Section 2 Atoms, Molecules , and Ions

Elements, Compounds and Mixtures

Begin Chapter 2
[Silberberg]

Elements:

have definite compositions and constant properties.
Elements cannot be broken down further by an ordinary chemical change.
All of an element's atoms share a property (the number of protons)
No further separation (reduction) is possible by chemical means.

Examples: He, Fe, Ar, Al
With multiple atoms of one type: H₂, O₂, N₂, S₈, O₃

Compounds:

have definite compositions and constant properties.
Compounds cannot be physically separated into more components.
However, compounds can be separated by chemical means into multiple elements.
Compounds generally are made up of 2 or more elements in fixed ratio and bonded relationship.

Law of Definite Proportion

Law of definite composition: Compounds always contain the same elements in a constant proportion by mass.

Molecule: a particle composed of two or more atoms in a consistent fashion. This usually excluded metals because they bind in a non-consistent fashion (metal alloys are mixtures, not compounds). Molecules are the smallest unit of a compound.

Law of Multiple Proportions

When two elements form a series of compounds, the ratios of the masses of the second element that combine with 1 gram of the first element can always be reduced to small whole numbers.

Pure substance:

is material that has definite composition and constant properties. A pure substance can be an element or a compound.

Mixture:

A Material that can be separated using physical methods into two or more substances. Mixtures have variable composition and properties.

Solution:

An Homogeneous Mixture.

Homogeneous mixtures (Solution):

A mixture that has uniform properties through out the material

Examples:

Air: a gaseous homogeneous mixture of nitrogen, oxygen and other trace gases.

Salt water: a liquid homogeneous mixture of liquid water and solid salt.

Alloy: a solid homogeneous mixture of two or more solid metals.

Coffee: a homogeneous mixture of water (aqueous mixture), coffee bean extract, cream and sugar

Heterogeneous mixtures:

A mixture whose properties vary through the material

Examples:

Salt and ice: Heterogeneous mixture (a small sample may be mostly ice or mostly salt)

Iced coffee: an heterogeneous mixture consisting of an aqueous coffee extract with cream and sugar along with small chunks of ice (frozen water)

End Chapter 1
[Browm/LeMay]

First Recorded Western Concept of Atom

Democritus (460-370 B.C.) – Greek philosopher

Proposed that matter cannot be divided indefinitely,
but that there was a limit to the size that matter could be divided.
That limit was a particle called an "atom".

Dalton Model of the Atom (Historical)

Proposed in 1803

All matter composed of tiny, indestructible, indivisible particles.

· Elements composed of tiny, indivisible, indestructible particles called atoms.
(Not true, atoms composed of protons, neutrons, and electrons which are composed of smaller particles.)

· All atoms of an element are identical and have the same properties.
(Not true, atoms of one element vary in mass, and number of neutrons, the different versions are called isotopes)

· Atoms of different elements combine to form compounds

· Compounds contain atoms in small whole number ratios

· Atoms may combine in more than one ratio to form different compounds.

· Chemical reactions involve reorganizing the atoms, changing the way they are bound together.

· Atoms themselves are not changed in chemical reactions.

Electrons and Protons

Thompson Model of the Atom (Historical)

Cathode rays were observed to be influenced by magnetic fields in the 1870's.
The cathode ray was shown to be deflected by an electric field in 1897.
These were decided to be small negative charged particles termed electrons, (e^-)

In 1886, a positive ray (canal ray) was observed.
Hydrogen gas produced the smallest positive particles, this was called the proton. (p⁺)

J.J. Thompson was able to measure the charge to mass ratio of the electron and proton.

In 1911, Robert Millikan was able to measure the charge on an electron. This allowed the determination of the mass of the electron and proton.

Electron = 9.11 x 10^-28 g Charge = negative

Proton = 1.67 x 10^-24 g Charge = positive

Electron and proton have the same relative charge with opposite sign.

In 1903, Thompson proposed the
Atom consisted of a sphere of positive charge with negative electrons scattered through it. (Plum pudding model, or raisin pudding model)
This has been shown to be incorrect.

