download ATOMS - Mr. Deets

Unit 3
Language of Chemistry
Part 1
Zumdahl: Chapter 4
Holt: Chapter 3
ATOMS: The Building
Blocks of Matter
Objectives
1.
2.
3.
4.
5.
Law of conservation of mass
Law of definite proportions
Law of multiple proportions
Dalton’s Atomic Theory
How Dalton’s Atomic Theory relates to 1,
2, & 3
Atomic Theory Foundations
Law of Conservation of Mass – mass is
neither created or destroyed during a
chemical or physical change
Law of Definite Proportions – a compound
contains the same proportions by mass
regardless of the size of the sample
Example:
NaCl – always 39.34% Na & 60.66% Cl
Atomic Theory Foundations
Law of Multiple Proportions – if two or more
different compounds are composed of the
same two elements, then the ratio of the
masses of the second element combined
with a certain mass of the first element is
always a ratio of small whole numbers
Example
CO2 and CO : ratio of oxygen is always 2:1
JJ Thomson
Showed atoms could emit negative particles
“plum pudding” model
Electrons were embedded in a positively
charged spherical cloud
Rutherford
• Shoots alpha particles (Helium atoms) at
gold foil
• Expected to pass right through
• Particles are deflected
• Leads to idea of a dense positively
charged center with e- orbiting around it
Ernest Rutherford
• Gold foil experiment
Dalton’s Atomic theory
1. All matter is composed of atoms
2. Atoms of an element have the same size,
mass and properties; atoms of a different
element have different sizes, masses and
properties
3. Atoms cannot be divided, created or destroyed
4. Atoms of different elements combine in simple
whole number ratios
5. Chemical reactions combine, separate, or
rearrange atoms
Modern Atomic Theory
• Atoms can be divided
• Atoms of the same element can have
different masses
• All else remains the same
Structure of the Atom
Objectives
Discovery of the Electron
Rutherford’s Experiments
Protons, Neutrons, Electrons
Atomic Structure
• Electron = no mass; negative charge
• Proton mass = hydrogen atom; positive
• Neutron mass = hydrogen atom; no charge
• Dalton’s Model
• JJ Thompson’s Plum Pudding Model
The Electron
• Mass of 9.109 x 10
-31
• Negative charge
kg
The Proton
• Mass = 1.673 x 10
-27
• Positive charge
kg
The Neutron
• Mass = 1.675 x 10
-27
• No charge
kg
Comparing Theories
Dalton
See notes
Thompson
Plum pudding
model
Rutherford
Concept of the
nucleus
Electrons
scattered thru
positively
charged cloud
Positively
charged
Solid sphere
Isotopes
II.A.2(c) – compare the
characteristics of
isotopes of the same
element
IV.B.2(d) – calculate the
weighted average atomic
mass of an element from
isotopic abundance,
given the mass of each
contributor
Isotopes
Def: atoms of the same element that have
different masses
Example: hydrogen
protium – 1 proton in nucleus
deuterium – 1 proton; 1 neutron
tritium – 1 proton; 2 neutrons
*Nuclide – general term for any isotope
Writing Isotopes
mass number Element Symbol
atomic number
Example – Uranium 235
235
92
U
Average Mass Number
Def: the weighted average of the atomic
masses of the naturally occurring isotopes
of an element
Like calculating a “weighted” grade
(decimal % of each isotope x mass of that
isotope)
Sample Calculation
Isotope
oxygen - 16
oxygen - 17
% abundance
99.762
0.038
Atomic Mass
15.994915
16.999131
oxygen - 18
0.200
17.999160
The
Periodic
Table
IV.B.2(c) – use the periodic table
to determine the atomic number,
atomic mass, mass number, and
number of protons, electrons and
neutrons in isotopes of elements
Atomic Number
• Atomic Number
– # of protons in nucleus
• Element Symbol
• Element Name
• Atomic Weight
• Electron
Configuration
3
Li
Lithium
6.941
[He]2s
1
Calculating
Protons,
Electrons &
Neutrons
IV.B.2(c) – use the
periodic table to
determine the atomic
number, atomic mass,
mass number, and
number of protons,
electrons and neutrons
in isotopes of elements
Mass Number
Def: the number of protons and neutrons in
the nucleus of an isotope
mass # - atomic # = # of neutrons
Example – oxygen
Mass # (16) – atomic # (8) = # of neutrons (8)
Formulas
II.A.2(a) – use the IUPAC symbols of the
most commonly referenced elements
III.A.1(a) – distinguish between chemical
symbols, empirical formulas, molecular
formulas and structural formulas
III.A.1(b) – interpret the information
conveyed by chemical formulas for
numbers of atoms of each element
represented
III.A.(d) – provide the interconversion of
molecular formulas, structural formulas,
and names, including common binary
and ternary acids
Chemical Symbols
Molecular Formulas
Empirical Formulas
Structural Formulas
Percent
Composition
III.A.1(c) – calculate the
percent composition of a
substance given its
formula or masses of
each component element
in a sample
% Composition
Sample Problems
Zumdahl – Ch 5
Holt – Ch 7
III.A.1(d) - provide the
interconversion of
