chemistry 101 : chap. 2 atoms, molecules and ions (1) atomic theory of matter (2) the discovery of...

Chemistry 101 : Chap. 2Chemistry 101 : Chap. 2

Atoms, Molecules and Ions

(1) Atomic Theory of Matter

(2) The discovery of Atomic Structure

(3) The Modern View of Atomic Structure

(4) Atomic Weight

(5) Periodic Table

(6) Molecules and Molecular Compounds

(7) Ions and Ionic Compounds

(8) Naming Inorganic Compounds

The Atomic Theory of MatterThe Atomic Theory of Matter

The history of development of atomic theory of matter begins in ancient Greece. However, modern atomic theory has it’s origin in a burst of scientific discovery between 1870 and 1930.

Democritus (460 ~ 370 BC)

Democritus proposed atomic theory of matter.

He and other Greek philosophers believed that material world must be made up of hard and tiny indivisible particles that they called atomos, which are in constant motion.


Aristotle (384 ~ 322 BC)

Aristotle proposed 4 element theory of matter.

Earth

Water

Fire

Air

hot dry

wet cold

The school of thought laid out by Socrates, Plato and Aristotle dominated the western philosophy for 2000 years and the atomic theory of matter was completely buried.


John Dalton (1766 ~ 1844)

Dalton proposed that all matter is made up of atoms and stated that elements are the simplest form of matter.

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.


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 a 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

12g C + 16g O CO or 12g C + 32g O CO2 16g : 32g = 1:2

The Discovery of Atomic StructureThe Discovery of Atomic Structure

After Dalton’s atomic theory, not much of progress had been made andno one had direct evidence for the existence of atom. Then, things startedto change in late 1800s…

William Crooks (1832 ~ 1919): Cathode-ray tube (CRT) [1879]

A high voltage between two electrodes in a partiallyevacuated tube generates electrical discharge(cathode ray)


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

determine the charge to mass ratio

charge/mass = 1.76 108 C/g

He discovered that cathode raysare negatively charged particles, whichhe originally called ``corpuscles’’ . He won a Nobel prize in physics [1906].


Robert Millikan (1868 ~ 1953) : Determine the charge of electron [1907]

Millikan’s oil-drop experiment

Measured charge = 1.60 10-19 C

Electron mass = charge/[charge/mass]

= 9.10 10-28 g

The machine on the right hand side isthe original apparatus Millikan usedto perform his oil-drop experiment.He won a Nobel prize in physics [1923].


Ernest Rutherford (1871~1937): Discovery of nucleus [1911]

Rutherford’s -particle [4He2+]scattering experiment

He directed his graduate student Hans Geiger and undergraduate studentErnest Marsden to carry out -paticleexperiment. He won a Nobel prize in chemistry [1908].


Radioactivity: Generation of - particles

- ray: particles with +2 charge - ray: particles with 1 charge

- ray: high energy radiation with no charge


From the scattering experiment….

(1) Most -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.

Early Models of an Atom Early Models of an Atom

J. J. Thomson’s model “plum-pudding model”

Rutherford’s model

+

Rutherford's Model:

Electrons are negatively charged, but atoms as a whole are neutral.

Modern View of Atomic Structure Modern View of 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 behavior.

A convenient unit (non-SI) to describe the dimensions of atoms and molecules is Angstrom (Å). 1 Å = 1 10-10 m = 100 pm

Modern View of Atomic Structure Modern View of Atomic Structure

Properties of subatomic particles

Particle Charge (C) Mass (g) Mass (amu)

Proton +1.60 10-19 (+1) 1.6727 10-24 1.0073

Neutron 0 ( 0) 1.6750 10-24 1.0087

Electron 1.60 10-19 (1) 9.1097 10-28 5.486 10-4

Every atom has an equal number of protons and electrons so that it hasno electrical charge

Atomic Mass Unit (amu)

1 amu = 1/12 of the mass of carbon (12C) atom

= 1.66054 10-24 (g)

Modern View of Atomic StructureModern View of Atomic Structure

The characteristics of each atom are determined by the numbers ofproton, neutron and electrons.

Atomic Number: The number of protons in the nucleus of an atom. Mass Number: 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.

Hydrogen:1 proton

Helium:2 protons

2 neutrons

Lithium:3 protons

4 neutrons

Beryllium:4 protons

5 neutrons

Modern View of Atomic StructureModern View of Atomic Structure

Same information : An element is defined by the number of protons

Atomic WeightAtomic Weight

Atomic Mass Unit (amu) = 1.66054 10-24 g

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

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?

mass of C-12 mass of C-13fractional abundance ofC-12

(0.9893)(12 amu) + (0.0107)(13.00335) = 12.01 amu

This is the massof carbon atom shown in theperiodic table

Atomic WeightAtomic Weight

Example: Boron has two naturally occurring isotopes: 10B (10.01 amu)

and 11B (11.01 amu). If the average atomic weight of Boron

is 10.81, what are the fractional abundances of the two isotopes?

