(L2)atoms, ions and molecules

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Atoms, Ions and Molecules

Mr. Rumwald Leo G. LecarosSchool of Chemical Engineering and ChemistryMapua Institute of Technology

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1) Dalton’s Atomic Theory2) Atoms and Molecules3) Structure of the atom4) Subatomic particles5) Atomic Number, Mass Number and

Isotopes6) Ions and Ionic Compounds

Outline

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Matter is composed of atoms. ◦ Atoms

Composed of electrons, protons and neutrons.◦ Molecules

Combination of atoms.◦ Ions

Charged particles

The Structure of Matter

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John Dalton -1808 Postulates

1) An element is composed of extremely small, indivisible particles called atoms.

2) All atoms of a given element have identical properties that differ from those of other elements.

3) Atoms cannot be created, destroyed, or transformed into atoms of another element.

4) Compounds are formed when atoms of different elements combine with one another in small whole-number ratios.

5) The relative numbers and kinds of atoms are constant in a given compound.

Dalton’s Atomic Theory

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Law of conservation of mass◦ Matter is conserved in chemical reactions

Law of constant composition◦ The elements that a compound is composed of

are present in fixed and precise proportion by mass.

Law of multiple proportions◦ When the same elements can form two different

compounds, the ratio of masses of one of the elements in the two compounds is a small whole number relative to a given mass of the other element.

Fundamental Laws of Matter

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Dalton’s atom had no features other than mass.

Many discoveries have demonstrated that the atom is not featureless or indestructible, but is composed of other parts. (subatomic parts)

J.J. Thomson and Ernest Rutherford

Components of the Atom

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First evidence for subatomic particles came from the study of the conduction of electricity by gases at low pressures.◦ J.J. Thomson, 1897◦ Rays emitted were called cathode rays◦ Rays are composed of negatively charged

particles called electrons.◦ Electrons carry a negative charge (-1) and have

a very small mass (1/2000 the lightest atomic mass)

Electrons

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Cathode Ray Apparatus

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J.J. Thomson modified the cathode ray tube in 1897 by adding two adjustable voltage electrodes.◦ Studied the amount that the cathode ray beam

was deflected by additional electric field. Thomson used his modification to measure

the charge to mass ratio of electrons.◦ e/m = -1.75881 x 108 coulomb/g of e◦ He named the cathode rays electrons◦ “discoverer of electrons”◦ TV sets and computer screens

The Discovery of Electrons

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Robert A. Millikan won the 1stAmerican Nobel Prize in 1923 for his famous oil-drop experiment.

In 1909 Millikan determined the charge and mass of the electron.

The Discovery of Electrons

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Millikan determined that the charge on a single electron = -1.60218 x 10-19coulomb.

Using Thomson’s charge to mass ratio we get that the mass of one electron is 9.11x 10-28g.◦ e/m = -1.75881 x 108coulomb◦ e = -1.60218 x 10-19coulomb◦ Thus m = 9.10940 x 10-28g

The Discovery of Electrons

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Eugene Goldstein noted streams of positively charged particles in cathode rays in 1886.◦ Particles move in opposite direction of cathode

rays. ◦ Called “Canal Rays” because they passed through

holes (channels or canals) drilled through the negative electrode.

Canal rays must be positive.◦ Goldstein postulated the existence of a positive

fundamental particle called the “proton”.

Canal Rays and Protons

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Canal Rays and Protons

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Ernest Rutherford directed Hans Geiger and Ernst Marsden’s experiment in 1910.◦ α-particle scattering from thin Au foils ◦ Gave us the basic picture of the atom’s structure.

Rutherford and the Nuclear Atom

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Ernest Rutherford, 1911 Bombardment of gold foil with α-particles

(helium atoms minus their electrons◦ Expected to see the particles pass through the foil◦ Found that some of the alpha particles were

deflected by the foil◦ Led to the discovery of a region of heavy mass at

the center of the atom

Protons and Neutrons – The Nucleus

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Rutherford’s major conclusions from the -particle scattering experiment

1. The atom is mostly empty space.2. It contains a very small, dense center

called the nucleus.3. Nearly all of the atom’s mass is in the

nucleus.4. The nuclear diameter is 1/10,000 to

1/100,000 times less than atom’s radius.