Protons and Neutrons in Nucleus with Electrons outside the nucleus

Rutherford Model of the Atom (Historical)

Rutherford was studying radiation: alpha, beta, and gamma (which he discovered)

In an alpha scattering experiment by Rutherford using thin gold foil,
a very small number of alpha particles were scattered backwards.
This was taken to mean that very small dense centers existed in the atom.

In 1911, Rutherford proposed that the atom consisted of a small dense atomic nucleus containing protons surrounded by a largely empty space with moving electrons.

He also proposed that the atomic nucleus also contained neutral particles in addition to protons.
The neutral particles are called neutron (n^o).
These were discovered in 1932 by James Chadwick.

The atom had a diameter about 1 x 10^-8 cm

The nucleus has a diameter of 1 x 10^-13 cm (and a density about 10¹³ – 10¹⁴ g/cm³)

Subatomic Particles

Particle	Symbol	Location	Relative charge	Mass (amu)	Mass (g)
Electron	e^-	Outside nucleus	-1	5.486 x 10^-4	9.11 x 10^-28 g
Proton	p^+	Inside nucleus	+1	1.0073	1.67 x 10^-24 g
Neutron	n^o	Inside nucleus	0	1.0087	1.67 x 10^-24 g

Structure	Size
Atom	10^-8 cm
nucleus	10^-13 cm

Defining Property of Elements – Number of Protons

Atoms consist of three particles: Electrons, Protons, and Neutrons.

Elements are defined by number of protons.

For neutral elements: # electrons = # protons

Each element has an atomic number (Z), which is the number of protons in that element.

Isotopes

Most elements have several different numbers of neutrons in the atomic nucleus.
About 20 elements have only one fixed number of neutrons in the atomic nucleus.
The atoms with different number of neutrons are referred to as isotopes of the elements.

This is why the atomic masses of elements are not whole numbers,
they are averages of the naturally occurring isotopes.

To refer to an isotope we state the element followed by the mass number.
(the element is associated with a specific atomic number which we don't have to specify)

Some isotopes are not stable and will radioactively decay into other elements.

We should be able to state number of electrons, protons, and neutrons in any isotope.

Atomic Notation

Each element defined by number of protons in atomic nucleus.
The atomic number (Z) is the number of protons in the nucleus.

The total of protons and neutrons in the atomic nucleus is the mass number (A).

Atomic notation designates atomic and mass numbers associated with an isotope.
Mass number on top and atomic number on bottom before the element symbol: ^A_ZSy

In Isotopes, mass number varies, atomic number is constant.

Examples: ¹₁H; ⁴₂He; ⁷₃Li; ⁹₄Be; ¹¹₅B; ¹²₆C; ¹⁴₇N; ¹⁶₈O; ¹⁹₉F; ²⁰₁₀Ne; ²³₁₁Na

Number of Protons equals atomic number (Z).

Number of Neutrons equals mass number minus atomic number (A – Z).

Number of Electrons equals number of protons (Z). (atoms have a neutral charge)

Example of Isotopes: Hydrogen

Most Common isotope: ¹₁H Hydrogen-1
Contains: 1 proton & 0 neutrons

Minor natural isotope: ²₁H Hydrogen-2 (deuterium)
Contains: 1 proton & 1 neutron

Radioactive isotope: ³₁H Hydrogen-3 (tritium)
Contains: 1 proton & 2 neutrons

Atomic Mass

Definition

Carbon-12 is the reference standard having a mass of 12 atomic mass units (amu)

1 amu is the mass of 1/12 the mass of carbon-12

That means that the mass of carbon-12 is 12.000000…

All other masses are relative to this and will not be exact.

Example:

Magnesium (Mg)-26 atomic mass is 25.983 amu

Simple and Weighted Averages

Simple average (SⁿV_i)/n
[equal to weighted averaged with identical weights]

Weighted average Sⁿ W_i·V_i

Were V is the Value, and W is the Weighting Factor

Atomic Mass of an Element

Atomic Mass is a weighted average using the relative abundance (on Earth) of each naturally occurring isotope of an element.

Example:
Boron has two isotopes of mass 10 and 11.
The relative abundance of both isotopes is 19% of Boron-10 and 81% of Boron-11

Atomic mass of Boron = 0.19 x 10 + 0.81 x 11 = 10.81 amu

Periodic Repetition of Properties

Dmitri Mendeleev proposed that the properties of elements repeat at regular intervals when they are arranged in order of increasing atomic mass.

