CHEM 110

Chapter One and Two

2023

Ms F Nevondo 1
Chemistry

The study of matter and its transformations study of physical &

chemical properties of matter - what changes occur in these properties,
as the result of a chemical reaction, & how these changes may be
observed - why the reaction involved does or doesn’t occur

Focus: How matter interacts at the atomic/molecular level

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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 the building blocks of matter.
• Each element is made of the same kind of atom.
• A compound is made of two or more different kinds of elements.

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 Pure Substances, Elements and Compounds
Pure Substance
• Matter that has distinct properties and a composition.
• This does not vary from sample to sample, e.g. salt.

Element
• A substance that cannot be decomposed into simpler substances, e.g.
oxygen gas.

Compound
• A substance that is composed of two or more different elements, so it
contains two or more different kinds of atoms, e.g. water.
• The elemental composition of a pure compound is always the same and this is known
as the Law of Constant Composition (or Law of Definite Proportions) Joseph Proust

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Mixtures
•Amixture is a combination of two or more substances in which
each substance retains its own chemical identity and can be
separated from each other.

There are two types:

• homogeneous: a mixture which is uniform throughout, e.g.

vanilla ice-cream.

• heterogeneous: a mixture which does not have the same

composition, properties and appearance throughout, e.g. muesli.

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Classification of Matter

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Properties of Matter
Physical Properties
• Can be observed without changing a substance into another
substance.
• Boiling point, density, mass, volume, etc.

Chemical Properties
• Can only be observed when a substance is changed into another
substance.
• Flammability, corrosiveness, reactivity with acid, etc.

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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)

Brown, LeMay, Bursten, Murphy, Langford, Sagatys: Chemistry 2e © 2010 Pearson Australia

Chemical Reactions (Chemical

Change)
In the course of a chemical reaction, the
reacting substances are converted to new
substances.

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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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Molecular Comparison of Substances and
Mixtures

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Separation of Mixtures

1. Distillation

Separates a
homogeneous
mixture on the
basis of differences
in boiling point.

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2. Filtration
Separates solid substances from liquids and solutions.

3.Chromatography
Separates substances on the basis of differences in solubility in
a solvent.
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Scientific Measurement
• Used to measure quantitative properties of matter
• SI (Système International d’Unités) units

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Metric System
• Prefixes convert the base units into units that are appropriate for
the item being measured.

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SI Units - Length and Mass

• The SI base unit of length is the metre (m).

• Mass (m) is a measure of the amount of material in an object.

The SI base unit of mass is the kilogram (kg).

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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

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 Temperature
Celsius
• Represented by °C
• Based on freezing point of water as 0°C and boiling point of
water as 100°C

Kelvin (SI unit)

• Represented by K (no degree sign)
• The absolute scale
• Units of Celsius and Kelvin are equal in magnitude

Fahrenheit (the English system) (°F)

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Equations for Temperature Conversions

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Temperature Conversions

• Ice cream is best stored at -24°C. What is this temperature on

the Kelvin scale?

K =°C+ 273.15

K= െ24 ι‫ ܥ‬+ 273.15 = 249.15 ‫ܭ‬

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Practice

• Convert the temperature reading of the ice cream(-14°C) into

the corresponding Fahrenheit temperature.

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Derived SI Units Volume

• The most commonly used metric units for volume are the litre
(L) and the millilitre (mL).
• A litre is a cube 1 dm long on each side.
• A millilitre is a cube 1 cm long on each side.

• 1dm3 = (1 dm) x (1 dm) x (1 dm) = 10 cm x 10 cm x 10 cm =

1000 cm3 = 1 L

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Derived SI Units Density
• Density is a physical property of a substance.
• Ratio of mass to volume
• It is determined through the following formula:

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Practice
• Calculate the density of a ruby which has a volume of 1.8 cm3
and a mass of 8.7g.

