PowerPoint to accompany
Chapter 1
Introduction: Matter,
Measurement and Molecules
Dr V Paideya
Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
CHEMISTRY
“study of matter & changes it undergoes”
“matter is anything that has mass and takes up space”
- study of physical & chemical properties of matter
- what changes occur in these properties, in the course of/as the result of a
chemical reaction, & how these changes may be observed
- why the reaction involved does (or doesn’t…) occur
be able to understand & explain such
macroscopic changes from an atomic/molecular
(submicroscopic) perspective
States (Phases) of Matter
- solid, H2O(s); liquid, H2O(l); gas, H2O(g)
- phase transitions occur @ specific P/T values,
governed by properties of atoms/molecules
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Matter
- atoms are building blocks of matter
- each element is made of same kind of atom/molecules
- compounds made of two or more different kinds of elements bonded together
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Pure Substances, Elements & Compounds
pure substance
-has distinct properties & unvarying/constant composition
eg. NaCl(s), H2O(l), HCN(g)
element
-substance that cannot be decomposed into simpler substances
eg. Cl2(g), Br2(l), I2(s); Ne(g), Hg(l), Au(s)
compound
-substance composed of 2 or more different elements
2 or more different kinds of atoms
eg. UF6(g), H2O(l), CaCO3(s)
Law of Constant Composition/Definite Proportions (Joseph Proust ca 1800)
“...elemental composition of pure substance is always the same…”
- different samples of pure compound have the same elemental
composition
- elements present in such samples have same proportion by mass
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Mixtures
- combination of 2 or more substances, in which each substance retains own
chemical identity & can thus be separated from each other
- 2 types:
heterogeneous:
- mixture of visibly different composition, properties or appearance
eg. sand in H2O(l) (s, l), sand & NaCl (s, s), petrol & H2O(l) (l, l)
homogeneous:
- mixture of visibly uniform composition, properties & appearance throughout
eg. NaCl(aq) (s,l), air (g,g), stainless steel (s,s), soda water (g,l)
Properties:
- physical: measurement without changing identity/composition eg. mass, , v
-chemical: must involve change in chemical identity eg. flammability, reactivity
- extensive: dependent on quantity of sample involved eg. mass, volume
- intensive: independent of quantity eg. , m.p., b.p.; useful for identification
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Classification of Matter
Figure 1.5
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Changes of Matter
Physical Changes
Changes in matter that do not change the
composition of a substance.
Changes of state, temperature, volume, etc.
Chemical Changes
Changes that result in new substances.
Combustion, oxidation, decomposition, etc.
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Chemical Reactions (Chemical Change)
In the course of a chemical reaction, the reacting
substances are converted to new substances.
Figure 1.6
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Compounds
Compounds can be broken down into more
elemental particles; for example, during the
electrolysis of water, the smaller particles
hydrogen gas and oxygen gas are created.
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Separation of Mixtures
1. Distillation
Separates a
homogeneous
mixture on the
basis of differences
in boiling point.
Figure 1.8
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Separation of Mixtures
2. Filtration
3. Chromatography
Separates substances on the basis of differences in
solubility in a solvent
Separates solid substances from liquids and solutions.
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The Scientific Method
A systematic approach to solving problems
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SI Units
Système International d’Unités
Uses a different base unit for each quantity
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Metric System
Prefixes convert the base units into units that
are appropriate for the item being measured.
Tera T 1012
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SI Units - Temperature
The Kelvin is the SI
unit of temperature.
It is based on the
properties of gases.
There are no negative
Kelvin temperatures.
K = C + 273.15
Figure 1.10
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Derived SI Units VolumeUnits of Measurement
Derived SI Units: Volume, V; E
-1 m10 d(eci)m 100 c(enti)m 1000 m(illi)
-1 L 1 dm3 (a cube 1 dm x 1 dm 1 x dm)
-1 mL 1 cm3(a cube 1 cm x 1 cm 1 x cm)
-1 kL 1 m3 (a cube 1 m x 1 m x 1 m)
ie. 1 m3 103 dm3 106 cm3 109 mm3
1 kL 103 L 106 mL 109 (micro)L
Derived SI Units: Density, or d; I
- physical property; units are usually
-g cm-3 (g/cm3, g mL-1, g/mL) for liquids
& solids
-g dm-3 (g/dm3, g L-1, g/L) for gases
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Derived SI Units Density
Density is a physical property of a substance
and is determined through the following
formula:
density =mass
volume
=mVor symbolically
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Uncertainty in Measurements
Different measuring devices have different
uses and different degrees of accuracy.