He presented 8 groups organized by the formula of the oxide

Periodic Law: When the elements are arranged in order of increasing relative mass, certain sets of properties recur periodically.

The Periodic Law Concept by Atomic Number

In 1913, nuclear charge (atomic number) of elements was shown to increase by 1 for each element.

Periodic law - the properties of elements recur in a repeating pattern when arranged according to increasing atomic number (only a few exceptions from increasing atomic mass)

In 1920's, Niels Bohr introduced concept of energy levels, this was used to shape the periodic table into the form used today.

The Periodic Table

Elements are arranged in order of increasing atomic number,
corresponding to increasing atomic mass.

The periodic table shows the name of the element, the symbol of the element, and the atomic number.

The periodic table arrangement reveals similar properties of the elements
along with electronic structure of the elements.

Nonmetals are generally in the upper-right corner of the table.
Metals are the bulk of the left and bottom of the Periodic Table.
The semi metals lie between the metals and the nonmetals.

Groups of elements with similar properties are located in columns.
These columns are labeled either with a sequential number 1-18
or with a Roman numeral and a letter, e.g. IA, IIA, IIIA, IVA, ... VIIIA, IB, ... VIIIB.

Physical States of the Elements

Standard state: 25°C and 1 atm pressure.

Gases are located on the upper-right of the periodic table.

Most elements are solid

Two elements are liquid: bromine and mercury

Periodic Table

The periodic table lists the atomic number near the top of the box and the atomic mass near the bottom of the box.

The isotopes are not normally listed.

Non-stable/radioactive elements list the mass number of one isotope in parenthesis.

However, even the stable elements have naturally occurring radioactive isotopes.

Different tables may list elements differently.

Groups and Periods of Elements

Row: Period or Series
Numbered 1 to 7

Vertical column: Group or Family
Numbered IA to VIIIA and IB to VIIIB (American System)
or 1 - 18 (IUPAC system)
The European system has different A and B numbers
We use American or IUPAC numbers.

Hydrogen can be placed in Group IA/1 or VIIA/17 or totally separate

Group IA/1 are the Alkali Metals

Group IIA/2 are the Alkali Earth Metals

Group VIIA/17 are the Halogens

Group VIIIA/18 are the Noble gases or Inert Gases

Other groups can be called by the first (top) element in the group. Example: Group IB/11 is called the copper group.

Metals, Nonmetals and semimetals

Elements were broadly classified as metals and nonmetals based on appearance and properties. A third category was added later of semimetals which has properties in-between metals and nonmetals. A semimetal is also called a metalloid.

Metal	Nonmetals	Semimetals (metalloid)
Bright luster High density High melting point Good conductor of heat Good conductor of electricity Malleable and Ductile Solid (1 exception, Mercury)	Dull appearance Low density Low melting point Poor conductor of heat Poor conductor of electricity Brittle Solid or gas (1 liquid)	Properties in-between metal and nonmetal
Majority of elements all solid except for mercury, Hg, which is liquid	17 elements 5 solid 11 are naturally gases 1 liquid; bromine, Br	8 elements, all solid Boron, B Silicon, Si Germanium, Ge Arsenic, As Antomony, Sb Tellurium, Tc Polonium, Po Astatine, At

Compounds and Chemical Formulas

Law of definite composition: Compounds always contain the same elements in a constant proportion by mass.

Molecule: a particle composed of two or more atoms in a consistent fashion. This usually excluded metals because they bind in a non-consistent fashion (metal alloys are mixtures, not compounds). Molecules are the smallest unit of a compound.

Chemical formula:
expresses the number of each type of atom in a molecule or compound.
Subscripts are used with the element symbol to express the number of atoms of that element present; however, the subscript "1" is omitted.

Chemical formula can be written generically showing only elements and number of atoms (written as a subscript). Or it can be written in a fashion that reveals information about the structure of the compound.