• The density of benzene at 15 °C is 0.8787 gcm-3. Calculate the

mass of 0.1730 dm-3 of benzene at this temperature.

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Scientific notation
• Is used to write very large or very small numbers
• The width of a human hair (0.000 008 m) is written as 8 x 10-6 m
• A large number such as 4 500 000 s is written as 4.5 x 106 s

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Scientific notation
• A number written in scientific notation contains a coefficient and a
power of 10.

• To write a number in scientific notation, the decimal point is moved

after the first digit
• The spaces moved are shown as a power of ten.

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Some Powers of Ten

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Practice
Select the correct scientific notation for each.

A. 0.000 0009
1) 9 x 107
2) 9 x 10-6
3) 0.9 x 10-6

B. 720 000
1) 7.2 x 105
2) 72 x 104
3) 7.2 x 10-4

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Known + Estimated Digits
In the length reported as 2.76 cm,

• the digits 2 and 7 are certain (known)

• the final digit 6 was estimated (uncertain)
• all three digits (2.76) are significant including the estimated
digit

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Uncertainty in Measurements
• Different measuring devices have different uses and different
degrees of accuracy.

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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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Guidelines for significant figures

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 Guidelines for significant figures
• All nonzero digits are significant
e.g. 123.45 has 5 significant figures

• Zeros between two significant figures are themselves significant

e.g. 103.405 has 6 significant figures

• Zeros at the beginning of a number are never significant

e.g. 00123.45 = 123.45 has 5 significant numbers

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Practice
• Determine the number of significant figures in each of
the following.
- 3.455 cm
- 0.00068 g
- 105 m
- 25.0 mL
- 0.10 mL

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Rounding off answers
In calculations

• Answers must have the same number of significant figures as the

measured numbers

• Calculated answers are usually rounded off

• Rounding rules are used to obtain the correct number of

significant figures

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Rounding off Procedure
1. Drop off the digit that follows if it is less than 5

e.g. 8.724 Æ 8.72

2. Add 1 to the preceding digit if it is equal or greater than 5

e.g. 8.727 Æ 8.73

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 Addition and subtraction
When adding or subtracting, answers are rounded to the
least significant decimal place.
102.50 two digits after decimal point
+ 0.231 three digits after decimal point
102.731 round to 102.73
Addition and subtraction
2.097 – 0.12 = 1.977 = 1.98
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Multiplication and division
When multiplying or dividing,

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.

Multiplication and division 1.4 x 8.011 = 11.2154 round to

11 (Limited by 1.4 to two significant figures in answer)

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 Exact numbers
• Exact numbers have an infinite number of significant figures.

Example: A coin issued after 1982 has a mass of 2.5 g. If we have three
such coins, the total mass is

3 x 2.5 g = 7.5 g

• In this case, 3 is an exact number and does not limit the number of significant
figures in the result.

• Similarly the average of two measured lengths

6.64 cm and 6.68cm = (6.64 + 6.68) ÷ 2

= 6.66 cm
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 Practice
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.

ii. Density (ȡ) of a pure liquid, in g cm-3, if 103.67 g of it is needed

to fill the box completely.

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 ȡ of this
liquid in g cm-3.
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 Precision and Accuracy
• Two ways to gauge the quality of a set of measured numbers
• 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.

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Describe accuracy and precision for each set

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Atomic Theory

• The theory that atoms are the fundamental building blocks of

matter came into being during the period 1803 to 1807 in the
work of an English schoolteacher, John Dalton.

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 Dalton’s Postulates
John Dalton (1766 ~ 1844)
Dalton’s Atomic Theory
(1) Each element is composed of atoms
(2) All atoms of a given element are identical, but they are
different from the atoms of all other elements
(3) Atoms are neither created nor destroyed in chemical
reactions.
(4) Compounds are formed from chemical combination of two
or more atoms.