Figure 1.12
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Significant Figures
All digits of a measured quantity,
including the uncertain, are called
significant figures.
The greater the number of significant
figures, the greater the certainty of the
measurement.
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Significant Figures All nonzero digits are significant, e.g. 123.45
Zeros between two significant figures are
themselves significant, e.g. 103.405
Zeros at the beginning of a number are
never significant, e.g. 00123.45 = 123.45
Zeros at the end of a number are
significant if a decimal point is written in the
number, e.g. 123.450 has six significant figures but 123450 has
only five significant figures
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Significant Figures
When addition or subtraction is
performed, answers are rounded to the
least significant decimal place.
When multiplication or division is
performed, answers are rounded to the
number of digits that corresponds to the
least number of significant figures in any
of the numbers used in the calculation.
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Precision and Accuracy
Accuracy refers to the proximity of a measurement
to the true value of a quantity.
Precision refers to the proximity of several
measurements to each other.
Figure 1.15
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Significant Figures1. any figure that is not zero is significant:
snot/units 845 mL _____ s.f.1243.29 mg _____ s.f.
2. zeroes between non-zero figures are significant:1906 mL _____ s.f.40501.09 J _____ s.f.
3. exact (“counting”) numbers by definition have an number of s.f., so physical
constants defined to be exact numbers do so also...:1 atm 101.325 kPa 760
mmHg; 0 OC 32 OF 273.15 K all _____ s.f.
4. leading zeroes (to the left of the first non-zero figure) are not significant:
snot/units 0.008 kg _____ s.f.0.003798 L _____ s.f.
5. trailing zeroes (to the right of the last non-zero figure) are significant only if the
number has a d.p.: 300.0 mm _____ s.f.0.0300 mm _____ s.f.
6. in measurements without a d.p., the number of s.f. is ambiguous:1200 mm ?? either: i) use snot
OR ii) 1200.
Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
Using Significant Figures in Calculations
- all calculations governed by two fundamental rules
multiplication/division
- number of s.f. in final answer is the LEAST of numbers of s.f. in each of
original measurements
addition/subtraction
- number of d.p. in final answer is the LEAST of numbers of d.p. in each of
original measurements
Eg. 1
Calculate
i) volume, in mm3, of a box of length 6.741 cm, breadth 2.441 x 10-1 m, & height
4.2 mm i) 6.9 x 104 mm3;
ii) density () of a pure liquid, in g cm-3, if 103.67 g of it is needed to fill the box
completely ii) 1.5 g cm-3
Eg. 2
An empty container of mass 23.29 g has a mass of 86.1 g when filled with 0.5000
dm3 of a pure liquid. Determine the of this liquid in g cm-3. 0.126 g cm-3
Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
Early Atomic Theory
John Dalton 1803 - 1807
-each element is composed of very small, indestructable, particles called atoms*
-all atoms* of given element are physically & chemically identical to each other, but
atoms of a particular element are different from atoms of all other elements
-Law of Conservation of Mass
-atoms are neither created or destroyed in chemical reactions
- mass reactants present @ start = mass products formed @ completion*
-Law of Constant Composition
-different samples of a pure compound have the same elemental composition
-elements present in such samples have same proportion by mass
-Law of Multiple Proportions
-if 2 elements (C & O) can combine to form 2 or more different compounds (CO &
CO2), the different masses of one element (O) combining with a fixed mass of the
other (C) can be expressed as a simple integral ratio
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Atomic TheoryThe theory that atoms are the fundamental building blocks of matter came
into being during the period 1803 to 1807 in the work of John Dalton.Dalton’s Postulates Each element is composed of extremely small particles called atoms.
All atoms of a given element are identical to one another in mass and other properties, but the atoms of a particular element are different from the atoms of all other elements.
Atoms of an element are not changed into atoms of a different element by chemical reactions; atoms are neither created nor destroyed in chemical reactions. This is the basis of the law of conservation of mass (or law of conservation of matter) which states that the total mass of substances present at the end of a chemical process is the same as the mass of substances present before the process took place.