Examples:

NaHCO₃ (baking soda, sodium bicarbonate, sodium hydrogen carbonate)

Na₂CO₃ (sodium carbonate)

CO₂ (carbon dioxide)

CO (carbon monoxide)

Co (cobalt)

CaCO₃ (limestone, calcium carbonate)

Ca(HCO₃)₂ (calcium bicarbonate, calcium hydrogen carbonate)

CoCl₂ (cobalt chloride)

H₂O (water)

H₂O₂ (hydrogen peroxide)

Diatomic elements: Some elements, under standard conditions, form a molecule consisting of two atoms of that element. These are called diatomic molecules. Naturally, occurring diatomic molecules are:

H₂, N₂, O₂, F₂, Cl₂, Br₂, and I₂

These can be remembered by the phrase: Have No Fear Of Ice Cold Beer
Representing respectively: H₂, N₂, F₂, O₂, I₂, Cl₂, Br₂

This also gives a clue to the physical states of these elements. Most of these elements are gases just as the words are just air. However, Beer is a liquid just as Bromine, and Ice is a solid just as Iodine.

Empirical Formula

Empirical Formula is the lowest whole number ratio of atoms of each element in the formula of a molecule or ions in an ionic compound. This formula can be, but doesn't have to be, smaller than the molecular formula or the formula unit of an ionic compound. The empirical formula is usually an experimentally calculated formula and is not generally used after the actual molecular formula is determined. The molecular formula represents the composition that is actually found in nature. For example hydrogen is always found as a diatomic molecule, H₂.

Covalent Compounds (Molecular Compounds)

Molecular compounds are composed of nonmetal elements.
The simplest representative particle in a compound is a molecule.
Molecules are sets of atoms bonded together in a specific structure.
This category includes both inorganic and organic compounds.
We are studying mostly inorganic covalent compounds.

Examples:
Water; H₂O (inorganic)
Methane (natural gas); CH₄ (organic)
Ammonia; NH₃ (inorganic)
Ethanol; CH₃CH₂OH (organic)

Ionic Compounds

Ionic compounds are compounds formed between a metals and nonmetals. The metals transfer electrons to the nonmetals, creating positive metal ions and negative nonmetal ions or a negative group of atoms (Polyatomic ion).

The compound is a specific ratio of these ions such that the compound is electrically neutral.

The smallest unit is called a formula unit, showing the ratio of ions that makes the compound neutral.

We are studying Ionic Compounds

Examples:
Sodium Chloride; NaCl
Sodium bicarbonate; NaHCO₃
Titanium (IV) oxide; TiO₂

Ionic Charges/Oxidation Number

Ion: an atom or group of atoms with a net electrical charge
[number of electrons has changed and is different from number of protons].

Ionic Charge is the amount of positive or negative charge that an ion has.
i.e. an atom that has lost two electrons has an ionic charge of +2
An atom that has received 3 electrons has an ionic charge of -3.
[The ionic charge carried by an atom is referred as the oxidation number.]

For Representative elements, the ionic charge of an ion can be predicted from the group number the element is in.

The number of valence electrons is shown by the group number. An atom will loose valence electrons, not core electrons.

Group Valence electrons Common
ionic charge
Group IA/1 1 +1 Alkali metals
Group IIA/2 2 +2 Alkaline earth metals
Group IIIA/13 3 +3 (Aluminum only)
Group VA/15 5 -3 (Nitrogen only)
Group VIA/16 6 -2 (for nonmetals)
Group VIIA/17 7 -1 Halogens (includes hydrogen when combined with metal)
Group VIIIA/18 8 0 Noble Gases do not form ions

Expect metals to lose electrons and become positive

Nonmetals gain electrons to become negative

Two or more ions or atoms having the same number of electrons are isoelectronic.
Example: S^2-. Cl^1-, and argon all have 18 electrons and are isoelectronic

Writing Ionic Chemical Formulas

The cation [positive metal ion] is always listed first,
the anion [negative nonmetal ion] is listed last.

The chemical formula (or formula unit) must be neutral, so the charge of the cations must equal the charge of the anions

The cationic charge is the ionic charge of the cation times the number of cations present.

The anionic charge is the ionic charge of the anion times the number of anions present.