Dalton proposed that all matter is made up of atoms and

stated that elements are the simplest form of matter.
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What can Dalton’s theory explain?
(1) Law of constant composition
• In a given compound, the relative numbers and kinds of atoms are
constant. [postulate 4]
(2) Law of conservation of mass
• The total masses of material present before and after chemical
reaction are identical [postulate 3]
(3) Law of multiple proportions
• If elements A & B combine to form more than one compound, the
masses of B which can combine with a given mass of A are in the
ratios of small whole numbers

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The discovery of atomic structure
After Dalton’s atomic theory, not much of progress had been made and no one
had direct evidence for the existence of atom. Then, things started to change in
late 1800s.
• William Crooks (1832 ~ 1919): Cathode-ray tube (CRT) [1879]

A high voltage between two electrodes in a partially evacuated tube

generates electrical discharge
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 The discovery of atomic structure
J. J. Thomson (1856 ~ 1940) : Discovery of electron [1897]

(1)Rays are the same regardless of the identity of the

cathode material
(2) Conduct quantitative analysis of the effect of
electric and magnetic field Æ determine the charge
He discovered that cathode rays are
to mass ratio
negatively charged particles, which he
originally called `corpuscles’’ . He won a
Nobel prize in physics [1906]. charge/mass = 1.76 × 108 C/g
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 The discovery of atomic structure
Robert Millikan (1868 ~ 1953): Determine the charge of electron
[1907]
Millikan’s oil-drop experiment

The machine on the right hand side is the original

apparatus Millikan used to perform his oil-drop
experiment.
He won a Nobel prize in physics [1923].
Measured charge = 1.60 × 10-19 C
Electron mass = charge/[charge/mass]
= Ms
9.10 × 10-28 g
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 The discovery of atomic structure
Radioactivity

The spontaneous emission of radiation by an atom was first observed by

Henri Becquerel. It was also studied by Marie and Pierre Curie.

Three types of radiation were discovered by Ernest Rutherford

• Į particles

• ȕ particles

• ܵUD\V

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Ernest Rutherford shot Į particles at a thin sheet of gold foil and
observed the pattern of scatter of the particles.

• From the scattering experiment….

(1) Most Į-particles simply pass through the gold foil.
(2) Small amount of scattering was observed at large
angles.

• Rutherford postulated that..

(1) Most of the total volume of an atom is empty space.
(2) Most of the mass of an atom and all of its positive
charge reside in a very small region,
called nucleus.

Rutherford also found the existence of protons inside of nucleus [1919].

Another particle in nucleus, neutron, was found by James Chadwick
in 1932.
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 Modern atomic structure
The list of subatomic particles has grown considerably since the discovery of
electrons, but only the electron, proton and neutron have a bearing on chemical
behaviour.
• Protons and electrons are the only particles that have a charge.

• Protons and neutrons have essentially the same mass.

• The mass of an electron is so small we ignore it.

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 Modern atomic structure
Properties of subatomic particles

• Every atom has an equal number of protons and electrons so that it has no electrical charge, the
atom is held together by force of Coulombic/ electrostatic attraction
• More than 99.99 % of atom mass is centred in the nucleus, where nucleons (protons) &
neutrons, are collectively bound together by strong nuclear force
Atomic Mass Unit (amu)
1 amu = 1/12 of the mass of carbon (12C) atom
= 1.66054 × 10-24 (g)
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 Modern atomic structure
The characteristics of each atom are determined by the numbers of
proton, neutron and electrons.

Atomic Number (Z): The number of protons in the nucleus of an atom.

Mass Number (A): The total number of protons plus neutrons in the atom
Isotopes : Atoms with identical atomic numbers but different mass
numbers such as C-14 and C-12.

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Atomic Numbers, Mass Numbers and
Isotopes

Mass Number (A)

12
6 C Symbol of element

Atomic Number (Z)

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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.

