Ch. 2: Atoms and ElementsCh. 2: Atoms and Elements

Dr. Namphol Sinkaset

Chem 200: General Chemistry I

I. Chapter OutlineI. Chapter Outline

I. Introduction

II. The Atomic View of Matter

III. The Nuclear Atom

IV. Structure of the Atom

V. Atomic Mass

I. MatterI. Matter

• Why study matter? Greeks first wondered about its origin

Leucippus and Democritus vs. Plato and Aristotle

Discover useful properties Explain behavior of the natural world

II. The Atomic View of MatterII. The Atomic View of Matter

• The word “atom” comes from the Greek word “atomos.”

• Idea was discarded, but then revived by the end of the 18th century.

• At that time, 3 natural laws were begging for an explanation.

II. The Law of Mass ConservationII. The Law of Mass Conservation

• In a reaction, matter is neither created nor destroyed.

• Credit Antoine Lavoisier.

II. Law of Definite ProportionsII. Law of Definite Proportions

• All samples of a given compound, regardless of their source or how they were prepared, have the same proportions of their constituent elements.

• Credit Joseph Proust

II. Sample ProblemII. Sample Problem

• e.g. Pitchblende is a source of uranium. If 84.2 g of pitchblende contains 71.4 g uranium, w/ oxygen as the only other element, how many grams of uranium can be extracted from 102 kg of pitchblende?

II. Law of Multiple ProportionsII. Law of Multiple Proportions

• John Dalton found that when two elements (A and B) form two different compounds, the masses of element B that combine with 1 g of element A can be expressed as a ratio of small whole numbers.

II. Law of Multiple ProportionsII. Law of Multiple Proportions

• What?• Compounds form when atoms of elements

form in whole number ratios like 1:2, 1:3, but not 1.34:2.66.

II. Dalton’s Atomic TheoryII. Dalton’s Atomic Theory

• John Dalton revived the idea of the atom to explain the natural laws that had everyone perplexed.

• His atomic theory worked so well that it was quickly accepted.

II. Postulates of Dalton’s TheoryII. Postulates of Dalton’s Theory1. Each element is composed of tiny,

indestructible particles called atoms.2. All atoms of a given element have the same

mass and other properties that distinguish them from the atoms of other elements.

3. Atoms combine in simple, whole-number ratios to form compounds.

4. Atoms of one element cannot change into atoms of another element. In a chemical reaction, atoms change the way that they are bound together with other atoms to form a new substance.

III. The Nuclear AtomIII. The Nuclear Atom

• Dalton’s theory treated atoms as permanent, indestructible building blocks that composed everything.

• A series of experiments were conducted that led to a new view of the atom.

III. Cathode RaysIII. Cathode Rays

• What conclusions about cathode rays can be made from these experiments?

III. Cathode RaysIII. Cathode Rays

• Using EM fields, J.J. Thomson measured the cathode ray particle’s mass to charge ratio.

• He estimated that cathode ray particles were about 2000 times lighter than a hydrogen atom.

• Result implies that atoms can be divided into smaller particles.

III. Cathode RaysIII. Cathode Rays• Using his famous oil drop experiment,

Robert Millikan calculated the charge of a cathode ray particle.

• His value is w/in 1% of today’s accepted value: -1.602 x 10-19 C.

• Mass was determined to be 9.109 x 10-28 g.

• Of course, cathode ray particles are now known as electrons.

III. Plum PuddingIII. Plum Pudding• If electrons are in all

matter, there must be positively-charged species as well.

• J.J. Thomson proposed the plum pudding model of the atom. Electron “raisins” “Pudding” of positive

charge

III. The Role of RadioactivityIII. The Role of Radioactivity

• Marie and Pierre Curie discovered radioactivity by accident.

• Ernest Rutherford used radium, an alpha () particle emitter.

• These -particles are dense and have a positive charge.

III. Rutherford’s III. Rutherford’s -Particle -Particle ExperimentExperiment

III. Conclusions from III. Conclusions from Rutherford’s ExperimentRutherford’s Experiment

• Most of an atom’s mass and all of its positive charge exists in a nucleus.

• Most of an atom is empty space, throughout which electrons are dispersed.

• By having equal numbers of protons and electrons, an atom remains electrically neutral.

• Note: neutrons discovered 20 years later.

III. Rise of the Nuclear AtomIII. Rise of the Nuclear Atom

IV. Structure of the AtomIV. Structure of the Atom

IV. Atomic NumberIV. Atomic Number

• The atomic number (Z) of an element equals the # of protons in the nucleus All atoms of an element have same, unique

atomic number!!

• Protons are responsible for an atom’s identity.

• e.g. All carbon atoms have 6 protons and all uranium atoms have 92 protons.

IV. Chemical SymbolsIV. Chemical Symbols

• Each element has a unique symbol.

• The symbol is either a 1 or 2 abbreviation of its name.

• e.g. carbon C; nitrogen N; chlorine Cl; sodium Na; gold Au

IV. Mass NumberIV. Mass Number

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

• e.g. A carbon atom with 6 neutrons has a mass number of 12.

IV. Depicting an AtomIV. Depicting an Atom

IV. IsotopesIV. Isotopes• The # of protons determines the identity

of the atom, but the # of neutrons has no effect.

• Thus, atoms of the same element can have different mass numbers.

• Since chemical properties are mainly due to e-, isotopes are almost identical chemically.

• Different isotopes of an element exist in certain percentages – natural abundances.

IV. IsotopesIV. Isotopes

• e.g. Determine the number of protons and neutrons in carbon-12, carbon-13, and carbon-14.

V. Masses of AtomsV. Masses of Atoms

• The mass of an atom is measured relative to the mass of a standard.

• The modern standard is carbon-12, which is assigned a mass of exactly 12 atomic mass units (amu).

• e.g. On this scale, a hydrogen atom has a mass of 1.008 amu.

• Note that 1 amu = 1.66054 x 10-24 g.

V. Atomic Mass or Atomic V. Atomic Mass or Atomic WeightWeight

• The abundances of the isotopes are determined by nature.

• The atomic mass or atomic weight is the average of the masses of all isotopes of an element weighted according to their abundance.