These two total charges must be equal

Example:

Al₂O₃ Al³⁺ (+3 x 2 = 6); O^2- (-2 x 3 = -6)

Simple Method:

Take the charge number of the cation and use it as the number of anions present

Take the charge number of the anion and use it as the number of cations present

Reduce if multiple of each other.

Example:

Lead (IV) oxide: Pb⁴⁺ and O^2-

Reversing the numbers yields Pb₂O₄

Dividing by the common multiple of 2 yields PbO₂

Formula Units Containing Polyatomic Ions

The chemical formula (or formula unit) must be neutral, so the charge of the cations must equal the charge of the anions

Procedure is the same as for binary ionic compounds, if there is multiple units of a polyatomic ion, then parenthesis is needed around the ion before the subscript number of ions is added.

Ammonium sulfite (NH₃)₂SO₃

Aluminum sulfate Al₂(SO₄)₃

Monoatomic Ions

Naming of Monoatomic Cations

Type 1 Ions

The name of the metal followed by the word "ion"
Al⁺³ is the aluminum ion
Ca⁺² is the calcium ion
K⁺ is the potassium ion

Main group elements usually form only one cation
[tin and lead are exceptions, both for +2 ions along with the predicted +4 ions]

Transition elements form several cations [Ag⁺, Zn²⁺, Cd²⁺ are exceptions]

Type 2 Ions

Stock system: the name of the metal is followed by roman numerals in parentheses indicating the ionic charge. (this is the simpler system and requires less memorization)

Iron (Ferrum): Fe⁺² is iron (II) ion and Fe⁺³ is iron (III) ion

Copper (cuprum): Cu⁺ is copper (I) ion and Cu⁺² is copper (II) ion

Latin system: If the metal has two ionic charges, the lower of the two receives an -ous suffix, the higher charge receives an -ic suffix on the latin name of the metal. (required knowing the latin name [one exception, mercury] of the metal and which charge is higher or lower.)

Iron (Ferrum): Fe⁺² is ferrous ion and Fe⁺³ is ferric ion

Copper (cuprum): Cu⁺ is cuprous ion and Cu⁺² is cupric ion

Tin (stannum): Su⁺² is stannous ion and Sn⁺⁴ is stannic ion

Naming of Monoatomic Anions

Nonmetal anions: nonmetal stem + -ide suffix

Examples:
Br, Bromine; Br^-, bromide
S, Sulfur; S^-2, sulfide
N, Nitrogen; N^-3, nitride
O, Oxygen; O^-2, oxide
H, Hydrogen H^-, hydride

Predicting Formulas of Monoatomic Cations

Predictable Charges

Alkali metals (Group IA) form a +1 cation
Alkaline Earth metals (Group IIA) form a +2 cation
Silver (Group IB) forms a +1 cation
Zinc and Cadmium (Group IIB) form a +2 cation
Aluminum (Group IIIA) forms a +3 cation

Variable Charges

Transition elements: can usually have a +2 charge along with other hard to predict charges. The other metals not listed with a predictable charge have multiple charges.

Predicting Formulas of Monoatomic Anions

Nonmetals usually gain electrons to become isoelectronic with a noble gas

Isoelectronic means having the same number of electrons or the same electron configuration.

So:
Halogens accept one electron to form an -1 anion
Group VIA nonmetals gain 2 electrons to form an -2 anion
Group VA nonmetals gain 3 electrons to form an -3 anion

Binary Ionic Compounds (Metal with Nonmetal)[Type I]

The smallest unit is called a formula unit.

The cation is always listed first, the anion is listed last.

The overall charge is neutral

Binary Compounds Containing a Transition Metal [Type II]

The smallest unit is called a formula unit.

Transition metals can have one of several ionic charges

We can predict the ionic charge of the transition metal using the charge of the anion.

The anion carries a predictable charge. This predictable charge can be used to find the charge on the cation.

Find the total negative charge and divide by the number of cations to get the ionic charge on the cation.

FeCl₃ has a negative charge of 3 x -1 = -3
3/1 Fe yields a +3 ionic charge on Fe.
Fe³⁺, ferric ion, or iron (III) ion

Naming Binary Ionic Compounds

Binary ionic compounds are named using the cation name followed by the anion name.