• The chemical properties of isotopes is largely similar, but

physical properties, & particularly the ones involving radioactive
“nuclei”, can be very different
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Practice

Element name Symbol Z A No of e-

73

Ruthenium

39

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 Atomic Mass
Atomic and molecular masses can be measured with great accuracy
with a mass spectrometer.
Atomic Mass Unit (amu) = 1.66054 × 10-24 g

C = 12 amu (exact), 1H = 1.0078 amu, 16O = 15.9949 amu

12

Average Atomic Masses : Weighted average of all the isotopes of an element found in nature.
Example : Naturally occurring carbon is composed of 98.93% 12C and 1.07 % 13C. What is the
average mass of carbon?

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 Practice:: 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.

• Naturally occurring Mg has three isotopes:

24Mg 78.99 %, 23.9850 amu
25Mg 10.00 %, 24.9858 amu
26Mg 11.10 %, 25.9826 amu

Calculate the average atomic mass of Mg

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 Practice
Chlorine has two naturally occurring stable isotopes: 35Cl 34.968853 amu
37Cl 36.965803 amu

If the (average) atomic mass of naturally occurring elemental Cl is 35.453 amu,

what are the % abundances of the two isotopes?

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 Periodic Table
• A systematic catalogue of elements.
• Elements are arranged in order of atomic number.
If the elements are arranged in order of increasing atomic number, their chemical
properties are found to show a repeating, or periodic, pattern.

Elements having similar properties are placed in vertical columns

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Periodic Table

Alkali
Metal

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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.
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 Molecules and Chemical Formulae
Chemical Compounds

Molecular Ionic

(1) Molecular compounds are composed of more than

one type of atom
H2O, NH3, CH3OH
(2) Most molecular substances contain only non-metallic
atoms
H2O, H2O2, CO, CO2, CH4

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Diatomic Molecules

These seven elements occur naturally as

molecules containing two atoms.
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 Types of Formulae
Empirical formulae: gives the lowest whole-number ratio of atoms of
each element in a compound, e.g. HO.

Molecular formulae: gives the exact number of atoms of each element

in a compound, e.g. H2O2.

Structural formulae: shows 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.

Structural formula Perspective drawing

Ball-and-stick model
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Space-filling model 69
Ions and Ionic Compounds
When atoms lose or gain electrons, they become ions.

Cation: An ion with a positive charge Anion: An ion with a negative

charge

Na+ Cl-
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Ions and Ionic Compounds

Cations are formed by elements on the left side of the periodic

chart.
Anions are formed by elements on the right side of the periodic
chart.
Metals tend to form Cations
Nonmetals tend to form Anions
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Ions and Ionic Compounds
Example: Determine the number of electrons, protons and neutrons
in each of the following ions
No. of protons No. of Neutrons No. of Electrons
14N3-

119Sn4+

52Cr6+

40Ca+

80Br-

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Ions and Ionic Compounds
Ionic Compounds : Cations (metals) and anions (non-metal)
combine to form ionic compounds

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 Ions and Ionic Compounds

Ionic compounds :
(1) Ionic compounds are generally combination of metals and nonmetals
NOTE: Molecular compounds are generally composed of nonmetals only (H2O , CH3OH ,
CH3CH2Cl , …)
(2) Ionic compounds are represented by empirical formulas
їƵƐĞsimplest whole-number ratio of cations and anions
NOTE: There is no discrete (or isolated) molecule of NaCl
(3) Ionic compounds are always neutral. Therefore, the total positive
charge equals the total negative charge
Mg2+ and N3- form Mg3N2 : 3× (+2) + 2×(-3) = 0
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 Ions and Ionic Compounds
Example : Find the empirical formula for the ionic compound made of
given cation and anion

Na, O ї

Al, O ї

Ca, O ї

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Naming Ions and Ionic Compounds

• 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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 Naming Ions and Ionic Compounds

Names of ionic compounds consist of the

cation name followed by the anion name

CaCl2 = calcium + chloride Æ calcium chloride

Names of Positive Ions (cations) :