Compounds are formed when atoms of more than one element combine; a given compound always has the same relative number and kind of atoms. This is the basis of the law of constant composition (or law of definite proportions) which states that the relative numbers and kinds of atoms are constant, i.e. the elemental composition of a pure substance never varies.
Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia
The Law of Multiple Proportions
Was deduced by Dalton from the preceding four
postulates and states that:
If two elements A and B combine to form more
than one compound, the masses of B that can
combine with a given mass of A are in the ratio
of small whole numbers.
Examples
H2O consists of 2 hydrogens and 1 oxygen
H2O2 consists of 1 hydrogen and 1 oxygen
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The Discovery of Atomic Structure
Cathode Rays & Particles (JJ Thomson, 1897)
- electrical discharges from cathode originally thought to be new form of radiation
- showed that radiation emitted was
- independent of cathode material used
- deflected by magnetic/electric fields
- findings consistent with model in which “beam” /”rays” composed of negatively
charged “particles”(-) with charge/mass ratio of - 1.7588 x 108 C g-1, or
-5.6857 x 10-9 g C-1
Electron Charge & Mass (Robert Millikan, 1909)
- oil drop experiment
- (-) charge on oil drops found always to be a multiple of
minimum value of 1.6(02) x 10-19 C ie. 1.602 x 10-19 C must
be charge of single electron (e-)
-mass of single e- determined to be 9.109 x 10-28 g
- only 1/1836 of mass of an H atom
- first subatomic particle
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Radioactivity
The spontaneous emission of radiation by
an atom was first observed by Henri
Becquerel. It was also studied by Marie
and Pierre Curie.
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Radioactivity
Three types of radiation were discovered by
Ernest Rutherford
particles
particles
rays
Figure 1.21
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Discovery of the Nucleus
Ernest Rutherford shot particles at a thin sheet of gold foil and
observed the pattern of scatter of the particles.The Nuclear Atom
Some particles were deflected at large angles.
This led Rutherford to postulate that the
atom had a nucleus.
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Modern Atomic Structure
- more than 99.99 % of atom mass & entire Q+ centred in atomic nucleus,
where nucleons (protons, p+ (Rutherford, 1919) & neutrons, nO (Chadwick,
1932) are collectively bound together by strong nuclear force
- atomic nucleus surrounded by much larger atomic volume, containing as
many e- as p+, so atom is electrically neutral & held together by force of
Coulombic/ electrostatic attraction
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element symbol
(number of e- in neutral atom)
atomic number (Z)
number of p+
mass number (A)
number of p+ & nO
ZA E
Atomic (Z) & Mass (A) Numbers
- atoms of different elements have different numbers of p+ in their nuclei
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Isotopes
Atoms with identical atomic numbers (Z) but different
mass numbers (A), or atoms with the same number of
protons which differ only in the number of neutrons are
called isotopes.
Examples:
116C
126C
136C
146C
carbon-12 isotope
carbon-14 isotope
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Isotopes
- atoms of same element having different numbers of nO in their nuclei
- ie. same Z, different A, or same Z, different N
- chemical properties largely similar, but physical properties, & particularly the “nucular”
ones involving radioactive “nuculei”, can be very different
- each Mg ( ) atom is one of three naturally occurring isotopes
- 24Mg; 25Mg; 26Mg
Mg-24 Mg-25; Mg-26
-
-three isotopes of H individually named:
protium (1H, 1 p+/0 nO); deuterium (2H, 1 p+/1 nO); tritium (3H, 1 p+/2 nO)
- four isotopes of C:
11C; 12C; 13C (NMR probes & MRI scanners); 14C(“radiocarbon” dating)
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Isotopes
Eg. 3 Complete the table below:
Experimentally..
- High Resolution Mass Spectrometry (p. 28) used for very precise (4-6 d.p.; 7-10
s.f. in total) measurements of the masses of an element’s isotopes & their naturally
occurring abundances
do: Q 1.44, 1.48 - 1.50
Element name Symbol p+ No e- A
79197Au
Ba 138
143 92
Pb 126
Krypton36
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Atomic Mass
Atomic and molecular masses can be
measured with great accuracy with a mass
spectrometer.