molecular formulas,
structural formulas,
and names, including
common binary and
ternary acids
Oxidation Numbers
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Pure element have an oxidation number of zero
Fluorine has an oxidation of -1 in all
compounds
Oxygen has an oxidation number of -2 except
when in a peroxide when its oxidation number
is -1
Hydrogen has an oxidation number of +1 except
when bonded to a metal
The algebraic sum of the oxidation numbers of
all atoms in a neutral compound is zero
The algebraic sum of the oxidation numbers of
all atoms in a polyatomic ion is equal to the
charge of the ion
These rules apply to covalently bonded atoms
Oxidation Numbers Cont’d
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Group 1 = +1
Group 2 = +2
Transition Metals = varies
Group 15 = varies (-3 usually)
Group 16 = varies (-2 usually)
Group 17 = varies (-1 usually)
General Rules for Naming
Chemical Compounds
• Cation (pronounced cat-ion)
– is named first
– has a positive charge/oxidation number
– name does NOT change
• Anion (pronounced an-ion)
– is named second
– has a negative charge/oxidation number
– ending is “-ide”
Binary Ionic Compounds
Type I
• Metal is from Group 1 or 2
Examples:
NaCl – sodium chloride
KF – potassium fluoride
CaBr2 – calcium bromide
Li2O – lithium oxide
Binary Ionic Compounds Type II
• Metal is a transition metal
• Many metals form more than one type of
positive ion
–
–
–
–
Iron: Fe2+ or Fe3+
Copper: Cu1+ or Cu2+
Lead: Pb2+ or Pb4+
Tin: Sn2+ or Sn4+
• Roman numeral indicates charge on the metal
ion (NOT the number of atoms)
• Look at 2nd element to determine charge on
transition metal
Examples of Type II
EX: FeCl2 – iron(II) chloride
FeCl3 – iron(III) chloride
CuO – copper(II) oxide
Cu2O – copper(I) oxide
SnF4 – tin(IV) fluoride
SnO2 – tin(IV) oxide
SnO – tin(II) oxide
**NOTE: sum of oxidation numbers = 0
Exceptions to the Rule
Aluminum – always Al
Cadmium – always Cd2+
Zinc – always Zn2+
Silver – always Ag+
3+
No Roman numerals needed
Binary Covalent Compounds or
Binary Molecular Compounds
• Formed between two non-metals
1.first element name doesn’t change
2.Second element ends with “-ide”
3.Don’t use “mono-” on first name
4.Use prefixes to indicate number of atoms
of element
Prefixes
Mono– 1
Hepta- 7
Di– 2
Octa- 8
Tri- 3
Nona- 9
Tetra- 4
Deca- 10
Penta- 5
Undeca – 11
Hexa- 6
Dodeca - 12
Examples of Binary Covalent
Compounds
CO – carbon monoxide
CO2 – carbon dioxide
CCl4 – carbon tetrachloride
SiO2 – silicon dioxide
SeBr2 – selenium dibromide
BF3 – boron trifluoride
P2O4 – diphosphorous tetroxide
P4O10 – tetraphosphorous decoxide
Polyatomic Ions
• Group of atoms with a shared
charge
III.A.1(c) – use the names,
formulas, and charges of
commonly referenced polyatomic
ionsl
Polyatomic Rule #1
• Common form ends with –ate
Example: nitrate, chlorate, sulfate
Polyatomic Rule #2
• If oxygen decreases by one, O-1,
changes ending to “-ite”
Example: nitrite, chlorite, sulfite
Polyatomic Rule #3
• If oxygen decreases by two, O-2, add
prefix “hypo-”
• Keep ending “-ite”
Example: hypochlorite
Polyatomic Rule #4
• If oxygen increases by one, O+1, add
prefix “per-”
• Keep ending “-ate”
Examples: perchlorate, permanganate
Polyatomic Ions
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Chlorate
Bromate
Iodate
Nitrate
Permanganate
Carbonate
Silicate
Selenate
Phosphate
Arsenate
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ClO3BrO3IO3NO3MnO4CO32SiO32SeO32PO43AsO43-
More Polyatomic Ions
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Acetate
Hydroxide
Bicarbonate
Bisulfate
Ammonium
Thiocyanate
Thiosulfate
Oxalate
Chromate
Dichromate
C2H3O2OHHCO3HSO4NH4+
SCNS2O32C2O42CrO42Cr2O72-
Binary
Acids
III.A.1(d) - provide the
interconversion of
molecular formulas,
structural formulas,
and names, including
common binary and
ternary acids
Binary Acids
• Hydrogen takes place of metal ion
• Use prefix hydro- in place of hydrogen
• Use suffix -ic on the root
Examples –
HCl – hydrochloric acid
HF – hydroflouric acid
H2S – hydrosulfuric acid
Oxyacids
III.A.1(d) - provide the
interconversion of
molecular formulas,
structural formulas,
and names, including
common binary and
ternary acids
Oxyacids
• contain hydrogen, a non-metal, and oxygen
• usually contain 3 or 4 oxygen atoms
• add suffix -ic to the stem
Examples –
HBrO3 – bromic acid
H2CO3 – carbonic acid
HClO3 – chloric acid
HNO3 – nitric acid
H3PO4 – phosphoric acid
H2SO4 – sulfuric acid
Oxyacids Cont’d
• Common name ends with -ate
• One extra oxygen add per- to stem name
Example –
HClO3 – chloric acid
HClO4 – perchloric acid
• One less oxygen change -ic ending to -ous
Example –
HClO3 – chloric acid
HClO2 – chlorous acid
Oxyacids cont’d
• Two fewer oxygens change -ic ending to
-ous AND add prefix hypoExamples HClO3 – chloric acid
HClO – hypochlorous acid