FeCl₃ is called iron (III) chloride or ferric chloride

Predicting Formulas of Binary Ionic Compounds

We can predict formulas using the predicted charges of the ions from the periodic table,

also, if a formula is found for one compound, then the other compounds in the group can be expected to have a similar formula.

Ternary Ionic Compounds

contains three elements, with at least one metal and one nonmetal [LiNO₃ ]

Usually contains a monoatomic cation and a polyatomic anion.

Usually has a name that ends in -ate or -ite

Calcium carbonate CaCO₃

Sodium chlorate NaClO₃

Potassium hydroxide KOH

Magnesium nitrite Mg(NO₂)₂

Ternary Compounds Containing a Transition Metal

Determine the ionic charge on the metal and then name using stock or latin systems

Fe₃(PO₄)₂ has a negative charge of 2 x -3 = -6
6/3 Fe yields a +2 ionic charge on Fe
Fe²⁺, ferrous ion, or iron (II) ion

Fe₃(PO₄)₂ is called iron (II) phosphate or ferrous phosphate

Naming Ternary Ionic Compounds

These are named using the name of the metal cation followed by the name of the anion.

Transition metals are named using either the stock or Latin systems.

Predicting Formulas of Ternary Ionic Compounds

Formulas can be predicted using the expected charge of the group of the representative elements.

Alternatively, and especially for transition metals, the formula can be predicted from other elements in the group.

Polyatomic Ions

Names usually end with the suffix -ate or -ite. When both -ate and -ite suffix exist with the same root, the -ate suffix has one more oxygen

There are a couple of -ide suffix names

Names of Ions

Anions
Names	Formula	Names	Formula
Sulfite	SO3^2-	nitrite	NO2^-
sulfate	SO4^2-	nitrate	NO3^-
hydrogen sulfate (bisulfate)	HSO4^-	Carbonate	CO3^2-
hypochlorite	ClO^-	hydrogen carbonate (bicarbonate)	HCO3^-
chlorite	ClO2^-	phosphate	PO4^3-
chlorate	ClO3^-	acetate	C2H3O2^-
perchlorate	ClO4^-	hydroxide	OH^-
		cyanide	CN^-
Cations
Ammonium ion	NH4^+

Binary Molecular Compounds [Covalent Compounds]

Molecular compounds are composed of nonmetal elements.

The simplest representative particle in a compound is a molecule.

The order of writing the formula is from most metallic elements to most nonmetallic.

Or in the order of

C, P, N, H, S, I, Br, Cl, O, F

Examples: CO₂, PH₃, NH₃, H₂O, SO₂ IF₆, Br₃O₈, Cl₂O₅,

Carbon dioxide, phosphorous trihydride, nitrogen trihydride (ammonia), dihydrogen oxide (water), sulfur dioxide, iodine hexaflouride, tribromine octaoxide, dichlorine pentaoxide

Naming Binary Molecular Compounds

IUPAC names, the second element has an -ide suffix (common names vary)

The number of atoms of an element is shown using the following prefixes

Atoms	Prefix	Atoms	Prefix
1	mono	6	hexa
2	di	7	hepta
3	tri	8	octa
4	tetra	9	nona
5	penta	10	deca

the prefix mono is usually not used unless it is needed to prevent confusion

However, CO and NO use mono in their names; carbon monoxide and nitrogen monoxide

Examples: CO₂, PH₃, NH₃, H₂O, SO₂ IF₆, Br₃O₈, Cl₂O₅,

Carbon dioxide, phosphorous trihydride, nitrogen trihydride (ammonia), dihydrogen oxide (water), sulfur dioxide, iodine hexaflouride, tribromine octaoxide, dichlorine pentaoxide

Hydrogen chloride (HCl), Hydrogen bromide (HBr)

Binary Acids

Acids: substances that release hydrogen ions (H⁺) in water

Binary acid: an aqueous solution containing a compound of hydrogen and one other nonmetal [HF(aq)]

The formula always begins with H [HCl, HF, HBr]

Naming: Start with hydro use the nonmetal stem with an -ic suffix, finish with the word acid

HCl (aq) hydrochloric acid [(chlor)ine -ic]