(1) Cations formed from metal atoms have the same name

as the metal.
Na+ Æsodium ion, Zn+ Æzinc ion, Al3+ Æaluminum ion

NOTE: Ions formed from a single atom are called monatomic ions
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 Naming Ions and Ionic Compounds
(2) If a metal can form different cations, the positive charge is indicated by a
Roman numerical in parenthesis following the name of the metal

Fe2+ їiron (II) ion Cu+ їcopper (I) ion

Fe3+ їiron (III) ion Cu2+ їĐŽƉƉĞƌ(II) ion

These ions are usually transition metals

NOTE: Metals that form only one cation

group 1A їNa+, K+, Rb+

group 2A їMg2+, Ca2+, Sr2+, Ba2+

and Al3+ (group 3A), Ag+ (group 1B), Zn 2+ (group 2B)

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Naming Ions and Ionic Compounds
(3) Cations formed from nonmetal atoms have names that end in -ium
NH4+ їammonium ion H3O+ їhydronium ion
NOTE: These ions are examples of polyatomic ions

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 Naming Ions and Ionic Compounds
Names of Negative Ions (anion) :
(1) The names of monatomic anions are formed by replacing
the ending of the name of the element with –ide.
H- hydrogen їhydride ion, O2- oxygen їoxide ion,
NOTE: polyatomic anions with common names ending with –ide
OH- їhydroxide ion, CN- їcyanide ion

(2) Polyatomic anions containing oxygen (oxyanions)

a) ending with –ate : reserved for the most common oxyanion
NO3- їnitrate ion, SO42- їsulfate ion
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b) Ending with –ite : used for oxyanion with the same charge, but one
fewer O atom than those ending with –ate.
NO2- Æ nitrite ion, SO32- Æ sulfite ion

c) If a series of oxyanions extends to more than two members,

use prefix per- (one more) or hypo- (one fewer)
ClO4- Æ perchlorate ion (one more than –ate)
ClO3- Æchlorate ion
ClO2- Æ chlorite ion
ClO- Æ hypochlorite ion (one fewer than -ite)

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 Naming Ions and Ionic Compounds
NOTE: Oxyanions with the maximum number of oxygen's

(i) Charges increase from right to left.

(ii) Second row elements (C, N) have maximum 3 oxygen atoms
and third row elements (P, S, Cl) have maximum 4 oxygen
atoms (row #2 + 1).
(iii) All names end with –ate except for ClO4-
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 Naming Ions and Ionic Compounds
(3) Anions derived by adding H+ to an oxyanion are named by adding as a prefix the word
hydrogen or dihydrogen.
CO32- : carbonate ion їHCO3- : hydrogen carbonate ion
PO43- : phosphate ion ї,2PO4- : dihydrogen phosphate ion

Type equation here.

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 Names of Binary Molecular Compounds
(1) The name of the element further to the left in the periodic table
appear first. (NOTE: Oxygen is always written last except when
combined with fluorine.)
(2) If both elements are in the same group, the one having the higher
atomic number is named 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

Cl2O : dichloro monoxide

NF3 : nitrogen trifluoride

N2O4 : dinitrogen tetroxide

P4S10: tetraphosphorous decasulfide

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Practice
Before you try to name a compound :

(1) Is the compound ionic or molecular?

(2) For ionic compounds, find the name of each ion. For
molecular compounds, find the number of each atom.
BF3 :
NiO :
KMnO4 :
SO :

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Practice
Write down the chemical formulas for the following compounds
(1) Sodium Nitride,

Q: Is this ionic or molecular?

Q: Is anion monatomic or polyatomic ion?

(2) Diphosphorus pentoxide,

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Practice

(1) NaClO :
(2) Fe2(CO3)3:
(3) SF6 :
(4) aluminium hydroxide :
(5) ammonium sulfate :
(6) NaH2PO4 :

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