Periodic TablePeriodic Table

If the elements are arranged in order of increasing atomic number, their chemical properties are found to show a repeating, or periodic, pattern.

group

period

Elements having similar properties are placed in vertical columns

Periodic TablePeriodic Table

Alkalimetal

Alkaline earthmetal

Halogen

raregas

= H2, N2, O2, F2, Cl2, Br2, I2

Transition metals

Molecules and Molecular Compounds


Chemical Compounds

Molecular Ionic

(1) Molecular compounds are composed of more than

one type of atom

H2O, NH3, CH3OH, O2

(2) Most molecular substances contain only non-metallic atoms O2, H2O, H2O2, CO, CO2, CH4



Chemical Formulars

(1) Molecular Formulas : Indicate the actual numbers and types

of atoms in a molecule Ex. C2H4O2

(2) Empirical Formulas : Indicate the relative number of atoms of

each type in a molecule Ex. CH2O

(3) Structural Formulas : H O

H – C – C – O – H

H



Picturing Molecular Compounds (Ex. Methane)

Structural Formula

Ball-and-stick model

Perspective drawing

Space-filling model

Ions and Ionic CompoundsIons and Ionic Compounds

Ion : Atoms can readily gain or loose electrons and become ions.

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

Na+ Cl


Metals tend to form Cations

Nonmetals tend to form Anions

Alkali Metals

Alkaline Earth Metals

Noble Gases

Halogens

I A II A

III A IV A VA VI A VIIA

VIII A

Which elements form cations and which form anions?


How many electrons each element can gain or loose?

Each element tends to have the same number of electrons as noble gases (rare gases).


Example: Determine the number of electrons, protons and neutrons in each of the following ions

16O2-

40Ca2+

58Fe3+

80Br

No. of Protons No. of Neutrons No. of Electrons


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

NaCl

Alternating positive and negativecharges


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

use simplest 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


Example : Find the empirical formula for the ionic compound made of given cation and anion

Na, O =>

Al, O =>

Ca, O =>

Naming Ions and Ionic CompoundsNaming Ions and Ionic Compounds

Names of ionic compounds consist of the cation name followed by the anion name

CaCl2 = calcium + chloride calcium chloride

Names of Positive Ions (cations) :

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

as the metal.

Na+ sodium ion, Zn+ zinc ion, Al3+ aluminum ion

NOTE: Ions formed from a single atom are called monatomic ions


(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+ iron (II) ion Cu+ copper (I) ion

Fe3+ iron (III) ion Cu2+ copper (II) ion

These ions are usually transition metals

NOTE: Metals that form only one cation group 1A Na+, K+, Rb+

group 2A Mg2+, Ca2+, Sr2+, Ba2+

and Al3+ (group 3A), Ag+ (group 1B), Zn2+ (group 2B)


(3) Cations formed from nonmetal atoms have names that

end in -ium

NH4+ ammonium ion H3O+ hydronium ion

NOTE: These ions are examples of polyatomic ions


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 hydride ion, O2- oxygen oxide ion,

NOTE: polyatomic anions with common names ending with –ide OH- hydroxide ion, CN- cyanide ion

(2) Polyatomic anions containing oxygen (oxyanions)

a) ending with –ate : reserved for the most common oxyanion

NO3- nitrate ion, SO4

2- sulfate ion


b) ending with –ite : used for oxyanion with the same charge,

but one fewer O atom than those ending with –ate.

NO2- nitrite ion, SO3

2- sulfite ion

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

use prefix per- (one more) or hypo- (one fewer)

ClO4- perchlorate ion (one more than –ate)

ClO3- chlorate ion

ClO2- chlorite ion

ClO- hypochlorite ion (one fewer than -ite)


NOTE: Oxyanions with the maximum number of oxygens

(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 # + 1).

(iii) All names end with –ate except for ClO4-

(3) Anions derived by adding H+ to an oxyanion are named by

adding as a prefix the word hydrogen or dihydrogen.

CO32- : carbonate ion HCO3

- : hydrogen carbonate ion

PO43- : phosphate ion H2PO4

- : dihydrogen phosphate ion


Halogen (7A)

Names of Binary Molecular Compounds

Names of Binary Molecular Compounds

(1) The name of the element farther 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 (1 mono-, 2 di-, 3 tri-, 4 tetra-, 5 penta-, 6 hexa- )

Cl2O : dichloro monoxide NF3 : nitrogen trifluoride

N2O4 : dinitrogen tetroxide P4S10 : tetraphosphorous decasulfide

Naming Compounds : ExamplesNaming Compounds : Examples

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


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,


(1) NaClO :

(2) Fe2(CO3)3 :

(3) SF6 :

(4) aluminium hydroxide :

(5) ammonium sulfate :

(6) NaH2PO4 :

chemistry 101 : chap. 2 atoms, molecules and ions (1) atomic theory of matter (2) the discovery of...

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