HF (aq) hydrofluoric acid [(fluor)ine -ic]

HBr (aq) hydrobromic acid [(brom)ine -ic]

Note: HCl (aq) is hydrochloric acid (a binary acid) while HCl (g) is hydrogen chloride (a binary molecular compound)

Ternary Oxyacids

Ternary oxyacids: Aqueous solutions of hydrogen bonded to a polyatomic anion;
these acids contain hydrogen, oxygen, and another element [H₂SO₄ (aq)]

Acids with the -ic ending are formed from polyatomic anions with the -ate ending

Example:

Nitrate [NO₃^-] à Nitric acid, HNO₃

Acids with the -ous ending are formed from polyatomic anions with the -ite ending

Examples:

Chlorite [ClO₂^-] à chlorous acid, HClO₂

Hypochlorite [ClO^-] à hypochlorous acid, HClO

Ternary Oxyacid Examples

Acid	Anion	Acid	Anion
Sulfuric acid H2SO4	Sulfate ion SO4^2-	Sulfurous acid H2SO3	Sulfite ion SO3^2-
Nitric acid, HNO3	Nitrate ion, NO3^-	Nitrous acid, HNO2	Nitrite ion, NO2^-
Perchloric acid	Perchlorate ion	Chlorous acid	chlorite ion
Chloric acid	Chlorate ion	Hypochlorous acid	Hypochlorite ion
Phosphoric acid	Phosphate

Note: sulfur and phosphorus are exceptions to using the stem of the anion in naming the acid; the stem expands in going from the anion to acid.

Organic chemistry

Goal:

To be able to name the small normal alkanes and identify and name several functional groups.

Organic chemistry is the study of carbon-containing compounds. Many, but not all, of these compounds are biomolecules, molecules that participate in life processes.

Alkanes: Saturated Hydrocarbons

Hydrocarbons

Hydrocarbons are compounds composed of hydrogen and carbon.

alkanes

Alkanes are saturated hydrocarbons, meaning that they only have carbon-carbon single bonds with no double or triple bonds.

normal hydrocarbons (straight-chain hydrocarbons, unbranched hydrocarbons)

Normal hydrocarbons are straight-chain, unbranched, hydrocarbons. The term "straight-chain" does not mean linear since the bond angles are 109° and the compound can rotate around forming many shapes. The term "straight-chain" means unbranched.

Selected normal hydrocarbons

Name	Formula	# structural isomers
Methane	CH4	1
Ethane	C2H6	1
Propane	C3H8	1
Butane	C4H10	2
Pentane	C5H12	3
Hexane	C6H14	5
Heptane	C7H16	9
Octane	C8H18	18
Nonane	C9H20	35
Decane	C10H22	75

Hydrocarbon Derivatives

Hydrocarbon Derivatives are hydrocarbons that contain other elements in addition to carbon and hydrogen.

Functional groups

Functional groups are atoms or groups of atoms attached onto hydrocarbons. These functional groups exhibit characteristic chemistry.

Common Functional groups

Name	Functional Group	General Formula
Halohydrocarbons	-X (F, Cl, Br, I)	R-X
Alcohols	-OH	R-OH
Ethers	-O-	R-O-R'
Aldehydes	O || -C-H	O || R-C-H
Ketones	O || -C-	O || R-C-R'
Carboxylic Acids	O || -C-O-H	O || R-C-O-H
Esters	O || -C-O-	O || R-C-O-R'
Amines	-NH2 -NH- | -N-	R-NH2 R-NH-R' NRR'R"

R, R', R" represent any hydrocarbon fragment.

Alcohols

Alcohols are characterized by the presence of the hydroxyl group (-OH).

Naming and classifying alcohols

The name is obtained by replacing the -e of the parent hydrocarbon with -ol.

The location of the hydroxyl group is specified by a number, where necessary.

Alcohols are also classified by the number of hydrocarbon fragments bonded to the carbon where the hydroxyl group is attached.

Primary alcohol - has one hydrocarbon group on carbon attached to alcohol group.

Secondary alcohol - has two hydrocarbon groups on carbon attached to alcohol group.

Tertiary alcohol - has three hydrocarbon groups on carbon attached to alcohol group.