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. History of the Atom

V. Subatomic Particles

VI. Atomic Mass

VII. Atoms and the Mole

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. 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

IV. Atomos vs. Fire, Air, Earth,

Water

IV. 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.

IV. Law of Mass Conservation

• In a reaction, matter

is neither created

nor destroyed.

• Credit Antoine

Lavoisier (1789).

IV. Law of Mass Conservation

IV. 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:

IV. 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.

IV. 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.

IV. 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.

IV. 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.

IV. Cathode Rays

• What conclusions about cathode rays can be

made from these experiments?

IV. 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.

IV. 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.

IV. 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

IV. 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.

IV. Rutherford’s a-Particle

Experiment (1909)

IV. 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.

IV. Rise of the Nuclear Atom

V. Subatomic Particles

• Therefore, all atoms are made up of

protons, neutrons, and electrons.

V. 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.

V. 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

V. 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.

V. 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.

V. Depicting an Isotope

V. Sample Problem 1.1

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?

V. 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-

VI. 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.

VI. 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.

VI. Sample Problem 1.2

• Lithium has an atomic mass of 6.941

amu and has two naturally-occuring

isotopes: 6Li and 7Li. These isotopes

have isotopic masses of 6.0151 amu

and 7.0160 amu, respectively. What

are the percent natural abundances (to

2 decimal places) of these isotopes of

lithium?

VII. How Much vs. How Many

• Countable vs. not countable

• In the lab, we say, “How much water do

we need?”

• However, matter is particulate, so it is

countable.

• When matter interacts, it does so on a

particle by particle basis.

VII. Using Mass to “Count”

• We can relate counts to mass.

e.g. 21-30 shrimp vs. 8-10 shrimp

• In chemistry, we use the mole.

A mole is the amount of material that

contains 6.02214 × 1023 particles.

Defined by the # of atoms in exactly 12 g of

pure carbon-12.

This number is Avogadro’s number.

VII. Avogadro’s Number

• We can use Avogadro’s number as a conversion factor to calculate # of atoms.

6.022 x 1023 atoms

1 mole atoms 6.022 x 1023 atoms

1 mole atomsOR

VII. Grams and Moles

• The mass of 1 mole of atoms of an element is its molar mass. The value of an element’s molar mass in g/mole is numerically equal to the element’s atomic mass in amu (from periodic table).

• In general, one mole scales up to a “touchable” amount of an element or molecule.

VII. Mole of Atoms & Compounds

• 1 Fe atom weighs 55.85 amu, so 1 mole

of Fe atoms weighs 55.85 g.

• 1 O atom weighs 16.00 amu, so 1 mole

of O atoms weighs 16.00 g.

• Same applies for compounds.

1 molecule H2O weighs 18.02 amu, so 1

mole H2O weighs 18.02 g.

VII. Sample Problem 1.3

• If a silicon chip has a mass of 5.89 mg,

how many Si atoms are in the chip?