Ch. 1: Atoms

Dr. Namphol Sinkaset

Chem 200: General Chemistry I

I. Chapter Outline

I. Introduction

II. Particulate View of the World

III. The Scientific Approach

IV. Measurement in Science

V. History of the Atom

VI. Subatomic Particles

VII. Atomic Mass

I. Real-Life Legos®

• Everything is comprised of small parts connected into a complex whole.

• The structure of the whole determines its properties.

II. Particles

• We will approach chemistry with two key principles in mind:

1. Matter is particulate.

2. Structure of particles determines properties of matter.

• Chemistry seeks to understand properties of matter by studying the structure of particles that compose it.

II. Matter

• Matter is anything that occupies space and has mass.

• Everything around you is composed of matter – desk, book, air.

• Remember: matter is particulate.

II. Atoms and Molecules

• Atoms are the basic particles that compose ordinary matter.

• Atoms can bind to one another in specific arrangements to yield molecules.

• For example, a water molecule is comprised of 1 oxygen atom and 2 hydrogen atoms.

II. Structure and Properties

• Boils at 30 °C• Feels like gasoline• Doesn’t dissolve salt

• Boils at 100 °C• Feels like water• Dissolves salt

II. Classifying Matter

• Any sample of matter is called a substance.

• Matter can be classified by state or by composition.

• State determined by relative positions and interactions of particles.

• Composition determined by types of particles.

II. States of Matter

II. States of Matter

• solid: strong particle attractions, pack in fixed locations, only vibrate in place, not compressible

• liquid: slightly weaker particle attractions, pack in non-fixed locations, fixed volume, assume shape of container

• gas: weak particle attractions, free to move, large distances between particles, compressible

II. Composition of Matter

• Can also classify matter by the kinds of particles out of which it is comprised.

• If there is only one type of particle, then it is a pure substance.

• If there is more than one type of particle, then it is a mixture.

II. Types of Matter

III. The Scientific Approach

• a.k.a. scientific method, is a flexible process of creative thinking and testing aimed at an objective

III. Differences Between Hypothesis and Theory

• Hypothesis not thoroughly tested• Theory more “developed”• Hypothesis does not predict• Experiments on hypothesis test

hypothesis itself• Experiments on theory test predictions

of theory• Theory can be expanded to many

related situations

III. Differences Between Hypothesis and Theory

• Compare the two statements below.

• “Methane reacts w/ oxygen to form carbon dioxide and water.”

• “Hydrocarbons undergo a combustion reaction w/ oxygen to form carbon dioxide and water.”

IV. Measurement in Science

• Observations in the lab can be qualitative or quantitative.

• Science is powerful because many observations can be assigned an accurate number.

• e.g. Hot/cold vs. 262 °C/12 °C

IV. Units

• All measured quantities have a number and a unit!!!!

• Without a unit, a number has no meaning in science.

• e.g. a person’s height is 6 feet, 4 inches, not 6-4.

• ANY ANSWER GIVEN W/OUT A UNIT WILL BE GRADED HARSHLY.

IV. SI Units

• Scientists have agreed to use the Système International d’Unités, a.k.a. SI units

IV. Derived SI Units

• Combinations of fundamental SI units are used to describe other quantities.

• e.g. speed is distance per time, so its SI unit is m/s

V. History of the Atom

• The Greeks were the first to wonder about matter.

• Greek philosophers around 430 B.C.E. debated what made up the world around them.

• Leucippus and Democritus vs. Plato and Aristotle

V. Atomos vs. Fire, Air, Earth, Water

V. Revival of the Atom

• The idea of the atom was discarded and forgotten about for almost 2000 years.

• In the late 18th and early 19th centuries, three natural laws baffled everyone.

• John Dalton resurrected the idea of the atom to explain what was observed.

V. Law of Mass Conservation

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

• Credit Antoine Lavoisier (1789).

V. Law of Mass Conservation

V. Law of Definite Proportions• All samples of a given compound have

the same proportions of constituent elements. Credit Joseph Proust (1797)

• e.g. Ammonia has 14.0 g N for every 3.0 g of H:

𝑀𝑎𝑠𝑠𝑟𝑎𝑡𝑖𝑜=14.0𝑔𝑁3.0𝑔 𝐻

=4.7𝑜𝑟 4.7 :1

V. Law of Multiple Proportions• In 1804, 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.

V. Dalton’s Atomic Theory

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

• His atomic theory (1808) worked so well that it was quickly accepted.

V. Postulates of Dalton’s Theory

1. 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 only change the way that they are bound together with other atoms.

V. 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.

V. Cathode Rays

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

V. Cathode Rays

• Using EM fields (late 1800s), 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.

V. Cathode Rays• Using his famous oil drop experiment,

Robert Millikan (1909) 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.

V. 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

V. The Role of Radioactivity

• Henri Becquerel and Marie Curie discovered radioactivity by accident.

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

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

V. Rutherford’s a-Particle Experiment (1909)

V. Conclusions from Rutherford’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.

V. Rise of the Nuclear Atom

VI. Subatomic Particles

• Therefore, all atoms are made up of protons, neutrons, and electrons.

VI. 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.

VI. 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

VI. 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.

VI. 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.

VI. Depicting an Isotope

VI. Sample Problem

1. What are the atomic number, mass number, and symbol for the carbon isotope with 7 neutrons?

2. How many protons and neutrons are present in an atom of potassium-39?

VI. Ions

• Atoms can lose or gain electrons and become ions.

• Ions are charged particles.• Positively-charged particles = cations.

e.g. Li Li+ + 1e-

• Negatively-charged particles = anions. e.g. F + 1e- F-

VII. Atomic Mass

• Postulate #2 of Dalton’s atomic theory stated that all atoms have the same mass.

• With the existence of isotopes, this can’t be true, but we can calculate an average mass.

• The average mass of an element is called the atomic mass.

VII. Atomic Mass

• Atomic masses are calculated using a weighted average of all isotopes of an element.

• The natural abundance is used to weight each isotope in the calculation.

𝐴𝑡𝑜𝑚𝑖𝑐𝑚𝑎𝑠𝑠=∑𝑛

( 𝑓𝑟𝑎𝑐𝑡𝑖𝑜𝑛 𝑖𝑠𝑜𝑡𝑜𝑝𝑒𝑛)×(𝑚𝑎𝑠𝑠𝑖𝑠𝑜𝑡𝑜𝑝𝑒𝑛)