Figure 1.23
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Average Atomic Mass(commonly called Atomic Mass) We use average masses in calculations, because we use
large amounts of atoms and molecules in the real world.
Average atomic mass is calculated from the fractional
abundance of each isotope and mass of that isotope.
For example, the average atomic mass of C -
made up mostly of 12C (98.93%) and 13C (1.07%) - is 12.01
u.
extremely small SI masses of individual atoms (~4 x 10-22 g) too awkward
for everyday usage, so masses expressed in unified atomic mass units
(amu, u):
1 amu (u) = 1.66054 x 10-24 g 1 g = 6.02214 x 1023 amu (u)
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Average Atomic Masses of Naturally Occurring Elements
-use average masses in “real world” calculations, as even smallest weighable sample
(~1 g 10-6 g) involves gobsmackingly large (~1015, or 10 quadrillion) numbers of
atoms
-no AAMs calculated as “weighted average” of an element’s isotopic masses (IMs) &
naturally occurring abundances
AAM= (IM x % ab/100) or (IM x fr ab)
Eg. 4
a)Naturally occurring Mg has three isotopes: 24Mg 78.99 %, 23.9850 u25Mg 10.00 %, 24.9858 u
Calculate its AAM. 26Mg 11.10 %, 25.9826 u
b) Naturally occurring Pb has four isotopes: 204Pb 1.40 %, 203.973037 206Pb 24.10 %,205.974455207Pb 22.10 %, 206.975885
Calculate its AAM. 208Pb 52.40 %, 207.976641
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Eg. 5
Chlorine has two naturally occurring stable isotopes: 35Cl 34.968853 u 37Cl 36.965803 u
If the (average) atomic mass of naturally occurring elemental Cl is 35.453 u, what
are the % abundances of the two isotopes?
http://www.sisweb.com/referenc/source/exactmaa.htm
do: antimony, chromium, & nickel AM from IM’s & abundances
copper, rubidium abundances from IM’s & AM*
The Periodic Table
- rapid post-Dalton growth in experiment-based chemical knowledge showed very
quickly that many elements could be grouped together on basis of
similarities in their physical & chemical properties
- arrangement of elements in order of Z showed that these similarities occurred
in repetitive/periodic patterns, & agreed so closely with experimentally acquired
data, that phys/chem properties of 2 “missing” elements were accurately predicted
before their being reported as formally discovered, &/or phys/chem properties
characterized
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Eg. 4
a) Naturally occurring Mg has three isotopes: 24Mg 78.90 %, 23.9850 u25Mg 10.00 %, 24.9858 u
Calculate its AM. 26Mg 11.10 %, 25.9826 u
AM = (IM x % ab/100)
= {[23.9850 x (78.90/100)] + [24.9858 x (10.00/100)] + [25.9826 x (11.10/100)] }
= 18.924 + 2.4986 + 2.8841
= 24.31 u (amu) [24.306]
b) Naturally occurring Pb has four:204Pb 1.40 %, 203.973037206Pb 24.10 %, 205.974455207Pb 22.10 %, 206.975885
Calculate its AM. 208Pb 52.40 %, 207.976641
AM = (IM x % ab/100)
= 2.856 + 49.640 + 45.742 + 108.98
= 207.2 u (amu)
[207.22]
back
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Eg. 5
Chlorine has two naturally occurring stable isotopes:35Cl 34.968853 u 37Cl 36.965803 u
If the (average) atomic mass of naturally occurring elemental Cl is 35.453 u, what are
the % abundances of the two isotopes?
Assume that the fractional abundance of 35Cl is z......then the fr ab of 37C is
1-z ( fr ab = 1)
AM = (IM x fr ab)
35.453 =[(34.968853) x z] + [(36.965803) x (1-z)]
=34.968853 z + 36.965803 - 36.965803 z
-1.5128 = -1.996950 z (35.453 - 36.965803 -1.5128)
z =0.7576 3 dp 6 dp
% ab 35Cl =75.76%
% ab 37Cl =24.24 %
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The Periodic Table
When one looks at the chemical properties of elements, one notices a
repeating pattern of reactivities.
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Periodic Table The rows are called periods.
The columns are called groups.
Elements in the same group have similar chemical
properties.