Rutherford and the Nuclear Atom

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James Chadwick in 1932 analyzed the results of α-particle scattering on thin Be films.

Chadwick recognized existence of massive neutral particles which he called neutrons.◦ Chadwick discovered the neutron.

Neutrons

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Atomic Structure

Solar system depiction of atomic structure.

◦ Emphasizes proton, neutron and electron distribution; does not accurately depict current accepted model of atomic structure.

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Atomic Structure Electrons are

depicted as clouds of negative charge surrounding the nucleus.

◦ The density of the small dots is related to the probability of finding an electron at a particular location.

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Subatomic Particles

Particle mass (amu) charge

Proton 1.007 +

Neutron 1.009 0

Electron 0.00055 –

Atomic Symbols Information regarding atomic structure is

written in scientific shorthand called the atomic symbol.

◦ E is the atomic symbol for element◦ Superscript A is the mass number.◦ Subscript Z is the atomic number.

ZAE

Atomic Symbols

Atomic Masses

Entry for carbon on the periodic table.

◦ Z = 6

◦ Relative atomic mass =   12.011 (~99% carbon-12)

◦ Element Symbol: C

612C

Isotopes Isotopes are atoms of an element that differ

in the number of neutrons in their nucleus.

◦ same Z but different A

Isotopic abundance is the mass percentage of an isotope in a naturally occurring element.

Isotopes

Mass spectrometers can measure the masses of atoms, isotopes, and molecules.

Isotopes

Mass spectrum showing carbon isotopes.

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One example of an isotopic series is the hydrogen isotopes.◦ 1H or protium is the most common hydrogen

isotope. one proton and no neutrons

◦ 2H or deuterium is the second most abundant hydrogen isotope. one proton and one neutron

◦ 3H or tritium is a radioactive hydrogen isotope. one proton and two neutrons

Mass Number and Isotopes

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1H, 2H, 3H◦ Hydrogen, deuterium, tritium◦ Different masses

Note that some of the ice is at the bottom of the glass –this is 2H2O

Isotopes of Hydrogen

Atomic Masses Relative atomic mass for an element is an

average of the atomic masses for the naturally occurring isotopes for an element.

Example Naturally occuring Cu consists of 2 isotopes.

It is 69.1% 63Cu with a mass of 62.9 amu, and 30.9% 65Cu, which has a mass of 64.9 amu. Calculate the atomic weight of Cu to one decimal place.

Atomic weight = Σ (relative abundance) (isotopic mass)

Atomic weight = (0.691)(639 amu) + (0.309)(649 amu)

Atomic weight = 63.5 amu

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Naturally occurring argon consists of three isotopes, the atoms of which occur in the following abundances: 0.34% 36Ar (35.9676amu), 0.07% 38Ar (37.9627), and 99.59% 40Ar (39.9624). Calculate the atomic weight of argon from these data.

Daily Exercise

Atomic weight = 39.95 amu

Ions Ions are formed when the number of

protons and electrons in an atom are not equal.

◦ Ions with more protons than electrons are called cations. net positive charge

◦ Ions with more electrons that protons are called anions. net negative charge

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A monatomic ion is derived from a single atom.

A polyatomic ion is derived from a group of atoms with an overall charge.

Ions

Ions

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Groups of atoms may carry a charge; these are the polyatomic ions◦ OH-◦ NH4+

Polyatomic Ions

Ions and Their Properties An element and its ion have the same

chemical symbol but different properties.

◦ Sodium metal atoms lose an electron to form sodium cations. Sodium metal reacts violently with water.

◦ Chlorine gas molecules gain electrons to form chlorine anions (chloride). Chlorine gas reacts violently with sodium metal.

◦ Ionic compounds containing sodium cation and chlorine anion dissolve in water without reacting.

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Determine the number of protons, electrons, and neutrons in:

(a)

(b)

Example