Nonmetals are on the right side of the periodic table (with
the exception of H).
Metalloids border the stair-step line (with the exception of Al
and Po).
Metals are on the left side of the chart.
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Groups
The above five groups are known by their names.
Table 1.7
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The Periodic TableMetals, Non-Metals, & Metalloids
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Molecules and Chemical Formulae
The subscript to the right of
the symbol of an element tells
the number of atoms of that
element in one molecule of
the compound.
Notice how the composition of
each compound is given by
its chemical formula.
Figure 1.29
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Diatomic Molecules
These seven elements occur naturally as molecules containing two atoms.
Figure 1.28
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Molecular CompoundsMolecular compounds are composed of molecules and almost
always contain only nonmetals.
Types of FormulaeEmpirical formulae give the lowest whole-number ratio of atoms
of each element in a compound, e.g. HO.
Molecular formulae give the exact number of atoms of each
element in a compound, e.g. H2O2.
Structural formulae show which atoms are attached to which
within the molecule, e.g. H-O-O-H.
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Picturing Molecules
Different
representations of
the methane (CH4)
molecule.
Figure 1.30
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Ions and Ionic Compounds
When atoms lose or gain electrons, they become ions.
Cations are positive and are formed by elements on the left side of the periodic chart.
Anions are negative and are formed by elements on the right side of the periodic chart.
Figure 1.31
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Ionic Compounds
Ionic compounds (such as NaCl) are
generally formed between metals and
nonmetals.
Figure 1.32
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Writing Formulae
Because compounds are electrically neutral, one
can determine the formula of a compound by:
writing the value of the charge on the cation
as the subscript on the anion.
writing the value of the charge on the anion
as the subscript on the cation.
Note: if the subscripts are not in the lowest
whole number ratio, simplify it, e.g. Ca2O2
would become CaO.
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Chemical NomenclaturePositive Ions (Cations)
a) Cations formed from metal atoms have the same
name as the metal, e.g. Na+ is the sodium ion.
b) If a metal can form different cations, the positive
charge is indicated by a Roman numeral in
parentheses following the name of the metal,
e.g. Au+ is the gold(I) ion and Au3+ is the gold(III)
ion.
c) Cations formed from nonmetal atoms have
names that end in -ium, e.g. NH4+ is the
ammonium ion.
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Chemical NomenclatureCommon Cations
Table 1.8
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Chemical NomenclatureNegative Ions (Anions)
a) The names of the monatomic anions are formed
by replacing the ending of the name of the element
with -ide, e.g. O2- is the oxide ion.
b) Polyatomic anions containing oxygen (called
oxyanions) have names ending in -ate or -ite, e.g.
SO42- is the sulfate ion and SO3
2- is the sulfite ion.
c) Anions derived by adding H+ to an oxyanion are
named by adding the prefix hydrogen or
dihydrogen, e.g. HCO3-is the hydrogen carbonate
ion.
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Chemical NomenclatureCommon Anions
Table 1.9
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Chemical NomenclatureMore on naming oxyanions
Examples:
ClO4-
perchlorate ion (one more O atom than chlorate)
ClO3-
chlorate
ClO2-
chlorite ion (one less O atom than chlorate)
ClO-
hypochlorite ion (one O atom less than chlorite)
Names of ionic compounds consist of the cation followed by the
anion name, e.g. Cu(ClO4)2 is copper(II) perchlorate, and CaCO3
is calcium carbonate.
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Chemical NomenclatureName and Formulae of Acids
1. Acids containing anions whose names end in -ide are named by changing the -ide ending to -ic, adding the prefix hydro- to this anion name, and then following with the word acid.
2. Acids containing anions whose names end in -ate or -ite are named by changing the -ate ending to -ic and the -ite ending to -ous and then adding the word acid.
Figure 1.36
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Chemical NomenclatureBinary Molecular Compounds
1. The name of the element farther to the left in the periodic table is written first.
2. If both elements are in the same group in the periodic table, the one having the higher atomic number is written first.
3. The name of the second element is given an -ide ending.
4. Greek prefixes are used to indicate the number of atoms of each element.
Examples
N2O4 is dinitrogen tetroxide
P4S10 is tetraphosphorus decasulfide.
Table 1.10